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 Essentials of volumetric analysis; an Int 
 
 3 1924 002 ■974""925"""" 
 
Cornell University 
 Library 
 
 The original of tliis book is in 
 tlie Cornell University Library. 
 
 There are no known copyright restrictions in 
 the United States on the use of the text. 
 
 http://www.archive.org/details/cu31924002974925 
 
WORKS OF DR. H. W. SCHIMPF 
 
 PUBLISHED BY 
 
 JOHN WILEY & SONS. 
 
 A Manual of Volumetric Analysis. 
 
 For the use of Pharmacists, Sanitary and Food 
 Chemists, as well as for Students in these 
 Branches. Fifth edition, rewritten and en- 
 larged, 8vo, XX -I- 725 pages, loa figures. 
 Cloth, Ss.oo. 
 
 Essentials of Volumetric Analysis. 
 
 An in troduc tion to the subject adapted to the 
 needs of Students of Pharmaceutical Chemistry. 
 Second edition, rewritten. Large i2mo, xL + 3s8 
 pages, 61 figures. Cloth, $1.50 net. 
 
 A Systematic Course of Qualitative Chemical 
 Analysis of lnorg;anic and Organic Substances 
 
 With Explanatory Notes. 8vo, vii+156 pag-es. 
 Cloth, I1.25 net. 
 
ESSENTIALS 
 
 OF 
 
 VOLUMETRIC ANALYSIS 
 
 AN INTRODUCTION TO THE SUBJECT, ADAPTED 
 TO THE NEEDS OF STUDENTS OF PHAR- 
 MACEUTICAL CHEMISTRY 
 
 EMBRACING THE SUBJECTS OF ALKALIMETRY, ACIDIMETRY, PRECIPI- 
 TATION ANALYSIS, OXIDIMETRY, INDIRECT OXIDATION, lODOM- 
 ETRY, ASSAY PROCESSES FOR DRUGS, ESTIMATION OF AL- 
 KALOIDS, CARBOLIC ACID, SUGARS, THEORY, APPLI- 
 CATION AND DESCRIPTION OF INDICATORS 
 
 BY 
 
 HENRY W. SCHIMPF, Ph.G., M.D. 
 
 Professor of Analytical Chemistry iti the Brooklyn College of Pharmacy 
 
 JIlUtBtratsb 
 
 SECOND EDITION— REWRITTEN AND ENLARGED 
 FIRST THOUSAND 
 
 NEW YORK 
 
 JOHN WILEY & SONS 
 
 London : CHAPMAN & HALL, Limited 
 
 1911 
 
/Copyright, 1903, 1911 
 
 BY 
 
 HENRY ^y. SCHIMFF 
 
 THE SCIENTIFIC PRESS 
 
 KOBEHT DRUMMOND AND COMPANY 
 
 BnOOKLYN, N. Y, 
 
PREFACE TO THE FIRST EDITION 
 
 Tke growing need for a short text-book which will make 
 the principles of volumetric analysis readily available without 
 going too deeply into detailed and discursive description has 
 led to the preparation of this elementary treatise. 
 
 In the following pages the aim is to present the principles 
 of this interesting and important subject in a form readily 
 intelligible to students and available for lecture-room and 
 laboratory work. The essential points are condensed within 
 the limits of a small book with the intention of furnishing 
 an outline which may serve as a practical guide as well as 
 an introduction to the more advanced and voluminous works 
 on the subject. 
 
 If presented in a suitable manner volumetric analysis 
 rarely fails to prove interesting to the student, because it gives 
 him a clear conception of the quantitative significance of 
 chemical equations and thus affords practical proofs of chemical 
 laws; it furthermore trains the student to make careful obser- 
 vations, to form habits of accuracy in manipulation, and since 
 the processes are easily carried out, enables him to arrive 
 readily at a definite numerical conclusion. 
 
 The subject-matter in this book is systematically arranged 
 as far as can be, and treated as concisely as is consistent with 
 clearness of expression. The processes are grouped under 
 five headings: Neutralization, Precipitation, Oxidation, Indirect 
 Oxidation, and lodometry. The principles underlying each 
 
IV PREFACE 
 
 group are definitely indicated, and their application illustrated 
 by numerous practical examples. Other subjects treated 
 include methods of calibration and of the accurate reading 
 of graduated instruments, the calculation of the results of 
 analyses, the preparation and standardization of volumetric 
 solutions. The indicators, their selection for special cases 
 and the ionic theory regarding their action, as well as assay 
 process for phenol, sugars and vegetable drugs also receive 
 special treatment. The author hopes that he has prepared 
 a book which will prove serviceable to those for whom it was 
 written and that it will be as generously received as were the 
 four editions of his Text-book of Volumetric Analysis. 
 
 Henry W. Schimpf. 
 
PREFACE TO THE SECOND EDITION 
 
 The favor with which the first edition of The Essentials 
 of Volumetric Analysis was received has encouraged me to 
 revise and enlarge the book. 
 
 In the present edition I have endeavored to keep in mind 
 the needs of students studying volumetric analysis and to 
 give them a competent and practical guide to this most interesting 
 subject. The main features of the former edition have been 
 retained, but the subject is brought up to date. The ionization 
 theory is described especially in its relation to indicators, and 
 the atomic weights used are those for 191 1 of the Inter- 
 national Committee on Atomic Weights (0 = i6). The book 
 contains sixty-one illustrations, and numerous useful tables. 
 
 The volumetric methods are arranged in a systematic 
 manner and comprise Alkalimetry, Acidimetry, Precipitation, 
 Analysis Involving the Use of Silver Nitrate, Sodium Chlorid, 
 and Potassium Sulphocyanate. Oxidation Methods involving 
 the Use of Permanganate, Bichromate and lodin. Reduction 
 Methods Involving the Use of Sodium Thisosulphate, Arsenous 
 Acid and Stannous Chlorid. 
 
 There are also given concise descriptions of methods for 
 assaying alkaloidal drugs, phenol, oils, sugars, formaldehyde, 
 and alcoholic liquids, together with a few simple gasometric 
 analyses, such as a pharmacist may find useful. 
 
 Henry W. Schimpf. 
 
CONTENTS 
 
 PAGE 
 
 List of Elements with their Atomic Weights xii 
 
 Table of Multiples of Atomic Weights and Combinations. . . xiii 
 
 CHAPTER I 
 Introduction r 
 
 CHAPTER n 
 
 General Principals of Chemical Combination upon which 
 Volumetric Analysis is Based 4 
 
 CHAPTER III 
 
 Volumetric or Standard Solutions 7 
 
 To Titrate. Residual Titration. 
 
 CHAPTER IV 
 
 Indicators 16 
 
 The Ionization Theory. The Ionization Theory of Indicators. 
 A Guide for the Selection of Indicators. 
 
 CHAPTER V 
 Apparatus Used in Volumetric Analysis 28 
 
 CHAPTER VI 
 
 On the Use of Apparatus 39 
 
 On the Reading of Instruments. Calibration of Instruments. 
 
 vii 
 
viii CONTENTS 
 
 CHAPTER VII 
 
 PAGE 
 
 Methods of Calculating Results 48 
 
 Table of Normal Factors, etc., of Alkalies, Acids and Alkali 
 Earths. On Stating Results. Table of Molecular Weights and 
 Normal Factors for the most Common Oxids. 
 
 CHAPTER VIII 
 
 Analysis by Neutralization 56 
 
 Alkalimetry. Preparation of Standard Acid Solutions. Esti- 
 mation of Alkali Hydroxids. Estimation of Alkali Carbonates, 
 Mixed Alkali Hydroxid and Carbonate. Estimation of Alkali 
 Bicarbonates when Mixed with Carbonates. Mixed Potassium 
 and Sodium Hydroxids. Estimation of Organic Salts of the 
 Alkalies. Table of Normal Factors, etc., of the Organic Salts of 
 the Alkalies. Estimation of Salts of the Alkali Metals. Esti- 
 mation of Mixed Hydroxids and Carbonates of Alkali Earths. 
 Acidimetry. Estimation of Acids. Preparation of Standard 
 Alkali Solutions. Table Showing Quantity to be taken for 
 Analysis in Direct Percentage Estimations. 
 
 CHAPTER IX 
 
 Analysis by Precipitation 105 
 
 Preparation of Decinormal Silver Nitrate, Decinormal Sodium 
 Chlorid, and Decinormal Sulphocyanate. Estimation of Soluble 
 Haloid Salts. Mohr's Method with Cbromate Indicator. Titra- 
 tion without an Indicator. Estimation of Haloid Acids. Esti- 
 mation of Cyanogen. Estimation of Silver Salts. Estimation 
 of Metallic Silver and Silver Alloys. Estimation of lodids by 
 Mercuric Chlorid Solution. Table of Substances Estimated by 
 Precipitation. 
 
 CHAPTER X 
 
 Analysis by Oxidation and Reduction 133 
 
 Preparation of Decinormal Potassium Permanganate. 
 Volumetric Analyses by Means of Potassium Permanganate. 
 On the Use of Empirical Permanganate Solution. Typical 
 Analyses by Permanganate. Direct Titrations. Estimation of 
 
CONTENTS ix 
 
 PAGE 
 
 Ferrous Salts. Estimation of Metallic Iron in Ferrum Reductum. 
 Estimation of Oxalic Acid and Oxalates. Estimation of Calcium. 
 Estimation of Hydrogen Dioxid and Barium Dioxid. Estimation 
 of Ferric Salts. Estimation of Nitrous Acid and Nitrites. Resid- 
 ual Titrations. Estimation of Hypophosphorous Acid and Hypo- 
 phosphites. Estimation of Calcium Salts. Estimation of Lead 
 Acetate and Subacetate. Estimation of Manganese Dioxid. 
 Estimation of Nitrates, Chromates, and Chromic Acid. Esti- 
 mation of Tin. Estimation of Copper. 
 
 Volumetric Atmlysis by Means of Potassium Dichromale. 
 Preparation of Decinormal Potassium Bichromate. Estimation 
 of Ferrous Salts. Table of Substances Estimated by Perman- 
 ganate or Bichromate. 
 
 Analysis by Indirect Oxidation. Preparation of Becinormal 
 lodin. Starch Solution. On the Use of Sodium Bicarbonate 
 in Titrations with lodin. Estimation of Arsenous Compounds. 
 Estimation of Antimony Compounds. Estimation of Sulphurous 
 Acid and Sulphites. Estimation of Sodium Thiosulphate. Hydro- 
 gen Sulphid and Sulphids. Estimation of Metallic Iron in 
 Reduced Iron. Table of Substances Estimated by INIeans of 
 lodin Solution. 
 
 Estimation of Substances Readily Reduced. lodometry. Esti- 
 mations Involving the Use of Sodium Thiosulphate V.S. Prepa- 
 ration of Becinormal Thiosulphate V.S. Estimation of Free 
 lodin. Indirect lodo metric Estimations. Estimation of Free 
 Chlorin or Bromin. Estimation of Available Chlorin. Estima- 
 tion of Hydrogen Bioxid. Bistillation Methods. Estimation of 
 Manganese Bioxid. Chromic Acid and Chromates. Estimation 
 of Alkali lodids. Digestion Methods. Estimation of Chlorates, 
 Bromates and lodates. Estimation of Ferric Salts. 
 
 Chloromelry, Reduction Methods, Involving the Use of Arsenous 
 Acid Solutions. Preparation of Standard Arsenous Acid V.S. 
 Iodized Starch Test Paper. Estimation of Free Halogens. Esti- 
 mation of Available Chlorin. Chlorometric Assay of Manganese 
 Bioxid. 
 
 Reduction Methods Involving the Use of Stannous Chlorid V.S. 
 Estimation of Iron by Means of Stannous Chlorid. Estimation 
 of Mercuric Salts. 
 
X CONTENTS 
 
 PART II 
 CHAPTER XI 
 
 PAGE 
 
 Estimation op Alkaloids 249 
 
 Table of Factors for Alkaloids. Gordin's Modified Alkali - 
 metric Method for Titrating Alkaloids. 
 
 CHAPTER XII 
 
 Assaying of Vegetable Drugs and their Preparation 257 
 
 Separation of Alkaloids and Use of Immiscible Solvents. 
 Kebler's Modification of the Keller Method. Assay of Galenical 
 Preparations. Lloyd's Method. Katz's Method. 
 
 CHAPTER XIII 
 Estimation of Phenol 269 
 
 Preparation of Decinormal Bromin Solution. 
 
 CHAPTER XIV 
 
 Some Technical Methods for Fats, Oils and Waxes 275 
 
 The Acid Value. The Saponification Number. Volatile Fatty 
 Acid Number. Reichert's Number. The Reichert-Meissl Num- 
 ber. Hubl's Number. 
 
 CHAPTER XV 
 
 Estimation of Suc-^rs 287 
 
 Preparation of Fehling's Solution. Determination of the End- 
 point. Estimation Starch after Inversion. Estimation of Maltose 
 in Malt Extracts. Estimation of Diastasic Value of Malt Extract. 
 
 29s 
 
 CHAPTER XVI 
 
 Estimation of Formaldehyde 
 
 The Ammonia Method. The Ammonium Chlorid Method. 
 Oxidation Method by Means of Hydrogen Dioxid. The lodo- 
 metric Method. The Cyanid Method. 
 
 CHAPTER XVII 
 
 Estimation of Alcohol in Tinctures and Beverages 303 
 
 Alcoholometric Table 
 
CONTENTS xi 
 
 PART III 
 CHAPTER XVIII 
 
 PAGE 
 
 The Nitrometer 307 
 
 The Law of Charles. The I^aw of Boyle. 
 
 CHAPTER XIX 
 
 Assay of Nitrites 311 
 
 Spirit of Nitrous Ether. Amyl Nitrite. Sodium Nitrite. 
 Nitric Acid in Nitrates. 
 
 CHAPTER XX 
 
 Hydrogen Dioxid 317 
 
 Use of Nitrometer. Use of Urea Apparatus. The Hypo- 
 chlorite Method. The Hypobromite Method. Table Showing 
 Weight in Milligrams of H2O2 corresponding to one cc. of Moist 
 Oxygen. 
 
 CHAPTER XXI 
 
 ICSTIMATION OF SOLUBLE CARBONATES BY THE USE OF THE NITROM- 
 ETER 322 
 
 CHAPTER XXII 
 
 ESTIMAITON OF UREA IN URINE 323 
 
 The Doremus' Ureometer. The Hinds-Doremus Ureometer. 
 Squibb's Urea Apparatus. 
 
 APPENDIX 
 
 Description of Indicators Alphabetically Arranged 328 
 
A LIST OF THE MORE COMMON ELEMENTS WITH THEIR 
 SYMBOLS AND ATOMIC WEIGHTS 
 
 Atomic Weight* 
 based on 0= i6. 
 
 Atomic Weight 
 
 based on H=i. Atomic Weight 
 
 Approximate 
 
 Aluminium Al. 
 
 Antimony Sb. 
 
 Arsenic As. 
 
 Barium Ba. 
 
 Bismuth Bi. 
 
 Boron B. 
 
 Bromin Br. 
 
 Cadmium Cd. 
 
 Calcium Ca. 
 
 Carbon C. 
 
 Chlorin CI. 
 
 Chromium Cr. 
 
 Cobalt Co. 
 
 Copper Cu. 
 
 Fluorin F. 
 
 Gold Au. 
 
 Hydrogen H . 
 
 lodin I. 
 
 Iron Fe. 
 
 Lead Pb. 
 
 Lithium Li. 
 
 Magnesium Mg. 
 
 Manganese Mn. 
 
 Mercury Hg. 
 
 Molybdenum Mo. 
 
 Nickel Ni. 
 
 Nitrogen N. 
 
 Oxygen O. 
 
 Phosphorus P. 
 
 Platmum Pt. 
 
 Potassium K. 
 
 Silver Ag. 
 
 Sodium Na. 
 
 Strontium Sr. 
 
 Sulphur S. 
 
 Tin Sn. 
 
 Zinc Zn. 
 
 27.1 
 120.2 
 
 74.96 
 137-37 
 208.0 
 
 II .0 
 
 79.92 
 II 2 , 40 
 
 40.09 
 
 12.0 
 
 35-46 
 
 52.0 
 
 58-97 
 63-57 
 19.0 
 
 197.2 
 1 .008 
 
 126.92 
 
 55-82 
 207 . 10 
 
 6-94 
 24.32 
 
 '54.93 
 
 200.0 
 96.0 
 58.68 
 14.01 
 16.00 
 31-04 
 
 195-2 
 39-1 
 
 107,88 
 23.00 
 
 87-63 
 32.07 
 119. o 
 65-37 
 
 26 
 119 
 
 74 
 136 
 206 
 
 9 
 3 
 3 
 4 
 4 
 10.9 
 
 79-36 
 III. 6 
 39-8 
 II .91 
 35-18 
 51-7 
 58-56 
 63 .1 
 18.9 
 
 195-7 
 1 .000 
 
 125-9 
 
 55-5 
 205-35 
 6.98 
 
 24.18 
 
 54-6 
 198-5 
 
 95-3 
 
 58-3 
 
 13-93 
 
 15.88 
 
 30.77 
 193-3 
 
 38.86 
 107. 12 
 
 22.88 
 
 86.94 
 
 31-83 
 118. 1 
 
 64.9 
 
 27 ,0 
 
 120.0 
 
 75-0 
 
 136.0 
 
 206.0 
 
 II .0 
 
 80.0 
 
 III .0 
 
 40.0 
 
 12.0 
 
 35-5 
 52.0 
 58.0 
 63.0 
 19.0 
 
 196,0 
 1 .0 
 
 126.0 
 56.0 
 
 206 
 7.0 
 24.0 
 
 55-0 
 200.0 
 
 95 o 
 58-0 
 J4,o 
 16.0 
 31.0 
 
 194.0 
 39-0 
 
 107.0 
 23.0 
 87.0 
 32.0 
 
 118. o 
 65.0 
 
 ' Internationa! Atomic Weights, 191 1 
 
TABLE OF MULTIPLES OF SOME ATOMIC WEIGHTS AND 
 COMBINATIONS IN FREQUENT USE. 
 
 
 I 
 
 2 
 
 3 
 
 4 
 
 5 
 
 6 
 
 7 
 
 8 
 
 9 
 
 H 
 
 1.008 
 
 -^ .016 
 
 3-024 
 
 4.032 
 
 5 -040 
 
 6.048 
 
 7.056 
 
 8.064 
 
 9-072 
 
 O 
 
 16.000 
 
 32. 
 
 48. 
 
 64. 
 
 80. 
 
 96. 
 
 112. 
 
 128. 
 
 144- 
 
 OH 
 
 17.008 
 
 34016 
 
 51.024 
 
 68.032 
 
 85.04 
 
 102.048 
 
 119-056 
 
 136.064 
 
 153-072 
 
 HsO 
 
 18.016 
 
 36.032 
 
 54.048 
 
 72.064 
 
 90.08 
 
 108.096 
 
 126. 112 
 
 144. 12S 
 
 162 . 144 
 
 N 
 
 14.01 
 
 28.02 
 
 42 -03 
 
 56.04 
 
 70.05 
 
 84.06 
 
 98.07 
 
 112 .08 
 
 126 .09 
 
 NHj 
 
 17-034 
 
 .1 4 . 064 
 
 51.102 
 
 68.128 
 
 85.170 
 
 102. 204 
 
 119.238 
 
 136.272 
 
 153-306 
 
 NH, 
 
 18.042 
 
 36.084 
 
 54-126 
 
 72.168 
 
 90.210 
 
 108.252 
 
 126 . 294 
 
 144-336 
 
 162.378 
 
 NO- 
 
 46.01 
 
 92.02 
 
 138.03 
 
 184.04 
 
 230.05 
 
 276.06 
 
 JI2.O7 
 
 388-08 
 
 414.09 
 
 NO3 
 
 62.01 
 
 124.02 
 
 186.03 
 
 248.04 
 
 310.05 
 
 372-06 
 
 434-07 
 
 496.08 
 
 558.09 
 
 C 
 
 12.00 
 
 24- 
 
 36. 
 
 48. 
 
 60. 
 
 72. 
 
 84- 
 
 96. 
 
 108. 
 
 CO2 
 
 44.00 
 
 88. 
 
 ■ 32. 
 
 176. 
 
 220. 
 
 264. 
 
 308. 
 
 352. 
 
 396. 
 
 CO3 
 
 60.00 
 
 120. 
 
 180. 
 
 240. 
 
 300. 
 
 360. 
 
 420. 
 
 480. 
 
 540. 
 
 CN 
 
 26.01 
 
 52.02 
 
 78.03 
 
 104.04 
 
 130. OS 
 
 156.06 
 
 182-07 
 
 208-08 
 
 234-09 
 
 CI 
 
 35.46 
 
 70.92 
 
 106.38 
 
 141 .84 
 
 177-3 
 
 212 . 76 
 
 248-22 
 
 283.68 
 
 319-14 
 
 Br 
 
 79.92 
 
 159.84 
 
 239.76 
 
 319-68 
 
 399-6 
 
 479-52 
 
 559-44 
 
 639-36 
 
 719.28 
 
 I 
 
 126.92 
 
 253-84 
 
 380.76 
 
 507.68 
 
 634-6 
 
 761 -52 
 
 888.44 
 
 1015.36 
 
 1 142 . 28 
 
 S 
 
 32.07 
 
 64.14 
 
 96.21 
 
 128.28 
 
 160.35 
 
 192.42 
 
 224.49 
 
 256.56 
 
 288-63 
 
 SO3 
 
 80.07 
 
 160.14 
 
 240.21 
 
 320.28 
 
 400 . 3 5 
 
 480. 42 
 
 560.49 
 
 640.56 
 
 720.63 
 
 SOi 
 
 96.07 
 
 192.14 
 
 288.21 
 
 384-28 
 
 480.35 
 
 586.42 
 
 672.49 
 
 768,36 
 
 864.63 
 
 PO. 
 
 9S04 
 
 190.08 
 
 285.12 
 
 380.16 
 
 475-20 
 
 580.24 
 
 665.28 
 
 760.32 
 
 855-36 
 
 P20S 
 
 142.08 
 
 284. 16 
 
 426.24 
 
 568.32 
 
 710.40 
 
 852-48 
 
 994-56 
 
 1136.64 
 
 1278.72 
 
ABBREVIATIONS AND SIGNS 
 
 cc cubic centimeter 
 
 gm gramme, 15:4323s grains 
 
 gr grain 
 
 at. wt atomic weight 
 
 V.S volumetric solution 
 
 T.S .test solution. 
 
 U. S. P United States Pharmacopoeia. 
 
 — normal. 
 
 I 
 
 — decinormal. 
 
 10 
 
 N 
 
 — centmormal. 
 
 100 
 
 — semi-normal. 
 
 2 
 
 2 
 
 T^ or 2N double-normal. 
 
THE ESSENTIALS OF VOLUMETRIC 
 ANALYSIS 
 
 CHAPTER I 
 
 INTRODUCTION 
 
 In a chemical analysis the aim is to determine the nature 
 of the chemical substances contained in a given compound 
 or to ascertain their quantities. In the former case the 
 analysis is a qtmlilative, in the latter a quantitive, one. 
 
 The quantitive analysis of a substance may be made 
 either by the gravimetric or the volumetric method. 
 
 The Gravimetric Method consists in separating and 
 weighing the constituents either in their natural states or in 
 the form of new and definite compounds, the composition 
 of which is known to the analyst. From the weights of 
 these new compounds the analyst can calculate the quantities 
 of the original constituents. 
 
 Example. To determine the quantity of silver in a solu- 
 tion by the gravimetric method we proceed as follows : 
 
 Ten grams of a solution containing silver in the form of 
 silver nitrate (AgNOs) is placed into a beaker, and, after 
 slightly acidulating with nitric acid, is treated with hydro- 
 chloric acid, drop by drop, until no further precipitation 
 occurs. The precipitate which consists of silver chlorid (AgCl) 
 is then separated by filtration, thoroughly washed, dried and 
 weighed. Its weight is found to be 0.69 gm. The calcu- 
 lation is then made as follows: 143.34 gms. of silver chlorid 
 
2 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 represents 107.88 gms. of silver or 169.89 gms. of silver nitrate, 
 as the equation shows: 
 
 AgN03+HCl = AgCl+HN03. 
 
 169.89 143-34 
 
 Therefore, 0.69 gm. of silver chlorid will represent 
 107. 
 
 143-34 
 
 Xo.69 = o.5i9 gm. of silver, 
 
 169.89 , „ , ., . 
 
 or Xo.69=o.8i7 gm. of silver nitrate. 
 
 143-34 
 
 The Volumetric Method. This method depends upon the 
 use of solutions {standard solutions) which are of khoAvn strength 
 and paying attention to the valume of such a solution which 
 must be added to the substance under analysis to perform 
 with it and complete a certain reaction. Thus, if we conduct 
 an analysis by means of such a solution, and can express by 
 a chemical equation the reaction which takes place, we can 
 readily and accurately calculate the quantity present of the 
 substance under analysis. 
 
 Example. If a silver solution is to be analyzed by this 
 method it is treated with a standard solution of sodium chlorid, 
 added slowly from a burrette until no more silver chlorid is 
 precipitated. Each cc. of this standard solution will precip- 
 itate a certain weight of silver as silver chlorid, and hence 
 by noting the number of cc. used to complete the precipitation, 
 the weight of the silver in the solution analyzed is easily 
 ascertained. 
 
 N 
 The — sodium chlorid solution is generally used for this 
 
 purpose. It is made by dissolving -^ of the molecular weight 
 of the salt (in grams) (5.846 gms.) in sufficient water to 
 make 1000 cc. 1000 cc. of this solution will precipitate -^ 
 of the atomic weight of silver (in grams) (10.788 gms.), and 
 
INTRODUCTION 3 
 
 hence each cc. of the sodium chlorid solution represents 
 
 0.010788 gm. of metallic silver, and by multiplying this figure 
 
 by the number of cc. of the sodium chlorid solution used, 
 
 the quantity of silver in the solution under analysis is ascer- 
 
 N 
 tained. If in the above analysis 100 cc. of the — sodium 
 
 10 
 
 chlorid solution were used, then 0.010788X100=1.0788 gms. 
 
 of metallic silver. 
 
 The reaction is illustrated by this equation: 
 
 AgNOs + NaCUAgCl + NaNOg. 
 10)107.88 10)58.46 
 
 1000) 10.788 gms. 1000) 5.846 gms. = 1000 cc. — V.S. 
 — 10 
 
 0.010788 gm. 0.005846 gm. = I cc. " " 
 
 From the examples given it will be seen that the gravi- 
 metric operations consume no little time, and require the 
 exercise of considerable skill. The washing of the precipitate 
 must be thoroughly performed in order that it be freed from 
 all adhering matter. The drying also is a matter of some 
 consequence and must be performed in such a manner as 
 to prevent the admixture of dust or the decomposition of 
 the precipitate by excessive heat. A very accurate balance 
 is also required. 
 
 The volumetric operations, on the other hand, do not 
 require that the substance to be determined be separated in 
 the' form of a compound of known composition and weighed 
 in the drj' state; in fact, the substance may be accurately 
 estimated when mixed with many others. It therefore obviates 
 the necessity for the frequent separations and weighings which 
 the gravimetric method demands, and enables the analyst to 
 do the work in a very short time. 
 
 The instruments needed for volumetric work are few and 
 simple, and comparatively little skill is required. Further- 
 more, the results obtained are in most instances more accurate. 
 
CHAPTER II 
 
 GENERAL PRINCIPLES OF CHEMICAL COMBINATION 
 UPON WHICH VOLUMETRIC ANALYSIS IS BASED 
 
 I. When substances unite chemically the union always 
 takes place in definite and invariable proportions. Thus when 
 silver nitrate and sodium chlorid are brought together, 169.89 
 parts (by weight) of silver nitrate and 58.46 parts (by weight) 
 of sodium chlorid will react with each other, producing 143.34 
 parts of a curdy white precipitate (silver chlorid). 
 
 These substances will react with each other in these pro- 
 portions only. 
 
 If a greater proportion of silver nitrate than that above 
 stated be added to the sodium chlorid, only the above pro- 
 portion will react, the excess remaining unchanged. 
 
 The same is true if sodium chlorid be added in excess 
 of the above proportions. For instance, if 200 parts of silver 
 nitrate be mixed with 58.46 parts of sodium chlorid, 169.89 
 parts only will react with the sodium chlorid, while 30.11 
 parts of silver nitrate will remain unchanged. Again, when 
 potassium hydroxid and sulphuric acid are mixed potassium 
 sulphate is formed, 11 2. 2 parts of potassium hydroxid and 
 98.1 parts of sulphuric acid being required for complete 
 neutralization. These two substances unite chemically m these 
 proportions only. 
 
 The equation is 
 
 2KOH-l-H2S04=K2S04-F2H20. 
 
 112. 2 yS.I 
 
GENERAL PRINCIPLES OF CHEMICAL COMBINATION 5 
 
 In other words, 112. 2 parts of KOH will neutralize 98.1 
 parts of H2SO4, and consequently 98.1 parts of H2SO4 will 
 neutralize 112. 2 parts of KOH. 
 
 Oxalic acid and sodium carbonate react upon each other 
 in the proportions shown in the equation 
 
 H2C2O4 • 2H2O + NaaCOs = Na2C204 + CO2 + 3H2O. 
 
 126.05 ^°6 
 
 126.05 parts of crystallized oxalic acid are neutralized by 
 106 parts of anhydrous sodium carbonate. 
 
 2. Definite chemical compounds always contain the same 
 elements in exactly the same proportions, the proportions 
 being those of their atomic weights, or some multiple of these 
 weights. 
 
 Thus sodium chlorid (NaCl) contains 23 parts of metallic 
 sodium and 35.46 parts of chlorin, these being the atomic 
 weights of sodiimi and chlorin respectively. 
 
 Potassiiun sulphate (K2SO4) contains twice 39.1 = 78.2 
 parts of potassium, 32.07 parts of sulphur, and four times 
 16 = 64 parts of oxygen. 
 
 Potassium hydroxid (KOH) contains 39.1 parts of potas- 
 sium, 16 parts of oxygen, and one part of hydrogen. Hydro- 
 chloric acid (HCl) contains one part of hydrogen and 35.46 
 parts of chlorin. 
 
 Upon these facts the volumetric methods of analysis are 
 based. 
 
 It has been shown that 98.1 gms. of sulphuric acid will 
 neutralize 11 2. 2 gms. of potassium hydroxid. It is therefore 
 evident if a solution of sulphuric acid be made containing 
 49.05 gms. of the pure acid in 1000 cc, that one cc. of this 
 solution will neutralize 0.0561 gm. of potassium hydroxid. 
 In estimating alkalies with this acid solution, the latter is 
 
6 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 added from a burrette, in small portions, until the alkali is 
 neutralized, as shown by its reaction with some indicator. 
 
 Each cc. of the acid solution required before neutralization 
 is complete indicates 0.0561 gm. of KOH, and the number 
 of cc. used multiplied by 0.0561 gm. gives the quantity of 
 pure KOH in the sample analyzed. 
 
 One cc. of the same solution will neutralize 0.04 gm. of 
 sodium hydroxid (NaOH), 0.053 gm. of anhydrous sodium 
 carbonate (Na2C03), etc. 
 
 If a solution of crystallized oxalic acid be made by dis- 
 solving 63.02 gms. in sufficient water to make 1000 cc, we 
 \vill have a normal solution, the neutralizing power of which 
 is exactly equivalent to the above mentioned normal sulphuric 
 acid solution. 
 
 The strength of acids is estimated by alkali volumetric 
 solutions. A normal solution of potassium hydroxid containing 
 56.1 gms. in the liter will neutralize exactly i liter of the 
 normal acid solution; i cc. of this normal alkali will neutralize 
 0.03646 gm. of HCl, 0.06362 gm. of H2C204,or 0.04905 gm. 
 of H2SO4, etc. 
 
CHAPTER III 
 VOLUMETRIC OR STANDARD SOLUTIONS 
 
 Any solution employed in volumetric analysis for the 
 purpose of estimating the strength of substances, that is, any 
 solution the chemical power or titer of which has been deter- 
 mined, is designated a standard or volumetric solution. Such 
 a solution is said to be " titrated " (French titre = title or 
 power), and is sometimes also called a set solution or a stand- 
 ardized solution. It may be normal, decinormal, empirical, 
 or of any strength, so long as its strength is known. 
 
 When volumetric analysis first came into use the solutions 
 were so made that each substance to be estimated had its 
 own special volumetric solution, and this was usually of such 
 strength as to give the result in percentages. Thus a certain 
 strength of standard acid was employed for potash, another 
 for soda, and a third for ammonia, and in testing the acids, 
 each had its own special standard alkali. These solutions 
 were known as normal solutions; they are still to some extent 
 in use, and since solutions now designated as normal are of 
 an entirely different character, it is important that no miscon- 
 conception should exist when a normal solution is spoken of. 
 
 Normal Solutions are such as contain one liter (looo cc), 
 the molecular weight of the active reagent expressed in grams, 
 and reduced to the valency corresponding to one atom of 
 replacable hydrogen or its equivalent. 
 
 Thus in univalent or monobasic compounds the full molec- 
 ular weight in grams is contained in a liter of the normal 
 
 solution. 
 
 7 
 
S THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Example. Hydrochloric acid, HCl, having one replacable 
 hydrogen atom, its normal solution would contain the full 
 molecular weight, 36.46 gms. in a liter. A normal solution 
 of potassium hydroxid should contain 56.1 gms. of KOH in 
 a Hter, while that of sodium hydroxid should contain 40 gms. 
 of absolute NaOH. 
 
 Normal solutions of bivalent or dibasic compounds, contain 
 in 1000 cc. one-half the molecular weight in grams. Thus, 
 oxalic acid H2C204-|-2H20 = 126.05, having two replacable 11 
 atoms, one-half of its molecular weight in grams = 63.02 is 
 contained in a liter of its normal solution. For the same 
 reason a liter of a normal solution of sulphuric acid contains 
 
 —— = 49.05 gms., and a liter of a normal solution of sodium 
 
 106 
 carbonate Na2C03 contams = 53 gms., while in the case 
 
 of trivalent or tribasic compounds one-third of the molecular 
 weight in grams is contained in a liter of the normal solution. 
 Thus it will be seen that one cc. of any normal acid solution 
 will neutralize one cc. of any normal alkali solution, because 
 one molecule of a univalent acid will neutralize one molecule 
 of a univalent alkali, or a half a molecule of a bivalent alkali. 
 This is shown by the equations 
 
 HCl f NaOH = NaCl-FH20, 
 
 36.46 46 
 
 2HCI + Na2C03 = 2NaCl + H2O + CO2. 
 2)72.92 2)106 
 
 36-46 53 
 The value of a reagent as expressed by its hydrogen 
 equivalent is readily seen in the case of acids and alkalies 
 by reference to the chemical formula, but in such standard 
 solutions as potassium dichromate, potassium permanganate, 
 sodium thiosulphate, and others, the particular reaction in 
 any given analysis must be taken into account in making a 
 
VOLUMETRIC OR STANDARD SOLUTIONS 9 
 
 normal solution; for instance, when K2Cr207 is to be used 
 as a precipitating agent, its reaction is as follows ; 
 
 2Ba(C2H302) 2 + KaCraOy + H2O = 2BaCr04 + 2KC2H3O2 
 
 + 2HC2H3O2. 
 
 It is thus seen that one molecule of K2Cr207 will cause 
 
 the precipitation of two atoms of barium in the form of 
 
 chromate. Each atom of barium is chemically equivalent to 
 
 two atoms of hydrogen; therefore one-fourth of a molecule 
 
 of K2Cr207 is equivalent to one atom of hydrogen. And 
 
 therefore a normal solution of this salt, when used as a 
 
 precipitating agent, must contain in one liter one-fourth of 
 
 294.2 
 its molecular weight in grams; = 73-55 g^is. 
 
 If K2Cr207 is to be used as an oxidizing agent, the three 
 atoms of oxygen which it yields for oxidizing purposes must 
 be taken into account. When this salt oxidizes it splits up 
 into K20-|-Cr203-l-03. The tliree atoms of oxygen combine 
 with and oxidize the salt acted upon, or they combine with 
 an equivalent quantity of the hydrogen of an acid and liberate 
 the acidulous part, which then combines with the salt. As 
 the equations show, 
 
 6FeO-fK2Cr207 = K20-hCr203+3Fe203 or (FceOg); 
 
 6FeS04 + K2Cr207 + 7H2SO4 = 
 
 7H20 + K2S04 + Cr2(S04)3+3Fe2(S04)3; 
 
 7H2SO4 + K2Cr207 = 3804 + 7H2O + K2SO4 + Cr2(S04) 3. 
 
 Each of these atoms of oxygen are equivalent to two atoms 
 of hydrogen. Thus O3 is equivalent to He. 
 
 Hence a liter of a normal solution of K2Cr207, when used 
 as an oxidizing agent, contains one-sixth of its molecular 
 weight in grams. 
 
10 THE ESSENTIALS OF VOLUAFETRIC ANALYSIS 
 
 The same may be said of potassium permanganate when 
 used as an oxidizing agent. 
 
 2KMn04 has five atoms of oxygen which are available 
 for oxidizing purposes, and each of these is capable of taking 
 two atoms of hydrogen from an acid and liberating the acidulous 
 part. The hydrogen equivalent of this salt may therefore be 
 said to be one-tenth of the weight of 2KMn04, and a normal 
 solution of this salt contains 31.606 gms. in a liter. 
 
 Sodium Thiosulphate (Hyposulphite), Na2S203, is another 
 instance. The molecule of this salt has two atoms of sodium, 
 which have replaced two atoms of hydrogen of thiosulphuric 
 acid. Thus it would seem that a normal solution -should 
 contain one-half of the molecular weight in grams. But the 
 particular reaction of this salt with iodin is taken into account. 
 
 One molecule reacts with one atom of iodin, as seen in 
 the equation 
 
 2Na2S203 + h = sNal + Na2S406. 
 
 Since iodin is univalent, a molecule of the salt is equivalent 
 to one atom of hydrogen. 
 
 A normal solution of this salt therefore contains the 
 
 molecular weight in grams in a liter. 
 
 N 
 Decinormal Solutions, — , are one-tenth the strength of 
 10 ° 
 
 normal solutions. 
 
 N 
 Centinormal Solutions, — , are one-hundredth the strength 
 
 of normal solutions. 
 
 Seminormal Solutions, — , are one-half the strength of 
 
 N 
 Solutions, 
 
 normal solutions. 
 
 2 
 Double-normal Solutions, ^, are twice the strength of 
 
 the normal. 
 
VOLUMETRIC OR STANDARD SOLUTIONS 11 
 
 Empirical Solutions. A solution which does not contain 
 an exact atomic proportion of the reagent may be employed 
 as a volumetric solution after its strength or titer has been 
 determined. Such a solution is said to be empirical, and 
 solutions of this sort are very frequently used. To prepare 
 solutions of exactly normal strength is a tedious process and 
 often inconvenient. If the solution is approximately normal 
 and its strength accurately determined, it may be used as it 
 is. Agam, in the case of standard solutions of the caustic 
 alkalies, which, when not kept with all precautions, deteriorate 
 readily through absorption of carbon dioxid from the air, as well 
 as through their action upon the glass containers. To restore 
 the titer of such solutions by the introduction of more of the 
 alkali is an unnecessary waste of time, inasmuch as it is only 
 necessary to determine its exact strength and then use it as 
 it is. For instance, if an approximately normal solution of 
 potassium hydroxid is on hand, its strength is determined as 
 follows: 
 
 Ten cc. of an exactly normal oxalic or other acid solution 
 are put into a beaker, and after diluting with a little water, 
 and adding three or four drops of methyl orange, the empirical 
 potassium hydroxid solution is run in from a burette until 
 the color of the solution changes from red to yellow; the 
 number of cc. required is then noted. 
 
 Assuming that 10.4 cc. were required to neutralize the 
 
 100 
 10 cc. of normal acid, hence its strength is — or 0.9615, 
 
 that of a strictly normal solution, and the number of cc. used 
 
 100 
 in any estimation must be multiplied by — or 0.9615, 
 
 and then by the normal factor for the substance analyzed. 
 
 It is a good plan to have the factor marked on the label 
 of the bottle containing such an empirical solution. In this 
 case it would be Xo.96i5=normal. 
 
12 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Standard solutions for use in volumetric analysis are usually 
 solutions of acids, bases, or salts, and in two cases elements, 
 namely, iodin and bromin. 
 
 A standard solution of a base is usually used for the esti- 
 mation of free acids. 
 
 A standard solution of an acid is usually used for the 
 estimation of a free base, or the basic part of a salt, the acid 
 of which can be completely expelled by the acid used in the 
 sta,ndard solution. Example, carbonates. 
 
 A standard solution of a salt may be used as a precipitant, 
 or it may be used as an oxidizing or reducing agent. 
 
 That part of the reagent in a standard solution which 
 reacts with the substance under analysis is the active con- 
 stituent of the solution. As Ag in AgNOs is the active con- 
 stituent of the standard solution of silver nitrate, 
 
 AgNOs -hNaCl = AgCl-FNaN03, 
 
 or CI in NaCl, is the active constituent of the standard solution 
 of sodium chlorid. 
 
 If the reagent is a base, as KOH, the basic part K is the 
 active constituent. If the reagent is an acid, the active constit- 
 uent is the acidulous part, as SO4 in H2SO4. 
 
 If the action of the reagent is oxidizing, then that part 
 of the reagent which produces the oxidation is the active 
 constituent. 
 
 The valence of an acid is shown by the number of replace- 
 able hydrogen atoms it contains. Thus HCl is univalent, 
 H2SO4 is bivalent, which means that a molecule of HCl is 
 chemically equivalent to one atom of hydrogen, and a molecule 
 of H2SO4 is chemically equivalent to two atoms of hydrogen. 
 
 The valence of a base is shown by the number of hydroxyls 
 it is combined with. As KOH is univalent, Ca(0H)2 is 
 bivalent. 
 
VOLUMETRIC OR STANDARD SOLUTIONS 13 
 
 The valence of a salt is shown by the equivalent of base 
 which has replaced the hydrogen of the corresponding acid. 
 
 Thus NaCl, in which Na has replaced H of HCl, is uni- 
 valent. 
 
 K2SO4, in which K2 has replaced H2 of H2SO4, is bivalent. 
 
 Preparation of Volumetric Solutions. In preparing volu- 
 metric solutions it must be borne in mind that most salts when 
 dissolved in water cause a condensation in volume, through 
 reduction of temperature, while some substances, as for instance 
 sulphuric acid and alkali hydroxids, cause a rise in tempera- 
 ture and a consequent increase in volume. It is therefore 
 necessary, after making a solution, to set it aside for a short 
 time, in order to allow it to attain the proper temperature 
 before measuring it. 
 
 It is always the best plan to take a weighed quantity of 
 the salt, slightly greater than that required by theory, and 
 to dissolve it in less water than is needed for the finished 
 solution. This solution is titrated, its strength determined 
 and then diluted to the proper measure. 
 
 After dilution it should be again carefully titrated and 
 its titer verified. 
 
 All volumetric solutions should be made with distilled 
 water and with reagents which are of a high degree of purity. 
 
 Standard Temperature. A cubic centimeter is the volume 
 occupied by one gram of distilled water at its maxim um_ density 
 4° C. (39° F.). This, however, is not the cubic centimeter 
 used in volumetric analysis. It is convenient to use in analyses 
 of this sort a cubic centimeter which represents the volume of 
 one gram of distilled water at a temperature which is easily 
 attained and maintained at any season of the year. 
 
 The temperature at which measuring instruments are 
 graduated is the temperature at which volumetric solutions 
 should be prepared, and at which all volumetric operations 
 
14 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 should be conducted. 4° C. is a temperature at which it is 
 obviously impossible to work except during one or two months 
 of the year. For this reason the temperature of 16° C. (60.8° F.) 
 has been taken as the standard. The cubic centimeter used 
 in volumetric analysis, under this standard, is the volume 
 occupied by one gram of distilled water at the latter tempera- 
 ture. 
 
 The employment of this standard of temperature, though 
 long in vogue, is justly criticized as too low. In order to obtain 
 accurate results the temperature of the atmosphere in which 
 the titration is performed must not be too much at variance 
 with the temperature at which the instruments are graduated 
 and the solutions made. A temperature of 16° C. is one 
 which is exceedingly dif&cult to maintain in the warm months 
 of the year, therefore it has been suggested to take a higher 
 temperature as the standard. 
 
 The U.S.P. VIII recommends the employment of 25° C. 
 (77° F.) as the standard. This, while better than the lower 
 temperature, is regarded by many as being too high and the 
 use of 20° C. (68° F.) as the standard is now being very 
 favorably considered, this being nearer the average temperature 
 of the atmosphere in laboratories throughout the year. What 
 ever temperature is adopted, it is at this temperature that 
 the whole set of measuring instruments must be graduated, 
 and all titrations carried out. It would be obviously improper 
 to use a burette graduated at 16° C. and a flask or cylinder 
 graduated at 25° C, or to employ solutions at a temperature 
 which is different from that at which they are made. 
 
 To Titrate a substance means to test it volumetrically 
 for the amount of pure substance it contains. The term is 
 used in preference to " tested " or " analyzed," because these 
 terms may relate to qualitative examinations as well as quan- 
 titative, whereas titration applies only to volumetric analysis. 
 
VOLUMETRIC OR STANDARD SOLUTIONS 15 
 
 Residual Titration, Re-titration, sometimes called Back 
 Titration, consists in treating the substance under examina- 
 tion with standard solution in a quantity known to be in 
 excess of that actually required; the excess (or residue) is 
 then ascertained by residual titration with another standard 
 solution. 
 
 Thus the quantity of the first solution which went into 
 combination is found. 
 
 N 
 Example. Ammonium carbonate is treated first with — 
 
 I 
 
 H2SO4 in excess, and the excess then found by titration with 
 
 I 
 
 N 
 The quantity of the — KOH used is then deducted from 
 
 N 
 the quantity of — H2SO4 added, which gives the quantity of 
 
 the latter which was neutralized by the ammonium carbonate. 
 
 Titrations may be carried out in flasks, beakers, or in 
 white porcelain evaporating dishes. Flasks of the Erlenmeyer 
 pattern, see Fig. 24, ha^dng a short narrow neck and a broad 
 fiat bottom, are very desirable for this purpose; they admit 
 of shaking their contents without danger of loss, and permit 
 ready observation of color changes. If a flask is used the 
 tip of the burette should extend well into its neck in order 
 to prevent any loss of the reagent. The flask should be 
 rotated after each addition of the reagent, and when the end 
 of titration is near, any of the solution adhering to the sides 
 of the flask should be washed down with distilled water. 
 A white porcelain tile, or a sheet of white paper placed under 
 the flask or beaker, adds materially in the observation of the 
 color change. 
 
CHAPTER IV 
 INDICATORS* 
 
 In volumetric analysis the substance to be analyzed in 
 the state of solution is placed in a beaker and the standard 
 solution is added from a burette until a certain reaction is 
 produced. The exact moment when a sufficient quantity of 
 the standard solution has been added is known by certain 
 visible changes, which differ according to the substance analyzed 
 and the standard solution used. When such a visible change 
 occurs the " end-reaction " is reached. 
 
 The end-reaction manifests itself in various ways, as 
 follows : 
 
 1. Cessation of precipitation. 
 
 2. First appearance of a precipitate. 
 
 3. Change of color. 
 
 In some cases, however, the addition of the standard 
 solution to the substance under analysis does not produce 
 either a precipitate or a change of color; in such cases we 
 must resort to the use of an indicator. 
 
 An indicator is a substance which is used in volumetric 
 analysis, and which indicates by change of color, or some 
 other visible sign, the exact point at which a given reaction 
 is complete. 
 
 Generally the indicator is added to the substance under 
 examination, but in a few cases it is used alongside, a drop 
 of the substance being occasionally brought in contact with 
 a drop of the indicator. 
 
 * A more detailed description of the individual indicators is given in 
 the Appendix. 
 
 16 
 
INDICATORS 17 
 
 Thus in estimating an alkali with an acid volumetric solu- 
 tion the alkali is shown to be completely neutralized when 
 the litmus tincture which was added becomes faintly red or 
 the phenolphthalein colorless. Again, when haloid salts are 
 estimated with nitrate of silver solution, chroma te of potas- 
 sium is added as indicator. A white precipitate is produced 
 as long as any halogen is present to combine with the silver, 
 and when all is precipitated the chroraate of potassium acts 
 upon the silver nitrate, forming the red silver chromate, this 
 color thus showing that all the halogen has been precipitated. 
 
 INDICATORS COMMONLY USED 
 
 The principle indicators used are: 
 
 Tincture of Litmus, which shows acidity by turning red 
 and alkalinity by becoming blue. 
 
 Phenolphthalein Solution, which is colorless in acid solu- 
 tions and red in alkaline solutions, but is not reUable for 
 alkaline phosphates, bicarbonates or ammonia. 
 
 Methyl-orange Solution turns red with acids and yellow 
 with alkalies. It is not affected by carbonic acid, and is 
 therefore adapted for the titration of alkaline carbonates. 
 
 Rosolic-acid Solution is yellow with acids and violet-red 
 with alkalies. It is very sensitive to ammonia. 
 
 Tincture of Turmeric turns brown with alkalies, and the 
 yellow color is restored by acids. 
 
 Cochineal Solution turns violet with alkalies and yellowish 
 with acids. It is used chiefly in the presence of ammonia 
 or alkali earths. 
 
 Eosin Solution is red by transmitted light, and shows a 
 strong green fluorescence by reflected light. Acids destroy 
 this fluorescence and alkalies restore it. 
 
 Brazil-wood Test-solution turns purplish-red with alkalies 
 and yellow with acids. 
 
18 THE ESSENTIALS OF VOLUMETRIC ANAl-YSIS 
 
 Fluorescein Tect-solution shows a strong green fluorescence 
 by reflected light in the presence of the least excess of an 
 alkali. 
 
 Neutral Potassium-chromate Test-solution is used in the 
 titration of haloid salts with silver-nitrate solution. It indicates 
 that all the halogen has combined with the silver by producing 
 a red-colored precipitate (silver chroma te). 
 
 Potassium-ferricyanide Test-solution is used in the estima- 
 tion of ferrous salts with potassium-dichromate solution. It 
 gives a blue color to a drop of the solution on a white slab 
 as long as any iron salt is present which has not been oxidized 
 to ferric. 
 
 Many other indicators are also used. 
 
 THE IONIZATION OR DISSOCIATION THEORY 
 
 When a soluble salt dissolves in water, its molecules split 
 up or dissociate more or less completely into parts called 
 ions. This behavior of substances, on going into solution, 
 is known as electrolytic dissociation or ionization. 
 
 Ions are electrically charged atoms or groups of atoms 
 and are of similar composition to the substances formed from 
 the compound when an electric current is passed through the 
 solution. The electro-positive ions migrate to and collect 
 around the negative pole (cathode) and hence are called 
 cathions, while the electro-negative ions are called anions, 
 because they concentrate around the positive pole or anode. 
 The dissociation of a compound into its ions when an electric 
 current is passed through its solution, although spoken of 
 as electrolytic dissociation, is really not caused by the electric 
 current, since the dissociation into ions occurs at once 
 upon dissolving the substance in water and without the aid 
 of an electric current, the action of the current being the 
 transportation of the separated ions to the poles. 
 
INDICATORS 19 
 
 The dissociation of compounds into ions when dissolved 
 in water is illustrated in the following list: 
 
 Sodium chlorid into (Na + ) and (CI - ) 
 
 Potassium nitrate into (K+) 
 
 " hydroxid into (K + ) 
 
 " acetate into (K + ) 
 
 Sulphuric acid into (H + ) 
 
 or (H + ) 
 
 (NO3-) 
 
 (0H-) 
 
 (C2H3O2-) 
 
 (HSO4-) 
 
 (H + )and (SO4) 
 
 The extent of this dissociation depends upon the nature 
 of the substance and the degree of dilution; the greater the 
 dilution the more complete the dissociation. Furthermore, 
 strong acids and bases dissociate readily, even in compara- 
 tively concentrated solutions, while the weaker acids and bases 
 are more or less undissociated when dissolved, i.e., they are 
 not readily split up into ions. Their salts, however, are 
 immediately and completely ionized. Therefore, upon neu- 
 tralizing a weak acid or base, an ionizable salt is formed. 
 According to the theory of Arrhenius, the reactions of analytical 
 chemistry are chiefly reactions between ions and not between 
 atoms. 
 
 Strong acids, bases and salts exist in solution, not as 
 molecules but chiefly in the form of ions. The formation 
 of silxer chlorid by the reaction between silver nitrate and 
 sodium chlorid takes place according to the following equation : 
 
 Ag/NOgAq. -l-Na/ClAq. = AgCl( solid) -t-Na/NOgAq. 
 
 The state of dissociation being denoted by the vertical 
 
 line between the ions of the molecules. 
 
 + + - . 
 
 This theory also explains why K/CIO3 with Ag/NOs does 
 
 not form AgCl, in that the reaction involves the ion CIO3 
 and not the atom CI. 
 
20 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Theories of Indicators. In connection with the use of 
 indicators in neutralization analyses, the question as to the 
 cause of the color changes is one of considerable interest. 
 
 Two distinct views are held. Of these the Ionization 
 Theory of Osiwald has received almost universal acceptance. 
 According to this theory the color changes are ascribed to a 
 change in the indicator from a molecular to an ionic condition. 
 As exemplified in the case of phenolphthalein the colorless 
 molecule , 
 
 OCOC6H4C-(C6H40H)2 .... (I) 
 I 1 
 
 is dissociated into the red negative ion 
 
 OCOC6H4C(C6H40H)C6H40. . . . (II) 
 
 In the other and less known view on this subject (the 
 Chromophoric Theory),* the sensitiveness of the indicators 
 and its color changes is ascribed to a change in the consti- 
 tution of the indicator, involving a chromophoric group, under 
 the influence of hydrogen and hydroxyl ions. According to 
 this view the color change is due (in the case of phenol- 
 phthalein) to a change of constitution from the colorless lactoid 
 (I) with no chromophoric group, to the colored quinoid (III) 
 with a quinone chromophore. 
 
 (NaOOC-C6H4)(HOC6H4)C;C6H4:0, . . (Ill) 
 
 and that the ionization of the sodium salt is merely a coin- 
 cidence and not the cause of the color change. 
 
 Whichever of these views is the correct one, remains for 
 future investigations to prove. That of Ostwald, being most 
 generally accepted at the present time, is treated more fully 
 below. 
 
 * See Julius Stieglitz, Jour. Am, Chem. Soc , XXV, 1112 (1903). 
 
INDICATORS 21 
 
 The Ionization Theory of Indicators.* The indicators used 
 in allialimetry and acidimetry are compounds of feeble acid 
 or basic character, and hence not prone to dissociation in 
 solution, but when neutralized the salt formed ionizes the 
 instant of its formation; the ions so liberated give rise to 
 colors which differ from those of the original compounds. 
 
 Any feeble acid or base may be utilized as an indicator 
 if its ions have a color different from that of the un-ionized 
 compound. Strong acids or bases are not suited as indicators, 
 because they ionize while in a free state on dilution, and thus 
 give no color when neutralized. 
 
 A solution in which H ions predominate has an acid reac- 
 tion, while one in which OH ions predominate reacts alkaline. 
 
 Phenolphthalein is a feebly acid indicator, and in its 
 imdissociated state is colorless. It does not dissociate readily 
 unless neutralized, but when sodium hydroxid is added to 
 its solution, a sodium salt of phenolphthalein is formed which 
 immediately ionizes and the ions liberated impart to the solution 
 a briUiant red color. If now some acid is added the sodium 
 salt is decomposed and the acid phenolphthalein again set 
 free, and the solution becomes colorless. 
 
 If a few drops of phenolphthalein solution be added to 
 an acid solution, ionization of the former is prevented by 
 the presence of the stronger acid; if now some sodium hydroxid 
 solution is added, the OH ions of the latter unite with the 
 H ions of the acid, and when the acid is completely neutralized 
 the first drop of excess of alkali unites with the phenolphtha- 
 lein, forming a salt which immediately ionizes and produces 
 the characteristic red color which shows the end of the reaction. 
 
 In the titration of a feeble acid the end-point is often 
 
 * See Ostwald's ''Lehrbuch der Allgemeinen Chemie," i8gi, and "Scienlific 
 Foundations of Analytical Chemistry," 1900; also Walker's " Introduction to 
 Physical Chemistry," 1899. 
 
22 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 indistinct and is lacking in sharpness; this is because the 
 indicator used has a greater tendency to ionize than the acid 
 itself. In this case the H ions present just before the com- 
 pletion of the reaction are not in sufficient amount to fully 
 retard the ionization of the indicator, and hence the latter 
 dissociates partly before neutralization is complete and gives 
 rise to an indefinite end-reaction. Therefore it is necessary 
 when titrating a feeble acid that an indicator should be selected 
 whose alkali salt ionizes with the production of a distinct 
 color change, and whose tendency to ionize is less than that 
 of the acid. Phenolphthalein is a suitable indicator in this 
 case, provided a strong alkali be used for titrating. 
 
 Fixed alkalies readily yield ionizable salts with phenol- 
 phthalein, but ammonia does not. The latter being too weak 
 a base to yield with so feeble an acid, a salt which can with- 
 stand the hydrolytic action of the water in dilute solutions, 
 and as a consequence a larger excess of the ammonia must 
 be used to overcome this. Thus is accounted for the imperfect 
 color change of phenolphthalein when ammonia or its salts 
 are present and why the color becomes visible only after a 
 large excess of the alkali is added. 
 
 Paranitrophenol is also an acid indicator; it exists in 
 solution in the form of undissociated colorless molecules, yet 
 its electro-negative ion is intensely yellow in color. This com- 
 pound has a slight tendency to dissociate in dilute solutions, 
 but the presence of a trace of a stronger acid will overcome 
 this tendency and the solution remains colorless. If an alkali 
 is, however, added, a salt of paranitrophenol is formed which 
 immediately ionizes and exhibits the intense yellow color of 
 its liberated ion. 
 
 Other indicators exhibit a color in both the ionized and 
 the non-ionized state, but the colors in both conditions are 
 different, as in the case of litmus, lacmoid and methyl orange. 
 
INDICATORS 23 
 
 Methyl orange is both an acid and a base and will form 
 ionizable salts with either acids or alkalies; its indicator 
 characteristics are, however, due essentially to its basic char- 
 acter. 
 
 The salt which methyl orange forms with acids dissociates 
 into red ions; this, upon the addition of an alkali, returns to 
 its undissociated state, which is yellow. Because of its weak 
 basic character its compound with acids is readily decomposed 
 by alkalies, but it takes a strong acid or a relatively large 
 quantity of a feeble acid to dissociate it in its non-ionized 
 state, hence this indicator is very sensitive to alkalies, and 
 much less so to acids. 
 
 With reference to the acid character of the indicator the 
 explanation of its action is that the non-ionized indicator is 
 red, while its ion is yellow. Acids lessen its ionization and 
 produce a red color, while alkalies produce a highly ionizable 
 salt and hence a yellow color. When a weak but slightly 
 ionizable acid is added to the methyl orange solution, the H' 
 ions of the acid given up in excess of the amount required for 
 neutralization are not sufficient to yield enough of a non- 
 ionizable salt to produce a decided red color, hence a large 
 quantity of such a weak acid is required to give an acid 
 indication. This would explain why methyl orange is not 
 suitable as an indicator for weak acids, and why it is very 
 sensitive to alkalies. 
 
 Litmus is an acid indicator which slightly dissociates in 
 solution. Its non-ionized molecules are red, but its negative 
 ions are blue. If it is added to an alkali, a salt is formed 
 which at once ionizes and gives a blue color. If added to 
 an acid, ionization is prevented and the red color of the non- 
 ionized molecules appears. 
 
 From the above explanations it will be seen that indicators 
 cannot be indiscriminately used, and that no one indicator 
 
24 THE ESSENTIALS OF VOI-UMETRIC ANALYSIS 
 
 will be suitable for every titration. Hence the indicator must 
 be selected to suit each case. This selection is facilitated by 
 reference to the classification of indicators, according to F. 
 Glaser, Ztchr. f. a. Chem., 1899, 273 + . 
 
 Group I. Indicators Forming Fairly Stable Salts. The 
 members of this group comprise such indicators as are (i) of 
 a strong acid character and which react readily with weak 
 bases, or (2) of a feeble, basic character and which require 
 a strong acid to form a stable salt. Hence they will be found 
 to be very sensitive to alkalies, and are useful in the titration 
 of weak bases, as ammonia and the amine bases, as well as 
 strong bases and acids. The indicators of this group are 
 the following, arranged in the order of their sensitiveness 
 towards alkalies: 
 
 (i) lodeosin, Resazurin; (2) Tropaeolin OO, Luteol; (3) 
 Methyl and Ethyl Orange; (4) Congo Red; (5) Cochineal; 
 (6) Lacmoid. 
 
 Group II. Indicators Possessing Faint Acid Properties and 
 Yielding Salts which are Very Unstable. These are readily 
 decomposed by relatively feeble acids, and are in consequence 
 very sensitive towards acids, slightly so towards alkalies. 
 
 They are: (i) Rosolic acid; (2) Curcuma; (3) Phenol- 
 phthalein, Flavescin; (4) Alpha-naphtholbenzein. 
 
 Group III. Indicators Occupying a Place Midway between 
 the Other Two Groups. They are somewhat stronger acids 
 than those of Group II, but feebler than those of Group I. 
 
 They are fairly sensitive towards both acids and alkalies, 
 but are more sensitive towards acids than those of Group I, 
 and less so towards alkalies. They are : 
 
 (i) Fluorescein, Phenacetolin; (2) Haematoxylin, Gallein, 
 Alizarin; (3) Litmus; (4) Paranitrophenol. 
 
 This division of indicators into groups, as above, facilitates 
 the selection of an indicator suitable for the work in hand. 
 
INDICATORS 25 
 
 For instance, for titrating weak acids, a glance at the groups 
 will show that the members of Group II are best adapted 
 for this purpose. Again, weak bases will be best titrated by 
 the indicators of Group I. Strong acids or bases may be 
 titrated by means of any of the indicators. 
 
 The quantity of indicator taken in a titration is a matter 
 of considerable moment. The smallest quantity which will 
 produce a distinct color should be taken, but it is equally 
 important that the quantity be not too small for the volume 
 of liquid; for in high dilutions the hydrolytic action of water 
 asserts itself, and intermediate tints will result, which interfere 
 with the sharpness of the end color. 
 
 If too much of the indicator is used, the sensitiveness is 
 
 lessened, because acid or alkali must be added to convert the 
 
 indicator into a salt, or when formed to decompose it; i.e., 
 
 a minimum of excess of the titrating fluid would react with 
 
 a small portion only of the indicator and intermediate tints 
 
 would be produced, until sufficient of the titrating solution 
 
 has been added to neutralize all of the indicator present. This 
 
 is especially true when using centinormal solutions. 20 drops 
 
 of litmus added to 10 cc. of water require from 10 to 14 drops 
 
 N 
 
 of acid or alkali solution to produce a change of color. 
 
 100 
 
 Thus the indicator itself takes up some of the standard 
 solution, and hence the necessity for using as small a quantity 
 of the indicator as possible; usually from j to ^ drops of the 
 indicator may be taken to each 50 or 100 cc. of the fluid titrated. 
 
 The degree of dilution of the substance titrated is also a 
 matter of considerable moment. In very concentrated solu- 
 tions ionization does not so readily occur, while too great a 
 dilution diminishes the reactive ability of the ions because 
 of their greater separation, and also because of the hydro- 
 lytic dissociation of water itself into H" and OH ions which 
 
26 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 react acid and alkaline respectively, and which brings about 
 a premature dissociation of the indicator. 
 
 The Requirements of a Good Indicator, according to H. 
 A. Cripps, are: 
 
 I. The end-reaction should be marked by a prominent 
 change of color. 
 
 II. The smallest possible quantity of the reagent should 
 be required to effect this change. 
 
 III. High tintorial power, which of itself assists in the 
 fulfilment of the second requirement, less of the indicator 
 being required. 
 
 IV. The change of color should not be affected by the 
 impurities commonly present in the substance under examina- 
 tion, nor by the products of the reaction. 
 
 In addition to these requirements it is a distinct advantage 
 if the color reaction is equally decided in alcoholic as in aque- 
 ous liquids. 
 
 A GUIDE FOR THE SELECTION OF INDICATORS 
 
 For Hydroxids and Carbonates 
 
 Indicators not affected by COj Indicators affected by CO2 
 
 (Cold Titrations) (Hot Titrations) 
 
 Methyl Orange, Gallein, Phena- Phenolphthalein (useless in 
 
 cetolin, Congo Red, lodeosin, presence of NH3 or its salts), 
 
 Cochineal. Lacmoid, Rosolic Acid, Resazurin. 
 
 For Ammonia {NH3) For Ammonium Carbonate 
 
 Same in 
 also Phe: 
 phthalein. 
 
 Rosolic Acid, Methyl Orange, Same indicators as for ammonia, 
 
 Congo Red, Litmus, Gallein. also Phenacetalin and Phenol- 
 
INDICATORS 27 
 
 For Inorganic Acids For Organic Acids 
 
 H^0„ HCl, HNO3. Phenolphthalein (all) Rosolic 
 
 Phenolphthalein, Litmus, Rosolic Acid (except acetic, citric and 
 Acid, Methyl Orange. tartaric), Galle'in. 
 
 Phenolphthalein (neutralized to 
 NajHPO,). 
 
 Methyl Orange and Cochineal 
 (each neutralized to NaHzPO,). 
 H,SO,. 
 
 Rosolic Acid and Methyl Orange. 
 
 Phenolphthalein (after addition 
 of glycerin). Litmus and Turmeric 
 paper. 
 
CHAPTER V 
 
 APPARATUS USED IN VOLUMETRIC ANALYSES 
 
 The Burette is a graduated glass tube which holds from 
 25 to 100 cc. and is graduated in fifths or tenths of a cc, 
 
 and provided at the lower end with 
 a rubber tube and pinch-cock. The 
 use of this instrument is to accurate- 
 ly measure quantities of standard 
 solutions used in an analysis. It is 
 in an upright position when in use, 
 and the flow of the solution can be 
 regulated so as to run out in a 
 stream or flow in drops by pressing 
 the pinch-cock between the thumb 
 and forefinger. The quantity of 
 solution used can be read from the 
 graduation on the outside of the 
 tube. This is the simplest and most 
 common form of burette, and is 
 known as Mohr's (Fig. i). 
 
 The use of the pinch-cock in 
 Mohr's burette may be dispensed 
 with by introducing into the rubber 
 tube a small piece of glass rod, 
 which must not fit too tightly. By 
 firmly squeezing the rubber tube surrounding the glass rod a 
 small canal is opened, through which the liquid escapes. A 
 
 28 
 
 Fig. I. 
 
APPARATUS USKD IN VOLUMETRIC ANALYSES 
 
 29 
 
 very delicate action can in this way be obtained, and the flow 
 of the liquid is completely under the control of the operator. 
 (See Fig. 2.) 
 
 The greatest drawback to this burette is that it cannot 
 be used for permanganate or other solu- 
 tions that act upon the rubber. 
 
 This defect can be overcome by the use 
 of a burette having a glass stop-cock in 
 place of the rubber tubing and pinch- 
 cock. This form has the additional advan- 
 tage of being capable of delivering the solution 
 in drops while both hands of the operator are 
 disengaged (Fig. 3). 
 
 Another good arrangement is that in 
 which the tap is placed in an oblique position, 
 so that it will not easily drop out of place 
 
 (Fig- 4)- 
 
 These glass stop-cock burettes should be 
 
 emptied and washed immedi- 
 ately after use, especially if soda 
 or potassa solution has been 
 used; for these act upon the 
 glass and often cause the stop- 
 per to stick so firmly that it 
 cannot be turned or removed 
 without danger of breaking 
 the instrument. 
 
 The most satisfactory form 
 of glass stop-cock is that shown 
 in Fig. 5. 
 
 When a number of estima- 
 tions are to be made in which the same volumetric solution is 
 employed, the arrangement shown in Fig. 6 is very serviceable. 
 
 48- 
 
 49- 
 
 FiQ. 3. 
 
 Fig. .\. 
 
30 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 A T-shaped glass tube is inserted between the lower end 
 of the burette and the pinch-cock and connected by a rub- 
 ber tube with a reservoir containing the volumetric solution. 
 
 The tube which communicates with the reservoir is provided 
 with a pinch-cock, which, when open, allows the solution to 
 flow into and fill the burette in so gradual a manner that 
 
 Fig. s- 
 
 Fig. 6. 
 
 The burette is emptied in the usual 
 
 no bubbles are formed, 
 manner. 
 
 E. &= A. Automatic Burette (Fig. 7). This is used for 
 the same purpose as the foregoing. It is provided with a 
 side tube for connection with reservoir, and has an overflow 
 cup which prevents its being filled to above the zero mark. 
 The three-way stop-cock is so arranged that if turned one 
 
APPARATUS USED IN VOLUAJETRIC ANALYSES 
 
 31 
 
 way the inlet is opened and the liquid from the reservoir flows 
 into and fills the burette. If turned the other way the inlet 
 is closed and the outlet is opened and the burette may be 
 
 Fio. 7. 
 
 Fig. 8. 
 
 emptied. If the handle of the stop-cock is turned half-way 
 round, both openings are closed. 
 
 There are many other forms of automatic burettes. 
 
 When working with solutions which are readily altered 
 by contact with air, as for example, stannous chlorid, potas- 
 
32 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 sium, sodium, or barium hydroxid or ammonia, an arrange- 
 ment like that depicted in Fig. 8 is very serviceable. In this 
 
 the upper end of the burette is 
 connected with the reservoir by 
 means of a rubber tube, thus 
 making an air-tight combination 
 between the burette and the reser- 
 voir. Its utility may be further 
 enhanced by providing the reser- 
 voir with a soda-lime tube or some 
 other suitable absorption tube. 
 
 Another form of apparatus is 
 shown in Fig. 9. In this both 
 the burette and the reservoir are 
 provided with tubes containing 
 soda-lime to insure a protection 
 against the admission of CO2 
 and moisture from the air. 
 
 Pinch-cocks used with Mohr's 
 burettes are of various kinds. 
 See Figs. 10, 11 and 12. 
 
 The screw pinch-cock,* Fig. 
 12, is a very useful device; it 
 may be used like the ordinary 
 pinch-cock by pressure with the 
 fingers upon a-a, when a rapid 
 flow is desired, or the nut-screw 
 {]}) may be so adjusted as to 
 allow a slower flow or to deliver 
 the solution in drops, thus giving the operator the freedom 
 of both hands for other work. 
 
 Burette supports are of various forms. One of the best 
 
 * W. V. Hergendorf. 
 
APPARATUS USED IN VOLUMETRIC ANALYSES 
 
 33 
 
 for one or two burettes is shown in Fig. 13. It is made of 
 iron, can stand firmly upon an uneven surface, and does not 
 easily tip over. The burettes are fastened to it by means 
 of clamps, illustrated in Figs. 14 and 15. 
 
 A revolving burette-holder for eight burettes is shown in 
 
 Fig. 16. Burrette-supports are also made with white porcelain 
 base, which enables the operator more readily to see the change 
 of color in the liquid titrated. 
 
 Pipettes are of two kinds — those which are marked to 
 deliver one quantity only, and those which are graduated on 
 the stem like burettes. Their use is to measure out portions 
 of solutions with exactness. 
 
34 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Pipettes are filled by applying the mouth to the upper end 
 and sucking the liquid vp to the mark, then, by closing the 
 upper opening with the forefinger, the liquid is prevented 
 from running out, but may be delivered in drops or allowed 
 to run out to any mark by lessening the pressure of the 
 finger over the opening. 
 
 In using the pipettes of the first class (Fig. 17) the finger 
 
 Fig. 14. 
 
 Fig. 13. 
 
 Fig. 15. 
 
 is raised and the instrument allowed to empty itself entirely. 
 A drop or two, however, usually remains in the lower portion 
 of the instrument, which may be blown out. By inclining the 
 pipette and placing the point against the side of the vessel 
 which is to receive the liquid, the instrument may be emptied 
 more satisfactorily. 
 
 Pipettes of the second class (Fig. 18) are never emptied 
 completely when in use. The flow of the liquid is regulated 
 by the pressure of the finger over the upper opening, and 
 stopped at the desired point. 
 
APPARATUS USED IN VOLUMETRIC ANALYSES 
 
 35 
 
 A very convenient form of pipette is one which has attached 
 to its upper end a piece of rubber tubing, into which a short 
 piece of glass rod has been inserted. By squeezing the 
 
 Fin. i6. 
 
 rubber surrounding the glass rod firmly between the fingers, a 
 canal is opened and the liquid can be drawn up into the 
 pipette by suction with the lips and run out again. By re- 
 moving the pressure the canal closes and the flow of the 
 liquid is stopped at any point (Fig. 19). 
 
36 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The Nipple Pipette is very convenient for measuring small 
 quantities of liquids, such as i or 2 cc. (Fig. 20). 
 
 When a volatile or highly poisonous solution is to be 
 measured it is not advisable to suck it up with the mouth. 
 
 50 ca 
 
 JiQ cc. 
 
 Fig. 17. 
 
 Fig. 18. Fig. iq 
 
 The pipette in this case is filled by dipping it into the liquid 
 contained in a long, narrow vessel, until the liquid reaches 
 the proper mark on the pipette, then closing the upper opening 
 and withdrawing. When this is done the liquid which 
 adheres to the outside of the pipette should be dried off 
 before the measured liquid is delivered. 
 
 A French firm has introduced pipettes provided with suction 
 
APPARATUS USED IN VOLUMETRIC ANALYSES 
 
 37 
 
 pumps, shown in various forms by Fig. 21, which possess 
 the advantage over the ordinary forms provided with a com- 
 pressible rubber bulb, that the liquid can with perfect ease 
 be drawn up to the desired point on the scale, and with 
 absolute accuracy maintained at the same height as long as 
 may be desired. 
 
 The Measuring-flask is a vessel made of thin glass having 
 
 Fig. 20. 
 
 Fig. 21. 
 
 a narrow neck, and so constructed as to hold a definite amount 
 of liquid when filled up to the mark on the neck. These 
 flasks are of various sizes, holding 100, 250, 500, 1000 cc, 
 etc., but are generally called " Liter Flasks." (Fig. 22.) 
 
 Liter flasks are used for making volumetric solutions. 
 
 Those which have the mark below the middle of the neck 
 are to be preferred, because the contents can be more easily 
 shaken. 
 
38 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Liter flasks are sometimes made with two marks on the 
 neck very near together; the lower one is the Uter mark. If 
 the flask is reqxiired to deliver a liter, it must be filled to the 
 upper mark, the difference between the two measures being 
 
 Fig. 22. 
 
 Fig. 23. 
 
 Fig. 24. 
 
 the equivalent of the liquid which remains in the flask adhering 
 to the sides. 
 
 The Test-mixer, or Graduated Cylinder (Fig. 23) is for 
 measuring and mixing smaller quantities of solutions. They 
 are made of different sizes, holding 100, 250, 500 and 1000 
 cc, and graduated in fifths or tenths of a cc. 
 
 Titration Flasks. Titrations may be carried out in flasks 
 of any usual shape, or in beakers, or evaporating dishes, but 
 the flask illustrated in Fig. 24 is to be preferred. 
 
CHAPTER VI 
 ON THE USE OF APPARATUS 
 
 It is important that all apparatus used in volumetric 
 analysis be perfectly clean. Even new apparatus should be 
 cleansed by passing dilute hydrochloric acid through them 
 and then rinsing with distilled water. 
 
 If the burette, pipette, or other instrument is even slightly 
 greasy, the liquid will not flow smoothly, and drops of the 
 liquid will remain adhering to the sides, thus leading to 
 inaccurate results. 
 
 Greasiness may be removed with dilute soda solution. 
 If this fails the instrument should be allowed to remain for 
 some little time in a solution containing sulphuric acid and 
 potassium dichromate, which will radically remove all traces 
 of grease. 
 
 The burette or other measuring instruments should never 
 be filled with volumetric solution without first rinsing, even 
 if the burette be perfectly dry. 
 
 It is well to wash the inside of the instrument with two 
 or three small portions of the solution with which it is to be 
 filled. 
 
 The burette may be filled with the aid of a funnel, the 
 stem of which should be placed against the inner wall of the 
 burette, so that the solution will flow down the side and thus 
 prevent the formation of bubbles. 
 
 The burette should be filled to above the zero mark, and 
 the air-bubbles, if there are any, removed by gently tapping 
 
 with the finger. 
 
 39 
 
40 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 A portion of the liquid should then be allowed to run out 
 in a stream so that no air-bubbles remain in the lower part 
 of the burette. In the glass tap burette it can be easily seen 
 if any air is present, but in the pinch-cock burette it is some- 
 times necessary to take hold of the rubber tube between the 
 thumb and forefinger and gently stroke upward. ' Or the 
 glass nip at the lower end of the burette may be pointed upward, 
 and the pinch-cock opened wide so that a stream of the liquid 
 will pass through and force out any air that may be inclosed. 
 
 If the titration is to be conducted at a high temperature, 
 as in the estimation of carbonates, when litmus is used as 
 the indicator, or in the estimation of sugar by copper solution, 
 a long rubber tube should be attached to the lower end of 
 the burette. The boiling can then be done at a little distance, 
 and the expansion of the liquid in the burette avoided. The 
 pinch-cock is fixed about midway on the tube. 
 
 Hart calls attention to the fact that if the fluid in a burette 
 or pipette be rim out rapidly at one time and slowly at another, 
 different amounts of fluid are obtained. 
 
 This is due to the adhesion of the fluid to the inner sides 
 of the instrument, and reading before it has settled down. 
 It is therefore advisable always to deliver burettes slowly, 
 as more constant results are then obtained. 
 
 Solutions which are measured by means of pipettes should 
 be dilute, since concentrated solutions adhere to glass with 
 different degrees of tenacity, and hence the amount of fluid 
 delivered is slightly less than that measured. 
 
 The temperatiure of the solutions measured should be 
 taken into account, since aU liquids are affected by change 
 of temperature, expanding and contracting as the tempera- 
 ture is increased or reduced. 
 
 This change of volume in the case of standard solutions 
 does not exactly correspond to that in pure water; in fact, 
 
ON THE USE OF APPARATUS 41 
 
 some of them differ widely. The correction of the volume 
 of a standard solution for the temperature by the expansion 
 coefficient of water is not entirely satisfactory, but in the case 
 of very dilute solutions this may be done. 
 
 Casamajor (C. N., xxxv, i6o) gives the following figures 
 showing the relative contraction and expansion of water below 
 and above 15° C: 
 
 Degrees C. Degrees C. 
 
 8— .000590 17 + .C00305 
 
 9— .000550 18 + .000473 
 
 10— .000492 19 + .000652 
 
 11 — .000420 20 + .000841 
 
 1 2 — .000334 2 1 + .00 1039 
 
 13 — .000236 22 + .001246 
 14— .000124 23 + .001462 
 
 15 — normal 24 + .001686 
 
 16 + . 000147 25+.001919 
 
 By means of these numbers it is easy to calculate the volume 
 of liquid at 15° C. corresponding to any volume observed 
 at any temperature between 8° C. and 25° C. If 25 cc. of 
 solution had been used at 20° C, the table shows that i cc. 
 of water passing from 15° to 20° is increased to 1.000841 cc. 
 Therefore, by dividing 25 cc. by 1.000841, the quotient, 24.97 
 cc. is obtained, which represents the volume at 15° C. corre- 
 sponding to 25 cc. at 20° C. 
 
 These corrections are of value only for very dilute solutions 
 and for water, but useless for concentrated solutions. Slight 
 variations of atmospheric pressure may be disregarded. 
 
 ON THE READING OF INSTRUMENTS 
 
 In narrow vessels the surface of liquids is never level. 
 This is owing to the capillary attraction exerted by the sides 
 of the vessel upon the liquid, drawmg the edge up and forming 
 
42 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 a saucerlike concavity (Fig. 25). All liquids present this 
 concave surface except mercury, the surface of which is convex. 
 
 This behavior of liquids makes it difficult to find a distinct 
 level, and in reading the measure either the upper meniscus 
 (a) or the lower meniscus (b) may be used (Fig. 26). 
 
 The most satisfactory results are obtained if the lowest 
 point of the curve (&) is used, especially with light-colored 
 
 hr 
 
 m 
 
 Fig. 25. 
 
 Fig. 26. 
 
 Fig. 27. 
 
 solutions. But if dark-colored or opaque solutions are measured 
 it is necessary to use the upper meniscus (a) for reading. 
 
 In all cases the eye should be brought on a level with the 
 surface of the liquid in reading the graduation. 
 
 The eye is very much assisted by using a small card, 
 the lower half of which is black and the upper half white, 
 This card is held behind the burette, the dividing line between 
 white and black being about an eighth of an inch below the 
 surface of the liquid. The eye is then brought on a level 
 with it, and the lower meniscus can be distinctly seen as a 
 sharply defined black line against the white background (Fig. 
 27). 
 
ON THE USE OF APPARATUS 43 
 
 Erdmann's Float, Fig. 28, is an elongated glass bulb, which 
 IS weighted at its lower end with mercury, to keep it in an 
 upright position when floating. It is of such diameter that 
 
 Fig. 28. Fig. 29. Fig. 30. 
 
 it will slide easily up and down inside of a burette. There 
 is a ring at the top by which it can be lifted in or out by 
 means of a bent wire. Around its center a line is marked. 
 At this line instead of at the meniscus the reading is 
 taken. 
 
44 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 These floats are sometimes provided with a thermometer, 
 and they then register the temperature as well as the volume. 
 Others are provided with projecting points along the sides, 
 
 Fig. 31. 
 
 Fig. 32. 
 
 the object of which is to prevent it from adhering to the walls 
 of the burette. (See Fig. 29.) 
 
 For the purpose of facilitatmg the reading, special forms 
 of burettes are constructed which are provided with a dark 
 longitudiaal stripe on a white enameled background (Fig. 
 30); the reflection of the dark stripe with the meniscus pro- 
 duces the peculiar appearance shown in Fig. 3 1 . The narrowest 
 point is at the middle of the meniscus, and by reading from 
 this point very accurate measurements are obtained. The 
 
ON THE USE OF APPARATUS 45 
 
 same effect can be produced by holding behind an ordinary 
 burette a white flexible card having a heavy black longitudinal 
 stripe, about one-eighth inch in width. 
 
 Another form of burette designed for the purpose of 
 facilitating reading is that provided with white enameled 
 sides, leaving a strip of clear glass in front and back (Fig. 
 32). This form is especially adapted for use with dark- 
 colored liquids, such as iodin and permanganate. 
 
 CALIBRATION OF INSTRUMENTS 
 
 Burettes are made from tubes of nearly uniform width. 
 They are filled with distilled water at 15° C* (59° F.) to the 
 o mark, and then 25, 50 or 100 cc. run out, and another mark 
 made at the surface of the hquid. The distance between 
 these two marks is then divided into 25, 50 or 100 equal 
 parts, and the spaces again subdivided into fifths and tenths. 
 Now it is very rarely possible to obtain tubes of exactly the 
 same caliber throughout, and the divisions made as above 
 do not always represent exactly what they are intended to. 
 
 If the tube is wider at one point the divisions at that 
 point will contain more, and if it is narrower they will contain 
 less than they should. 
 
 Hence before using a new burette, or in fact any other 
 measuring instrument, it is essential that the error, if any, 
 should be determined. This is done as follows: 
 
 Fill the burette to the o mark with distilled water at 15° 
 C. (59° F.) and run out 10 cc. at a time into a small weighed 
 flask, and weigh after each addition of 10 cc. 
 
 Each 10 cc. should weigh exactly 10 gms., and every 
 
 * Instead of 15° C. (59° F.) the temperature 25° C. (77° F.) is recommended 
 because this more nearly approaches the ordinary temperature of the atmosphere 
 in temperate climes. 
 
46 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 deviation found should be noted and taken into consideration 
 in using the instrument. 
 
 Example 
 Flask weighed 20.0000 grams. 
 
 + 10 cc. 
 
 
 30.1005 " 
 
 + 20 cc. 
 
 
 40.0499 " 
 
 +30 cc. 
 
 
 49.8000 " 
 
 +40 cc. 
 
 
 59.9700 " 
 
 + SOCC. 
 
 
 70.0100 " 
 
 Thus the ist 10. cc weighed 10.1005 grams. 
 
 2d 10 cc. " 9-9494 " 
 
 3d 10 cc. " 9-750I " 
 
 4th 10 cc. " 10.1700 " 
 
 5th 10 cc. " 10.0400 " 
 
 Having obtained these data, a table like the following may 
 be constructed and kept in some convenient place where it 
 can be readily consulted whenever the burette it represents 
 is being used. It is not necessary to carry the figure beyond 
 the second decimal place. 
 
 No. of cc. 
 
 No. of cc. 
 
 No. of cc. 
 
 No. of cc. 
 
 No. of cc. 
 
 No. of cc. 
 
 as read on 
 
 as 
 
 as read on 
 
 as 
 
 as read on 
 
 as 
 
 burette. 
 
 corrected. 
 
 burette. 
 
 corrected. 
 
 burette. 
 
 corrected. 
 
 I 
 
 1. 01 
 
 14 
 
 14.06 
 
 27 
 
 26.79 
 
 2 
 
 2.02 
 
 IS 
 
 15-05 
 
 28 
 
 27 
 
 76 
 
 3 
 
 3 03 
 
 16 
 
 16.04 
 
 29 
 
 28 
 
 73 
 
 4 
 
 4.04 
 
 17 
 
 i7°3 
 
 30 
 
 29 
 
 7° 
 
 5 
 
 5-05 
 
 18 
 
 18.02 
 
 31 
 
 3° 
 
 71 
 
 6 
 
 6.06 
 
 19 
 
 19.01 
 
 32 
 
 31 
 
 72 
 
 7 
 
 7.07 
 
 20 
 
 20.00 
 
 33 
 
 32 
 
 73 
 
 8 
 
 8.08 
 
 21 
 
 20.97 
 
 34 
 
 33 
 
 74 
 
 9 
 
 9.09 
 
 22 
 
 21.94 
 
 35 
 
 34 
 
 75 
 
 10 
 
 10.10 
 
 23 
 
 22.91 
 
 36 
 
 35 
 
 76 
 
 II 
 
 11.09 
 
 24 
 
 23.88 
 
 37 
 
 36 
 
 77 
 
 12 
 
 12.08 
 
 25 
 
 24.85 
 
 38 
 
 37 
 
 78 
 
 13 
 
 13 -o? 
 
 26 
 
 25.82 
 
 39 
 
 38 
 
 79 
 
ON THE USE OF APPARATUS 47 
 
 There should be no greater deviation than 0.15 cc. A 
 burette which deviates more is best not used. In the 
 foregoing table there is a deviation of 0.30 cc. at one point. 
 
 In order to test the accuracy of a pipette, fill to the mark 
 with distilled water at 15° C. (59° F.); empty into a previously 
 weighed flask, weigh again and thus determine the weight of 
 the water measured. One gram is equal to i cc. 
 
 Liter flasks are tested as follows: 
 
 The flask, perfectly dry and clean, is counterpoised on a 
 balance capable of turning with .005 when carrying about 
 2000 grams; it is then filled to the mark with distilled water 
 at 15° C. (59° F.) and the increase in weight should be exactly 
 the number of grams as the cc. indicated at the mark. 
 
CHAPTER VII 
 
 METHODS OF CALCULATING RESULTS 
 
 N . 
 Each cc. of a — univalent volumetric solution contains 
 
 I 
 
 rTjVir of the molecular weight in grams of its reagent, and 
 
 will neutralize ttVit oi the molecular weight of a univalent 
 
 substance, or TjijVir of the molecular weight of a bivalent 
 
 substance. 
 
 N . 
 Each cc. of a — bivalent volumetric solution contains 
 I 
 
 xirV-ff of the molecular weight in grams of its reagent, and will 
 
 neutralize or combine with -jVinf of the molecular weight of 
 
 a bivalent salt, or roVff of the molecular weight of a univalent 
 
 salt. 
 
 N . 
 A — is only A the strength of a normal solution and will 
 
 neutralize only yV the quantity of salt, etc. 
 
 Normal and decinormal solutions of acids should neutralize 
 normal and decinormal solutions of alkalies, volume for vol- 
 ume. Decinormal solution of silver nitrate and decinormal 
 solution of hydrochloric acid or sodium chlorid should combine, 
 volume for volume, etc. 
 
 Rules for Direct Percentage Estimations: i. With normal 
 solutions tV or -5*0 (according to its atomicity) of the molec- 
 ular weight in grams of the substance is weighed for titration, 
 and the number of cc. of the V.S. required to produce the 
 desired reaction is the percentage of the substance whose 
 molecular weight has been used. 
 
 48 
 
METHODS OF CALCULATING RESULTS 49 
 
 Thus, if sodium hydroxid (NaOH) is to be examined by 
 titration with a normal acid solution ^V of its molecular 
 weight in grams, 4 gms. is weighed out, and each cc. of the 
 acid solution required represents one per cent of the pure salt. 
 
 If sodium carbonate (Na2C03) is to be titrated 5V oi its 
 molecular weight in grams, 5.3 gms. is taken. 
 
 2. With decinormal solutions -[^-^ or -^ of the molecular 
 weight in grams of the substance to be analyzed is taken, 
 and the number of cc. will, in like manner, give the percentage. 
 
 The following equations will serve to explain more fully: 
 
 .AT' 
 Sodium hydroxid with — sulphuric acid: 
 
 2NaOH + H2SO4 = NaaSOi + 2H2O. 
 
 2X40=80 2)98 
 
 10)40 ^g = 1000 CC. 
 
 4.0 gms. = 100 cc. 
 
 . N 
 Sodium carbonate with — sulphuric acid: 
 
 NaaCOs + H2SO4 = Na2S04 + H2O + CO2. 
 2)98 
 
 20) 106 ^g = 1000 CC. 
 
 5.3 gms. = 100 cc. 
 
 N 
 With — sulphuric acid: 
 10 
 
 2NaOH + H2SO4 = Na2S04 + 2H2O. 
 2X40=80 2)^ 
 
 ioo)_40 4g = 1000 cc. 
 
 0.40 gm. = 100 cc. 
 
 In the case of a trivalent substance as citric acid ^^ of 
 the molecular weight in grams is taken for analysis when a 
 normal solution is employed and ^lu when a decinormal 
 solution is used. 
 
 In other words, when it is desired that each cc. of the 
 
50 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Standard solution should represent one per cent of the substance 
 upon which it acts, the rule is to take for analysis as much 
 of the substance as is represented by loo cc. of the standard 
 solution. 
 
 In the case of substances whose percentage of purity is 
 high, it is advisable to take smaller quantities, in order to 
 avoid the use of excessive quantities of standard solution. 
 Thus sulphuric acid, which contains 92.5 per cent of absolute 
 sulphuric acid, would require under the above conditions 92.5 
 cc. of normal alkali solution. 
 
 In the case of this acid, if 4.9 gms. are taken for analysis, 
 each cc. of a normal alkali solution would represent one per 
 cent of H2SO4. 
 
 If half of this quantity, i.e., 2.45 gms. are taken for 
 analysis, each cc. of the normal alkali will represent two per 
 cent of H2SO4, and thus less of the standard solution will 
 be required. Again, if 0.49 gm. be taken, each cc. of the 
 standard alkali will represent 10 per cent of H2SO4. 
 
 In the case of liquids where it is not always convenient 
 to weigh off the exact quantity required for titration by the 
 direct percentage method, the liquid is diluted to a convenient 
 degree with water, and then a quantity of this diluted liquid 
 (representing the weight required of the substance) is measured 
 for analysis. 
 
 Example. A sulphuric acid solution of specific gravity 
 1.826 is to be analyzed. Two cc. are accurately measured and 
 diluted to 100 cc. and then 13.41 cc. of this solution (repre- 
 senting 0.49 gm. of the acid) are taken for analysis. 
 
 N 
 Each cc. of — NaOH V.S. required in the titration rep- 
 
 N 
 resents 10 per cent of absolute H2SO4. If — NaOH V.S. 
 
 10 
 
 is employed, each cc. will represent one per cent. To de- 
 
METHODS OF CALCULATING RESULTS HI 
 
 termine the amount of the dikited Hquid to be taken we 
 proceed as follows: 
 
 Two cc. of sulphuric acid, specific gravity 1.826, weigh 
 3.652 gms., therefore the 100 cc. of diluted acid contain this 
 weight, and i cc. of the same contains 0.03652 gm. 
 
 If 0.03652 gm. are contained in i cc, then 0.49 gm. 
 are contained in how many cc. ? 
 
 gm. cc. gm. 
 
 .03652 : 1 : : 0.49 :x, 
 
 x= 13.41 cc. 
 
 Factors or Coefficients for Calculating the Analyses. It 
 
 frequently occurs that from the nature of the substance, or 
 from its being in solution, this percentage method cannot be 
 conveniently followed. 
 
 The best way to proceed in such a case is to find the 
 factor. 
 
 The first step in all cases is to write the equation for the 
 reaction which takes place between the substance under 
 analysis and the solution used. 
 
 For instance, a solution of caustic potash is to be examined, 
 
 N 
 a — solution of sulphuric acid being used. 
 
 2KOH + H2SO4 = K2SO4 + 2H2O. 
 
 2 )1" 2)95 j,^ 
 
 56 49 = 1000 CC. — acid. 
 
 0.56 gm. .049 = I cf- — acid. 
 
 N 
 The factor for KOH when — solution is used is .056 gm., 
 
 N 
 that being the quantity neutralized by each cc. of the — acid. 
 
 N 
 If — acid were used the factor would be .0056 gm. 
 10 
 
52 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The number of cc. of the acid used to produce the desired 
 result, when multiplied by the factor, gives the quantity in 
 grams of KOH in the solution taken. 
 
 Example. If lo grams of caustic potash solution were 
 
 N 
 taken, and 40 cc. of — acid were required, the 10 gms. of 
 
 solution contained .056 gm.X4o = 2.24 gms. of pure KOH. 
 
 To find the percentage the following formula may be 
 used: 
 
 QXioo 
 W 
 
 = % 
 
 0- 
 
 Q =the quantity of pure substance foimd by calculation; 
 If = weight of substance taken. 
 
 If the above example is taken, we have 
 
 2.24X100 „ 
 ^^— = 22.4%. 
 
 Or the calculation may be made by proportion. 
 
 The quantity of the substance taken is always the first 
 term, and the quantity of pure substance found, the second 
 term. 
 
 The following rule is easily remembered: As the quantity 
 taken is to the quantity found, so is 100 to x, the percentage 
 of pure substance in the sample. 
 
 Three terms of the equation being given, the fourth is 
 found by multiplying the means and dividing the product by 
 the given extreme. By applying this rule to the above case 
 we have 
 
 10 : 2.24 : : 100 : X. ^ = 22.4%. 
 
METHODS OF CALCULATING RESULTS 
 
 53 
 
 TABLE SHOWING THE NORMAL FACTORS, ETC., OF THE 
 ALKALIES, ALKALI EARTHS, AND ACIUS. 
 
 Substance. 
 
 Formula. 
 
 Molecular 
 
 Normal 
 
 Weight. 
 
 Factor.* 
 
 40 
 
 0.040 
 
 106 
 
 °-°53 
 
 84 
 
 U.084 
 
 56.1 
 
 0.0561 
 
 138.2 
 
 0.0691 
 
 100. 1 
 
 . lOOI 
 
 17 03 
 
 0.01703 
 
 96.08 
 
 0.04804 
 
 56.. 
 
 0.02801; 
 
 74-1 
 
 0.0370s 
 
 100. 1 
 
 0.050 
 
 63.01 
 
 0.063 
 
 36.46 
 
 0.03646 
 
 98.08 
 
 . 04904 
 
 1 26 . 05 
 
 0.063 
 
 60.03 
 
 . 0603 
 
 Quantity of 
 
 Substance 
 
 to be takent 
 
 so that each 
 
 N 
 cc. of - V. S. 
 
 1 
 will indicate 
 I per cent. 
 
 Sodium hydroxid 
 
 Sodium carbonate 
 
 Sodium bicarbonate 
 
 Potassium hydroxid 
 
 Potassium carbonate 
 
 Potassium bicarbonate .... 
 
 Ammonia (gas) 
 
 Ammonium carbonate, normal 
 
 Lime 
 
 Calcium hydroxid 
 
 Calcium carbonate 
 
 Nitric acid 
 
 Hydrochloric acid 
 
 Sulphuric acid 
 
 Oxalic acid, crystallized . . 
 Acetic acid 
 
 NaOH 
 
 NazCOj 
 
 NaHCOj 
 
 KOH 
 
 KjCOs 
 
 KHCO3 
 
 NH3 
 
 CaO 
 
 Ca(OH)2 
 
 CaCOa 
 
 HNO3 
 
 HCl 
 
 HjSO^ 
 
 HjC20,-2HjO 
 
 HCoHjO, 
 
 4.0 
 
 5-3 
 8.4 
 
 5-61 
 
 6.91 
 
 10.01 
 
 703. 
 804 
 805 
 70s 
 
 3 
 646 
 
 9 
 3 
 °3 
 
 * This is the coefficient by which the number of cc. of normal solution used is 
 to be multiplied in order to obtain the quantity of pure substance present in the 
 material examined. 
 
 t This is the quantity of substance to be taken in direct percentage estimations^ 
 
 Each cc. of — acid or alkali V.S. employed will then indicate i per cent; in the case 
 
 I 
 of many of these substances it will, however, be better to take smaller quantities so 
 that less of the standard solution be required. Thus if one-half the quantity be 
 
 N 
 taken each cc. of the — V. S. will represent 2 per cent, if A of the quantity be taken 
 
 N 
 each cc. will represent 10 per cent. If. however, — solutions be used and A of the 
 
 quantity indicated in the table be taken each cc. will indicate i per cent. 
 
 On Stating Results. In reporting the results of volu- 
 metric work, it is customary to state the quantity of pure 
 substance found; thus in the case of salts, the quantity of 
 the anhydrous salt is reported. It is, however, often required 
 
54 THE ESSENTIALS OF V0LUM]:;TRIC ANALYSIS 
 
 to state the results according to the duahstic formulae of 
 Berzelius, that is, the metals are reported as oxids and the 
 acids as anhydrids. Thus if sodium carbonate is analyzed, 
 a statement of results by this method will give the quantity 
 of Na20, spoken of as the base, soda. If we look upon 
 sodium carbonate as Na20C02, we can readily see io6 gms. 
 of the anhydrous salt contain 62 gms. of Na20. 
 
 By reference to the following equations we see that 98.08 
 gms. of sulphuric acid will neutralize 62 gms. of Na20 or 
 106 gms. of Na20C02. 
 
 NaaO + H2SO4 = NaaSOi + H2O. 
 
 2)62 2)98.08 j^ 
 
 31 gms. 49.04 gms. =1000 cc. — V.S. 
 
 Na20C02+H2S04=Na2S04+H20 + C02. 
 
 53 gms. = to 1000 cc. — V.S. 
 
 I 
 
 N 
 Thus one cc. of — H2SO4 will represent 0.031 gm. of 
 
 Na20 and 0.053 g™- of Na20C02. 
 
 In the case of sodium bicarbonate (NaHCOs) two molecules 
 contain one molecule of the base, soda, as here shown. 
 
 2NaHC03 = Na20,H20 (C02)2. 
 
 According to this 62 gms. of Na20 represent two molecular 
 weights (168 gms.) of NaHCOa. In the case of ferrous sul- 
 phate, one molecule (FeS04) contains, according to this system, 
 FeO and SO3. In the same way, ferric salts contain Fe203. 
 In stating the results of analyses of acids according to this 
 system, the quantity of acid anhydrid found is reported, not 
 the quantity of the whole acid. Thus if sulphuric acid is 
 analyzed, the quantity of SO3 is reported. In the case of 
 phosphoric acid the quantity of P2O5 is stated, etc. 
 
METHODS OF CALCULATING RESULTS 
 
 55 
 
 TABLE SHOWING THE MCJLECULAR WEIGHTS AND NORMAL 
 FACTORS FOR THE MoS [' COMMON OXIDS 
 
 Name. 
 
 Soda , 
 
 Potash 
 
 Lime 
 
 Magnesia 
 
 Lithium oxid 
 
 Strontium oxid 
 
 Barium oxid 
 
 Zinc oxid 
 
 Lead oxid 
 
 Arsenous oxid 
 
 Antimonous oxid . . . . 
 
 Mercurous oxid 
 
 Mercuric oxid 
 
 Ferrous oxid 
 
 Ferric oxid 
 
 Silver oxid 
 
 Sulphuric anhydrid . . 
 Phosphoric anhydrid . 
 
 Nitric anhydrid 
 
 Carbonic anhydrid . . 
 
 Formula. 
 
 Molecular 
 Weight. 
 
 ^ Factor. 
 
 Na^O 
 
 62.0 
 
 0.031 gm. 
 
 K2O 
 
 94.^ 
 
 U.0471 " 
 
 CaO 
 
 56. 1 
 
 0.028 " 
 
 MgO 
 
 40.32 
 
 0.020 " 
 
 LijO 
 
 29.88 
 
 0.0149 " 
 
 SrO 
 
 103-63 
 
 0.0518 " 
 
 BaO 
 
 IS3-37 
 
 0.0767 " 
 
 ZnO 
 
 81.37 
 
 0.0407 " 
 
 PbO 
 
 223.1 
 
 0.1115 " 
 
 AsjO, 
 
 197.92 
 
 0.049s " 
 
 Sb^Os 
 
 288,4 
 
 0.0721 " 
 
 Hg,0 
 
 416.0 
 
 0.208 
 
 HgO 
 
 216.0 
 
 0.108 " 
 
 FeO 
 
 71.82 
 
 0.0718 " 
 
 Fe,03 
 
 1 59-. 64 
 
 0.0798 " 
 
 A&O 
 
 231.76 
 
 0.1159 " 
 
 SO3 
 
 80.07 
 
 U.040 " 
 
 P2O5 
 
 142.08 
 
 0.02368 " 
 
 N2O5 
 
 108.02 
 
 0.054 
 
 COj 
 
 44-" 
 
 0.022 " 
 
CHAPTER Vni 
 
 ANALYSIS BY NEUTRALIZATION 
 
 This is based upon the fact that when an acid and an 
 
 alkah react each loses its individuality and a neutral salt is 
 
 formed, i.e., a body which has neither the character of an 
 
 acid nor that of an alkali. This mutual neutralization of 
 
 + 
 acid and alkali is the result of a union of the H' ions of the 
 
 acid and the OH' ions of the alkaU, forming non-ionized water 
 (HOH). 
 
 An acid is a compound which in aqueous solution disso- 
 ciates (ionizes) into positive and negative ions, the positive 
 ion being H. Thus hydrochloric acid in an ionized state is 
 
 + - + 4- - 
 
 H" -l-CI". Sulphuric acid ionizes into H' -f-H' + (SO4). 
 
 An alkali is a basic compound which ionizes into positive 
 and negative ions, and in which the negative ion is (OH). 
 
 The reaction between hydrochloric acid and potassium 
 hydroxid. in accordance with this theory, is illustrated by 
 the following equation: 
 
 H-FCi + K + OH = K + ci+HOH. 
 
 An acid is generally recognized as such by its color reac- 
 tions with certain substances known as indicators; for example, 
 it turns blue litmus red, and decolorizes a red phenolphthalein 
 solution. Alkalies are recognized by their turning red litmus 
 
 56 
 
ANALYSIS BY NEUTRALIZATION 57 
 
 blue, and by producing a deep red color with phenol- 
 phthalein. 
 
 The strength of an acid solution is ascertained by noting 
 the quantity of alkali that is required to neutralize it. The 
 stronger the acid, the more alkali is required. The strength 
 of an alkali is estimated by the quantity of acid which is 
 required to neutralize it. The estimation of the strength of 
 acids is called acidimetry, while the estimation of alkalies 
 is called alkalimetry. 
 
 The principal alkaline substances which may be estimated 
 by means of standard acid solutions are the hydroxids and 
 carbonates of sodium, potassium, lithium and ammonium, and 
 the hydroxids and oxids of calcium, barium and strontium 
 and the alkaloids. 
 
 When an acid is brought in contact with an alkali, a 
 reaction takes place in which a neutral salt is formed. This 
 is known as neutralization, and takes place between definite 
 and invariable proportions of the reacting bodies; thus, if 
 1 12.2 parts of potassium hydroxid are mixed with 98.08 parts 
 of absolute sulphuric acid, the alkali as well as the acid will 
 be exactly neutralized. If only 80 parts of the acid have been 
 added the mixture would still be alkaline, for ,it requires 98.08 
 parts of the acid to neutralize 112.2 parts of potassium 
 hydroxid. If more than 98.08 parts of the acid have been 
 added, the mixture would be acid, and would consist of 
 potassium sulphate and free sulphuric acid. The reaction 
 is thus illustrated: 
 
 2KOH + H2SO4 = K2SO4 + 2H2O. 
 
 2K = 78.2 2H= 2.016 
 
 20 = 32.0 5=32.070 
 
 2H= 2.016 40 = 64.00 
 
 112. 216 98.086 
 
 Sodium hydroxid will unite with oxalic acid in the propor- 
 
u8 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 tion of 80.016 parts by weight of the former and 126.048 
 parts by weight of the latter, as the equation shows. 
 
 aNaOH + H2C2O4 ■ 2H2O = Na2C204 +4H2O. 
 
 2Na=46 6H-= 6.048 
 
 20 = 32 2C = 24.000 
 
 2H' = 2.016 60 = 96.000 
 80.016 126.048 
 
 Ammonia water unites with hydrochloric acid as per the 
 equation, 
 
 NH4OH + HCl = NH4CI + H2O. 
 
 3S-0S 36.46 
 Sodium carbonate with hydrochloric acid, 
 
 Na2C03 + 2HCI = 2NaCl + H2O + CO2. 
 
 106 72.92 
 
 Upon a careful perusal of the foregoing equations it will 
 be seen that since definite weights of acids neutralize definite 
 weights of alkalies, the quantity of a certain alkali in solution 
 can be easily determined by the quantity of an acid solution 
 of known strength required to neutralize it, and vice versa. 
 
 Referring to the first equation we see that 98.086 gms. 
 of H2SO4 neutralize 112. 216 gms. of KOH. If we prepare 
 a normal solution of H2SO4 we take half the molecular weight, 
 98.086=49.043 gms., to 1000 cc. Half the molecular weight 
 is taken because sulphuric acid is a bivalent acid. 1000 cc. 
 of this solution will neutralize 56.108 gms. of KOH; hence 
 I cc. will neutralize 0.056108 gm. of KOH. 
 
 Thus if 10 gms. of a solution of KOH be treated with 
 the above normal solution of H2SO4, and it is found that 
 25 cc. of the acid solution are required to neutralize the 
 alkali solution, the latter contains 25X0.0561 = 1.40+ gm. of 
 pure KOH. 
 
 Since the acid and alkali as well as the neutral salt which 
 
ANALYSIS BY NEUTRALIZATION 59 
 
 is formed are colorless, and no visible change takes place 
 during the reaction, it is necessary to add some substance 
 which by change of color will show when the neutralization 
 is complete. Such a substance is known as an indicator. 
 
 In the case of sodium hydroxid with oxalic acid (see the 
 second equation) we find that 126.048 gms. of crystallized 
 oxalic acid neutralizes 80.016 gms. of NaOH. Oxalic acid, 
 like sulphuric, is bivalent, therefore a normal solution of it 
 contains half the molecular weight in grams,, i.e., 63.024 gms. 
 in 1000 cc. 
 
 1000 cc. w.ill neutralize 40 gms. of NaOH; 
 I cc. will neutralize 0.040 gm. of NaOH. 
 
 The neutralizing power of all normal acids is exactly the 
 same, because they all contain in 1000 cc. the molecular 
 weight in grams of the .acid in the case of univalent acids, 
 and half of the molecular weight in grams of bivalent acids. 
 
 Thus I cc. of any normal acid will neutralize 0.0561 gm. 
 of KOH or 0.040 gm. of NaOH or y^^j-^ of the molecular 
 weight of any other imivalent alkali, or j^iVir of the molecular 
 weight of an alkali earth, the latter being bivalent. In like 
 manner all decLnormal solutions have a like neutralizing 
 power, their neutralizing equivalence is one-tenth that of 
 normal solutions. 
 
 Thus I cc. of a decinormal acid will neutralize 0.00561 
 gm. of KOH or 0.0040 gm. of NaOH, etc. 
 
 Alkalimetry 
 
 Preparation of Standard Acid Solutions. It is possible 
 to carry out the titration of most alkalies by means or one 
 standard acid solution, but the same standard acid is not 
 equally applicable in all cases; furthermore, the standard 
 acids are frequently employed for other volumetric operations 
 
60 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 than neutralization, and therefore it is advisable to have a 
 
 variety. 
 
 The standard oxalic acid solution is preferred by some, 
 
 because of the ease with which it may be prepared, provided 
 
 a pure oxalic acid is to hand. It does not, however, keep 
 
 very long, is unreliable for use with methyl orange, and is 
 
 inapphcable for the titration of alkali earths, because it forms 
 
 insoluble compounds with these metals. Standard hydrochloric 
 
 acid is the mbst desirable for alkali earths, because it forms 
 
 soluble compounds with them; its disadvantage, however, is 
 
 in its volatility and its consequent uselessness in hot titrations. 
 
 Standard sulphuric acid is preferred by most analysts as 
 
 being the best general standard. A pure acid can be gotten 
 
 without difficulty and the standard solution made from it is 
 
 unaffected by boiling, and can therefore be used in hot as 
 
 well as in cold titrations; it reacts sharply with the indicators 
 
 and it keeps its titer indefinitely. It is, however, not suited 
 
 for the titration of alkali earths, because it forms insoluble 
 
 compounds with them, which precipitate and are very annoying 
 
 to the operator. Ip. the preparation of standard solutions the 
 
 greatest care should be exercised in order that the product 
 
 be absolutely accurate. The slightest inaccuracy in the 
 
 strength of a standard solution will result in relative errors 
 
 in the analysis. It is customary to prepare one standard 
 
 solution, and then from this to adjust various others. For 
 
 example, a normal oxalic acid may be made first, and by 
 
 means of this a normal alkali solution, which in turn may 
 
 be utilized for the adjusting of other standard acid solutions. 
 
 N 
 Normal Oxalic Acid (H2C204- 21120 = 126.048; — V.S. 
 
 = 63.024 gms. in 1000 cc). Dissolve 63.024 gms. of pure 
 oxalic acid (see below) in enough water to make, at or near 
 15° C, exactly 1000 cc. 
 
AN.\LYSIS BY NEUTRALIZATION " 61 
 
 Pure oxalic acid, crystallized, is in the form of colorless, 
 
 transparent, clinorhombic crystals, which should leave no 
 
 residue when ignited upon platinum foil. It is completely 
 
 soluble in 14 parts of water at 15° C. If the acid leaves a 
 
 residue on ignition it should be purified by recrystallization, 
 
 as directed in the U. S. P. 
 
 N 
 One cc. of — oxalic acid V.S. is the equivalent of 
 
 Ammonia gas, NH3 0.017034 
 
 Potassium hydroxid, KOH 0.0561 
 
 Sodium hydroxid, NaOH 0.040 
 
 Potassium permanganate, KMn04 0.0316 
 
 Manganese dioxid, Mn02 0.04346 
 
 Calcium hydroxid, Ca(OH)2 o-037o53 
 
 /N 
 Decinormal Oxalic Acid I — 
 
 \io 
 
 V.S. = 6.3024 gms. in looo cc. ). 
 
 Dissolve 6.3024 gms. of pure oxalic acid in enough water to 
 make, at or near 15° C, exactly 1000 cc. 
 
 N 
 Normal Hydrochloric Acid (HC1 = 36.46; — V.S. = 36.46 
 
 gms. in 1000 cc). Mix 130 cc. of hydrochloric acid of sp. 
 gr. 1. 163 with enough water to measure, at or near 15° C, 
 1000 cc. 
 
 Of this hquid (which is still too concentrated) measure 
 carefully into a flask or beaker 10 cc, add 20 cc. of distilled 
 water and a few drops of phenolphthalein T.S., and then 
 gradually add from a burette sufficient recently prepared and 
 
 N 
 standardized — potassium or sodium hydroxid to just produce 
 
 a permanent faint pink tint. 
 
 N 
 Note the number of cc. of — alkali solution consumed 
 
 and then dilute the acid solution so that equal volumes of 
 
62 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 it and — alkali neutralize each other. It is usually advisable 
 
 to make two or three titrations, as just described, before 
 dilution, taking an average of the results. 
 
 Example. Assuming that the lo cc. of the acid solution 
 
 Fig. 33- 
 
 N 
 
 required 12 cc. of the — alkali, each 10 cc. of the acid must 
 
 be diluted to 12 cc, or the whole of the remaining acid in 
 the same proportion. 
 
 After the dilution a new trial should be made. 10 cc. 
 of the acid V.S. should require exactly 10 cc. of the alkali. 
 
ANALYSIS BY NEUTRALIZATION 63 
 
 This method is fairly satisfactory if an accurately stand- 
 ardized normal alkali hydroxid solution is at hand; the latter, 
 however, always contains a small quantity of carbonate, hence 
 methyl orange would be more desirable as an indicator. 
 
 Standardization by Means of Sodium Carbonate. Pure 
 anhydrous sodium carbonate may be obtained by heating 
 to dull redness a few grams of pure sodium bicarbonate for 
 about fifteen minutes. The resulting carbonate is practically 
 free from impurity. 
 
 The sodium bicarbonate loses on ignition one-half of its 
 carbonic acid gas: 
 
 2NaHC03-l-Heat = Na2C03 +CO2+H2O. 
 
 The bicarbonate should, however, be tested before igniting, 
 and if more than traces of chlorid, sulphate, or thiosulphate 
 are found, these may be removed by washing a few hundred 
 grams, first with a saturated solution of sodium bicarbonate, 
 and afterward with distilled water. 
 
 0.53 gm. of the pure anhydrous sodium carbonate is 
 accurately weighed and dissolved in about 200 cc. of water 
 in a flask and a few drops of methyl orange T.S. added as 
 indicator. The acid to be " set " or " standardized " is then 
 run into the sodium carbonate solution until a permanent 
 light-red color is produced. It should require exactly 10 cc. 
 
 of the — acid solution. 
 
 I 
 
 If 8 cc. of the acid solution are consumed to bring about 
 the required result, then every 8 cc. must be diluted to 10 
 cc, or the whole of the remaining solution must be diluted 
 in this proportion: 
 
 NagCOa + 2HCI = 2NaCl + H2O + CO2. 
 2)106 2)72.9 
 
 53 gms. 36.45 = 1000 cc, V.S. 
 
 0-53 gm- 
 
 = 10 ce. I 
 
61 THE ESSEXTIALS OF VOLUMETRIC ANALYSIS 
 
 This method may be employed as well for the standardization 
 of sulphuric or oxalic acid. 
 
 Other Methods for standardizing hydrochloric acid V.S. 
 are: (a) by means of silver nitrate (gravimetrically and volu- 
 metrically); (&) by means of borax; (c) by means of the 
 specific gravity; {d) by means of calc-spar. 
 
 N 
 Normal Sulphuric Acid (112804 = 98.07; -V.S.=48.675 
 
 gms. in 1000 cc). Mix carefully 30 cc. of pure concentrated 
 sulphuric acid (sp.gr. 1.835) ^i''^ enough water to make about 
 1050 cc, and allow the liquid to cool to about 15° C. 
 
 Titrate 10 cc. of this liquid in the manner described under 
 
 N 
 
 — hydrochloric acid, and dilute it so that equal volumes of 
 
 the acid and the alkali will neutralize each other. 
 
 The standardization of the normal sulphuric acid solution 
 may also be effected by the use of pure anhydrous sodium 
 carbonate, as described under normal hydrochloric acid V.S., 
 and by various other methods, among which are: (a) the 
 iodometric; (b) the specific gravity method; (c) the borax 
 method; (d) by precipitation with barium chlorid (gravi- 
 metrically). 
 
 Standard acid solutions are used in other strengths besides 
 
 N N 
 
 normal, namely, Half-normal — , Fifth normal — , Tenth-nor- 
 
 N N . . N 
 
 mal — , Twenlieih-normal — , Fiftieth-normal — ,and Hundredth- 
 10 20 50 
 
 normal^. 
 100 
 
 Estimation of Alkali Hydroxids 
 
 Potassium and sodium hydroxids are usually titrated with 
 
 N . 
 
 — sulphuric or hydrochloric acid; they are, however, so prone 
 
 to absorb carbon dioxid out of the air that they are seldom 
 free from carbonate, and hence the selection of a^l indicator 
 
ANALYSIS BY NEUTRALIZATION 65 
 
 is a matter of some importance. Phenolphthalein or litmus 
 may be employed, but it is then advisable to boil the solu- 
 tion while titrating, in order to drive off the liberated carbon 
 dioxid, because the latter gives an acid reaction with phenol- 
 phthalein and litmus and thus causes an end-reaction tint to 
 appear before neutralization is complete. It is better, usually, 
 to employ an indicator which is not affected by carbon dioxid. 
 Methyl orange is mostly preferred; cochineal and Congo red 
 are also useful. These indicators are especially serviceable 
 in the presence of carbonates in that they are not affected by 
 carbon dioxid, and can therefore be employed in direct titra- 
 tions without the use of heat. 
 
 The quantity of carbonate in a recent sample of sodium 
 or potassium hydroxid is so small usually that it is customary 
 to disregard it and to report the total alkalinity as hydroxid. 
 
 A definite quantity of the sample (from 0.5 gm. to i gm. 
 of the solid or an equivalent of a solution) is taken for analysis, 
 dissolved in 30 to 50 cc. of water in a white porcelain dish or 
 a beaker placed over a white surface, and a few drops of a 
 suitable indicator added. 
 
 The vessel is then placed beneath a burette containing 
 the standard acid solution and the latter run in, drop by drop, 
 until the last drop just causes the color to change. The 
 solution should be rotated or stirred after each addition of 
 the standard acid. 
 
 The alkali hydroxids are so exceedingly hygroscopic that 
 they take up water from the air while being weighed; it is 
 therefore difficult to make a direct weighing with any degree 
 of accuracy. 
 
 The best way is to take a small piece of the sample (about 
 I gm.), place it immediately in a tared stoppered flask and 
 take the weight accurately. It is then dissolved in water, 
 transferred to the porcelain dish or beaker and titrated. 
 
66 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Potassium Hydroxid (KOH = 56. i) . An accurately weighed 
 
 portion (preferably less than i gm.), is placed in a small beaker, 
 
 dissolved in 50 cc. of water, three drops of methyl orange added, 
 
 N 
 and the titration begun with — sulphuric acid and continued 
 
 until the yellow color of the solution is changed to red. Then 
 the burette is carefully read to see how much of the acid 
 solution was used. The number of cc. of the latter are 
 multiplied by the normal factor for KOH (0.05574) and the 
 result is the quantity of pure KOH in the sample taken for 
 analysis. 
 
 The following equation illustrates the reaction: 
 
 2KOH + H2SO4 = K2SO4 + 2H2O. 
 
 2 )112.2 2)98-07 j^ 
 
 56.1 gms. 49-03 gms., quantity in looo cc. of — acid V.S. 
 
 0.0561 gm. (the factor for KOH), quantity neutralized by 
 
 I cc. of — acid. 
 
 I 
 
 N 
 Thus 1000 cc. of — H2SO4 V.S. contammg 49.03 gms. 
 
 of absolute H2SO4 will neutralize 56.1 gms. of KOH. There 
 
 N 
 fore each cc. of — H2SO4 V.S. will neutralize 0.0561 gm. o: 
 
 pure KOH. 
 
 Example. In the above analysis let it be assumed that 
 0.915 gm. of potassium hydroxid were taken and that 15.3 
 cc. of the standard acid were required to neutralize it, then 
 0.0561 gm.X 15.3 =018583 gm., the quantity of pure KOH in 
 the 0.915 gm. taken. 
 
 The percentage is then calculated in this way: 
 
 0.915 : 0.8583 :: 100 :x; ^=93.8+per cent. 
 
 0.8583 X 1 00 _ g 
 0.915 93- • 
 
ANALYSIS BY NKUTRALIZATION 67 
 
 Sodium Hydroxid (NaOH = 4o). This is estimated in 
 exactly the same manner as described for potassium hydroxid, 
 the following equation being applied: 
 
 2NaOH + H2SO4 - Na2S04 + 2H2O. 
 
 2)80 2) 98.07 j^ 
 
 40 gms. 49-03 gms. = iooo cc— V.S. 
 
 .040 gm. I cc. — V.S. 
 
 The factor. 
 
 The official solutions of potassium and of sodium hydroxid 
 are estimated in this same manner, 10 cc. may be taken for 
 analysis, diluted with 20 cc. of water. 
 
 Ammonia Water (NH3 1120). Three cc. of ammonia water 
 
 are put into a stoppered weighing bottle and the weight taken. 
 
 40 cc. of water are then added and the solution titrated with 
 
 N 
 
 — sulphuric acid. As indicator, litmus, methyl orange or 
 
 rosolic acid may be used. Phenolphthalein is useless for 
 titrating ammonia, and even methyl orange and rosolic acid 
 are unsuitable in the presence of much salts of ammonium. 
 Because of the volatile character of ammonia its solutions 
 readily lose strength upon exposure. It is therefore best to 
 measure a quantity into a weighing bottle and find its weight 
 as directed for potassium hydroxids. If the specific gravity 
 of the ammonia solution is known, the weight of a given 
 volume is easily calculated, it being only necessary to mul- 
 tiply the volume in cc. by the sp.gr. Thus, if the sp.gr. of 
 an ammonia solution is 0.9585 and the volume taken is 3 
 cc, the weight of the 3 cc. is 3X0.9585 = 2.8755 gms. 
 
 N 
 In the titration with ^ sulphuric acid each cc. of the 
 
68 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 latter represents 0.017 SP^- ^^ NH3, as shown by the equa- 
 tion 
 
 2NH3.H20+H2S04=(NH4)2S04-1-2H20. 
 
 2)34 
 
 17 gms. 
 
 2)98.07 j^ 
 49.03 gms. = 1000 cc. — V.S. 
 
 ,017 
 Factor. 
 
 = xcJv.S. 
 
 I 
 
 N . 
 If 16.9 cc. of — acid were required in the above assay, 
 
 then C.017 gm.X 16.9 = 0.2873 gm., the quantity of pure NH3 
 in the 3 cc. (2 8755 gms.) of ammonia water taken. 
 
 The percentage is found as follows : 
 
 If 3 cc. of ammonia water weighing 2.8755 §™s. contain 
 0.2873 gm. of NH3, 100 gms. of ammonia water will contain 
 Xgm. of NH3, 
 
 0.2873X100 
 
 2-8755 
 
 = 9.99 per cent. 
 
 Stronger ammonia water and spirit of ammonia may be 
 estimated in the same manner. 
 
 Estimation of Alkali Carbonates 
 
 When carbonates are treated with acids carbonic-acid gas 
 is liberated. This gas shows an acid reaction with most 
 indicators, and the reaction will seem to be completed before 
 the alkali is entirely neutralized. 
 
 To avoid this, the titration may be conducted at the boiling 
 temperature (hot way) in order to drive off the carbon dioxid. 
 The standard acid being added until two minutes' boiling fails 
 to restore the color indicating alkalinity. If the titration is 
 conducted at a boiling temperature, it is advisable to attach 
 to the lower end of the burette a long rubber tube with a 
 pinch- cock fixed about midway on the tube. 
 
ANALYSIS BY NEUTRALIZATION 
 
 69 
 
 The boiling can then be done at a little distance from 
 
 the burette and the expansion of the standard solution therein 
 
 thus prevented. 
 
 Another method is to add to the carbonate a measured 
 
 excess of the standard acid, and then after boiling to drive 
 
 off the carbon dioxid, an indicator is added, and the excess 
 
 of standard acid determined by titration with a standard 
 
 alkali {residual titration way). 
 
 The quantity of the latter, 
 
 deducted from the quantity of 
 
 the standard acid taken, gives 
 
 the quantity of the latter which 
 
 reacted with the carbonate. Still 
 
 another method is to titrate the 
 
 carbonate direct, without heat 
 
 {cold way), using an indicator 
 
 which is not affected by carbon 
 
 dioxid. The best of the indica- 
 tors which are not so affected is 
 
 hiethyl orange; others are cochi- 
 neal and Congo red. When 
 employing methyl orange as an 
 indicator standard oxalic acid 
 solution should not be used, as 
 the end-reaction is very indefinite and unreliable. 
 
 The end-reaction with this indicator is at all events not 
 a clearly marked one, and considerable practice and an eye 
 for .color is required to detect the point at which yellow changes 
 to pale pink. It is a good plan to have on the bench two 
 vials, one containing an acid and the other an alkali tinted 
 with methyl orange, with which comparisons can be made. 
 
 Pottassium Carbonate (K2C03 = 138.2), Weigh carefully 
 one gram of the salt, dissolve in a small quantity of water in 
 
 Fig. 34 
 
70 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 a beaker or flask, add a few drops of methyl orange T.S., 
 and titrate with normal sulphuric acid imtil a faint orange- 
 red color appears. 
 
 K2CO3 + H2SO4 = K2SO4 + H2O + CO2. 
 
 2 )138-2 2)98.07 j^r 
 
 69.1 gms. 49.03 gms.= 1000 cc— V.S. 
 
 N 
 Each cc. of — H2SO4, therefore, represents 0.0691 gm. 
 
 of pure potassium carbonate. 
 
 If 14.3 cc. of the normal acid are required the salt contains 
 14.3X0.0691 gm. =0.98813 gm. of pure K2CO3 or 98.813 
 per cent. If it is desired to use litmus or phenolphthalein, 
 it will be necessary to boil the solution as described above. 
 
 Other alkali carbonates are estimated in exactly the same 
 manner as this, described for potassium carbonate. 
 
 Potassium Bicarbonate (KHCO3 =100.1). 
 
 2KHCO3 + H2SO4 = K2SO4 + 2H2O + CO2. 
 
 2)200.2 2)98.07 
 
 100. 1 gms. 49.03 gms. = 1000 cc. — V.S. 
 
 I 
 
 N 
 Each cc. of — acid V.S. =0.1001 gm. of KHCO3. 
 
 Sodium Carbonate (crystallized) (Na2C03.ioH20 = 286.i6). 
 
 NaaCOs-ioHaO + H2SO4 = Na2S04 + 11H2O + COe. 
 
 2)286.16 2)98.07 
 
 143.08 gms. 49.03 gms. = 1000 cc. — V.S. 
 
 N 
 
 Each cc. of — acid =0.143 S^- crystallized sodium car- 
 bonate. 
 
ANALYSIS BY NEUTRALIZATION 71 
 
 Sodium Carbonate (anhydrous) (Na2C03 = io6). 
 
 NaaCOg + H2SO4 = NagSOi + H^O + CO2. 
 
 2)106 2)98.07 ^ 
 
 53 gms. ' 49.03 gms.= 1000 cc. — V.S. 
 
 I 
 
 N 
 Each cc. of — acid =0.053 g™- Na2C03. 
 
 Sodium Bicarbonate (NaHC03 = 84). 
 
 2NaHC03 + H2SO4 = NaaSOi + 2H2O + 2CO2. 
 
 2)168 2)98.07 j^ 
 
 84 gms. 49.03 gms. = 1000 cc. — V.S. 
 
 N 
 
 Each cc. of — acid =0.084 g™- NaHCOa. 
 
 Lithium Carbonate (Li2C03 = 73.88). 
 
 Li2C03 + H2SO4 = LiaSOi + H2O + CO2. 
 
 2)73-88 2)98.07 j^ 
 
 36.94 gms. 49.03 gms. = *ooo cc. — V.S. 
 
 1 
 
 N 
 Each cc. of — acid = 0.03694 gm. Li2C03. 
 
 Ammonium Carbonate (N3HiiC205 = 157.03). Normal 
 ammonium carbonate has the formula (NH4)2C03, but the 
 normal salt loses upon exposure NH3 and H2O. The commer- 
 cial salt, therefore, generally is a mixture of bicarbonate and 
 carbamate. 
 
 (NH4) 2CO3 - NH3 = NH4HCO3 ; 
 
 (NH4)2C03 - H2O = NH4NH2CO2. 
 
 The commercial carbonate is therefore generally expressed 
 thus: 
 
 NH4HCO3.NH4NH2CO2 or N3H„C205. 
 This salt may be estimated by direct titration with normal 
 
72 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 or decinormal acid, using rosolic acid or methyl orange as 
 an indicator. 
 
 Two grams of the salt are taken, dissolved in about 50 
 
 N 
 cc. of water and titrated with — H2SO4 V.S. The reaction 
 
 is as follows : 
 
 2N3H11C2O5 + 3H2SO4 = 3(NH4)2S04 + 4CO2 + 2H2O. 
 
 6)314.0 6 6 )294.18 j^ 
 
 52.34 gms. 49'°3 gms. = iooo cc. — acid V.S. 
 
 N 
 
 Each cc. of — acid V.S. represents 0.052 gm. of N3 H11C2O5 
 
 or 0.017 gm. of NH3. 
 
 If in this titration 37.3 cc. of the standard acid are required 
 then the two grams of ammonium carbonate contained 0.052 
 gm.X37.3 = 1.939 gms. of the salt. 
 
 1.9^9X100 
 
 ^ =90.95 per cent. 
 
 If rosolic acid is used as indicator heat must be applied 
 to expel carbon dioxid. The estimation of the carbonic acid 
 may be effected by precipitating a definite weight of the salt 
 with barium chlorid, collecting the precipitated barium car- 
 bonate, dissolving it in a measured excess of normal hydro- 
 chloric acid and retitrating with normal alkali. 
 
 The method usually employed by skilled analysts {the 
 residual titration method), is to add a measured excess of the 
 standard acid solution, and thus convert the ammonium car- 
 bonate into the less volatile ammonium sulphate; then gently 
 boil to get rid of CO2, and titrate back with a standard alkali 
 V.S. (using litmus as an indicator) until the excess of acid is 
 neutralized. The quantity of free acid thus found, when 
 deducted from the amount of acid first added, gives the quantity 
 which was required to neutralize the ammonium carbonate. 
 
ANALYSIS BY NEUTRALIZATION 73 
 
 Thus 2 gms. in solution of ammonium carbonate are 
 
 N 
 treated with 50 cc. of — H2SO4 V.S., which is more than 
 
 sufficient to neutralize it; the solution is then gently boiled 
 
 to drive off CO2, a few drops of litmus tincture added, and 
 
 N 
 then titrated with — KOH V.S. until the litmus no longer 
 
 shows an acid reaction and the solution is neutral. 
 
 N 
 Let us assume that 12.7 cc. of the — KOH V.S. were 
 ' I 
 
 N 
 used. By deducting the 12.7 cc. from the 50 cc. of — acid 
 
 first added, we find 37.3 cc. of the acid went into combination 
 with the ammonium salt, the calculation is then made as 
 described above. 
 
 Mixed Alkali Hydroxid and Carbonate 
 
 If it is desired to ascertain the proportion in which these 
 exist in a mixture, we proceed as follows: 
 
 First determine the total alkalinity by means of normal 
 hydrochloric acid, using methyl orange as an indicator. Then 
 dissolve a like quantity of the mixture in 150 cc. of water 
 and add sufficient barium chlorid to precipitate all of the 
 carbonate as barium carbonate, and then add water to make 
 200 cc. and set aside to settle. When the supernatant liquid 
 is clear take one-fourth (50 cc.) of it, and titrate with normal 
 hydrochloric acid, using phenolphthalein as indicator.* The 
 number of cc. multiplied by 4 will be the quantity of normal 
 
 * The slight error which occurs in this method because the volume of the 
 precipitate is included in the measured liquid, may be overcome by using 
 the entire quantity of liquid, including the precipitate (instead of taking one- 
 fourth of it), and titrating with oxalic acid V.S. in the presence of phenol- 
 phthalein. Oxalic acid in very dilute solutions does not react with alkali 
 earth carbonates. 
 
74 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 acid required by the caustic alkali. The difference between 
 this and the number of cc. representing the total alkalinity 
 is calculated as carbonate. 
 
 Example. Assuming that we are analyzing a mixture of 
 sodium hydroxid and carbonate. 
 
 Two grams of the substance are dissolved in water and 
 titrated with normal acid solution. 43.2 cc. of the latter are 
 required. Another 2 grams is dissolved, treated with barium 
 chlorid as directed, and one-fourth of the clear solution titrated 
 with normal acid. 5.6 cc. are required; then 5.6X4=22.4 
 cc, representing the sodium hydroxid. 
 
 43.2 cc. = total alkalinity; 
 — 22.4X0.040=0.896 gm. sodium hydroxid. 
 
 20.8X0.053 = 1.1024 gms. sodium carbonate. 
 
 Another way is to filter the mixture after barium chlorid 
 has been added, titrate the filtrate with normal acid to find 
 the quantity of hydroxid, then dissolve the precipitated barium 
 carbonate in normal hydrochloric acid in excess, and retitrate 
 with normal alkali, thus ascertainmg the amount of carbo- 
 nate. 
 
 When the alkaline carbonate is present in very small 
 quantities the method of Lunge may be employed. 
 
 A few drops of phenacetolin solution are added to impart 
 a scarcely perceptible yellow to the Hquid. Normal acid 
 solution is then run in until a pale rose tint appears, indicating 
 that all the alkali hydroxid is neutralized; the volume of acid 
 is noted, and the titration continued; the red color is inten- 
 sified, and when the carbonate is entirely decomposed a 
 golden-yellow color results. 
 
 Considerable practice is required with solutions of known 
 composition to accustom the eye to the changes of color. 
 
ANALYSIS BY NEUTRALIZATION 75 
 
 Mixed Alkali Bicarbonates and Carbonates 
 
 Thompson' s Method. Take 2 grains of the salt and dissolve 
 in 100 cc. of water. Divide the solution into two equal parts 
 and titrate one portion with normal acid solution, using methyl 
 orange as indicator, and note the quantity required. We will 
 assume 13 cc. 
 
 Then treat the second portion with a measured excess 
 (say 25 cc.) of normal sodium hydroxid solution free from 
 CO2. This converts the bicarbonate into carbonate. Now 
 add an excess of pure neutral barium chlorid solution in 
 order to precipitate all the carbonate as barium carbonate, 
 and then titrate with normal acid, using phenolphthalein as 
 indicator, to determine the excess of sodium hydroxid. 15 
 cc. are required. Thus 
 
 25—15 = 10 cc, the equivalent of bicarbonate; 
 and 13—10 = 3 cc, the equivalent of carbonate; 
 10 X. 084 = .840 gm. sodium bicarbonate; 
 3 X. 053 = .159 gm. sodium carbonate. 
 
 Estimation of Alkalies In the Presence of Sulphites 
 
 This is accompHshed by adding hydrogen peroxid to the 
 solution in order to convert the sulphite into sulphate, and 
 then titrating in the usual way with normal acid. 
 
 Mixed Potassium and Sodium Hydroxids 
 
 These are estimated by treatment with tartaric acid solu- 
 tion, which converts them into bitartrates. The bitartrate of 
 potassium is almost insoluble in solution of sodium bitartrate 
 and hence may be separated by filtering. The sodium bitar- 
 trate is estimated in the filtrate by titration with normal 
 sodium hydroxid solution. The potassium is found by 
 difference. 
 
76 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Estimation of Organic Salts of the Alkalies 
 
 The tartrates, citrates and acetates of the alkaU metals are 
 converted by ignition into carbonates, the whole of the base 
 remaining in the form of carbonate. 
 
 Each molecular weight of a normal tartrate gives when 
 ignited one molecular weight of carbonate: 
 
 K2C4H406=K2C03. 
 
 Every two molecular weights of an acetate or an acid' 
 tartrate give one molecular weight of carbonate: 
 
 2KC2H302 = K2C03; 
 
 2KHC4H406=K2C03. 
 
 Every two molecular weights of a normal citrate give 
 three molecular weights of carbonate : 
 
 2K3C6H507 = 3K2C03. 
 
 These reactions are taken advantage of in volumetric 
 analysis, and the tartrates, citrates and acetates of the alkalies 
 are indirectly estimated by calculating upon the quantity of 
 carbonate formed by burning them, the quantity of carbonate 
 being found by titration in the usual manner. 
 
 The Process. Before igniting, the salt to be examined 
 should be thoroughly dried in a desiccator over calcium chlorid 
 or in a drying oven, the latter only for such salts as have no 
 water of crystallization in their composition. If the weight 
 is taken before and after, the amount of moisture present is 
 determined. One or two grams of the dried salt is weighed 
 accurately, placed in a porcelain crucible, and heat applied 
 gradually, until dull redness is reached and white fumes cease 
 to be given off. Upon applying heat to the salt, the latter 
 swells, fuses, and then boils, and if the heat is applied too 
 rapidly at this point, there is apt to be a considerable loss 
 
ANALYSIS BY NEUTRALIZATION 77 
 
 of material through sputtering. The completion of the ignition 
 is known to be reached when the black contents of the crucible 
 is dry and crisp. The crucible is then allowed to cool, and 
 its contents treated with boiling water to dissolve out the 
 alkali carbonate, and the solution filtered through a small, 
 wetted filter into a flask or beaker. The filtrate should be 
 perfectly colorless. If it has a yellow or brownish color it 
 indicates incomplete ignition and should be rejected, and a 
 fresh quantity of the salt subjected to ignition. The contents 
 of the crucible and the filter should be washed with several 
 small portions of water until the washings no longer show 
 an alkaline reaction. The filtrate mixed with the wash water 
 is now titrated with standard sulphuric or hydrochloric acid, 
 using methyl orange as the indicator. From the quantity of 
 carbonate found in the filtrate, the equivalent amount of the 
 organic salt may be calculated. The quantity of standard 
 acid employed is multiplied direct by the factor for the original 
 salt. 
 
 In the case of salts of the alkali earths,* residual titration 
 should be resorted to. The residue in the crucible being 
 dissolved in standard hydrochloric acid, and retitrated with 
 standard alkali. 
 
 Lithium salts, because of the sparing solubility of the 
 carbonate in water, should also be titrated by the residual 
 method. 
 
 Potassium Tartrate (K2C4H406= 226.2). Two grams of 
 the salt are placed in a platinum or porcelain crucible and 
 heated to redness in contact with the air until completely 
 charred; that is to say, until nothing is left in the crucible 
 but carbonate and free carbon. 
 
 The crucible is now cooled, and its contents treated with 
 
 * Organic salts of the alkali earths subjected to ignition as above are 
 reduced partly to oxids. 
 
7S THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 boiling water, which dissolves the potassium carbonate, the 
 carbon being separated by filtration. In order to obtain 
 every trace of carbonate it is well to wash the crucible with 
 several small portions of hot water, and add the washings 
 to the rest of the filtrate through the filter. 
 
 If the salt is completely carbonized the filtrate will be 
 colorless, but if the carbonization is not complete the solution 
 will be more or less colored and should be rejected, and a 
 fresh quantity of the salt subjected to ignition. 
 
 To the filtrate, which contains potassium carbonate, add 
 
 N 
 a few drops of methyl orange, and titrate with — sulphuric 
 
 acid V.S. until a light orange-red color appears and the car- 
 bonate is neutralized. 
 
 The following equations will explain the reactions: 
 
 K2C4H4O6 = K2CO3 + C2 + CO + 2H2O 
 
 226.2 138.2 
 
 then 
 
 K2CO3 + H2SO4 = K2SO4 + H2O + CO2 
 138.2 98.07 
 
 therefore 
 
 K2C4H4O6 = K2CO3 = H2SO4 
 
 2)226.2 2)1,38.2 2)98.07 
 
 113.1 gms. = 69.1 gms. = 49.03 gms. = 1000 cc. — V.S. 
 N 
 
 and each cc. of — H2SO4 represents 0.1131 gm. of potassium 
 
 tartrate. 
 
 Example. Two grams of potassium treated as described 
 
 N 
 above require 16.3 cc. of — H2SO4. It therefore contains 
 
 0.1131X16.3 = 1.8435 gms. 
 
 1.8435 X 100 
 
 ^ = 92.17 per cent. 
 
ANALYSIS BY NEUTRALIZATION 79 
 
 Potassium and Sodium Tartrate (KNaC4H406.4H20 
 
 = 282.22) {Rochelle Salt). This salt is treated in exactly the 
 same way as described for potassium tartrate. 
 
 When ignited the double tartrate is converted into a double 
 carbonate of potassium and sodium: 
 
 KNaC4H406 = KNaCOs + etc.; 
 210. 1 122. 1 
 
 then 
 
 KNaCOg + H2SO4 = KNaS04 + CO2 + H2O 
 therefore 
 
 KNaC4H406 = KNaCOs = H2SO4 
 
 2)210.1 2)122.1 2)98.07 
 
 105.05 61.05 49.03 = 100000 v.s. 
 
 I 
 
 N 
 and each cc. of — H2SO4 represents 0.10505 gm. of 
 
 KNaC4H406. 
 
 Example. If one gram of rochelle salt treated .as above 
 
 N 
 described requires 7 cc. of — sulphuric acid, it contains 
 
 0.10505X7=0.7353 gm. = 73.53 per cent. 
 
 Potassium Bitartrate (KHC4H4O6 = 188. i) {Cream of Tartar) 
 The estimation of this salt is affected in the same way as the 
 tartrate. 
 
 The bitartrate having but one atom of potassium in its 
 molecule, it takes two molecules to form one molecule of 
 carbonate. 
 
 2KHC4H4O6 = K2CO3 = H2SO4 
 
 2)376-2 2 )98.07 i^ 
 
 188. 1 gms. 49.03 gms. = iooo cc. — V. S. 
 
 Each cc. of - H2SO4 V.S. =0.1881 gm. of KHC4H4O6. 
 
 Another way of estimating bitartrate is to dissolve a weighed 
 
 N 
 quantity in hot water and titrate with — potassmm hydroxid 
 
80 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 until neutral, and thus the amount of tartaric acid existing 
 
 as bitartrate is found. The bitartrate is acid in reaction. 
 
 In detail the method is as follows: 
 
 Two grams of the bitartrate are dissolved in loo cc. of 
 
 hot water, a few drops of phenolphthalein T.S. added, and 
 
 N 
 then titrated with — KOH V.S. until a faint, pink color 
 
 indicates that all of the acid has been neutralized. Not less 
 than 10.6 cc. of the normal alkali should be required, corre- 
 sponding to 98.9 per cent of pure salt. 
 
 The following equation will show the reaction: 
 
 KHC4H4O6 + KOH = K2C4H4O6 + H2O. 
 
 • 188.1 sS.i =1000 cc. of — KOH V.S. 
 
 I 
 
 N 
 Each cc. of — KOH V.S. represents 0.1881 gm. of 
 
 KHC4H4O6. 
 
 If 10.6 cc. are required for neutralization, then io.6Xo.i88i 
 = 1.99+ gms.: 
 
 i.ggXiCKD 
 
 =99-5 per cent. 
 
 Potassium Citrate (K3C6H507 = 306.3). 
 
 2K3C6H5O7 = 3K2CO3 = 3H2SO4. 
 
 6)612.6 6)414.6 6)294.18 
 
 N 
 
 102. 1 gms. 69.1 = 49.03 gms. = 1000 cc. — acid. 
 
 N 
 Thus each cc. of — acid represents 0.1021 gm. of pure 
 
 potassium citrate. 
 
 Potassium Acetate (KC2H3O2 = 98. i) . In estimating potas- 
 sium acetate the salt is ignited and the residue treated in 
 
ANALYSIS BY NEUTRATJZATION 81 
 
 exactly the same manner as in the estimation of the citrates 
 and tartrates before mentioned. 
 
 2KC2H3O2 = K2CO3. 
 
 2)196^ N 
 
 98.1 gms. = 1000 cc. — HjSO^. 
 
 N 
 Each cc. therefore of — H2SO4 corresponds to 0.0981 
 
 gm. of potassium acetate. 
 
 Sodium Acetate (NaC2H302.3H20 = 136.09). 
 
 2(NaC2H302.3H20) =Na2C03, 
 
 N 
 Each cc. of — H2SO4 V.S. represents 0.13609 gm. of 
 
 crystallized sodium acetate. 
 
 Sodium Benzoate (NaC7H502= 144.05). 
 
 2NaC7H502 = Na2C03. 
 
 N 
 Each cc. of — H2SO4 V.S. represents 0.14405 gm. of sodium 
 
 benzoate. 
 
 Sodium Salicylate (NaC7H503 = 160.05). 
 
 2NaC7H303 = Na2C03. 
 
 N 
 Each cc. of — H2SO4 V.S. represents 0.16005 S^- of sodium 
 
 salicylate. 
 
 Lithium Citrate (Li3C6H507 = 209.82). As stated before, 
 the organic salts of lithium and those of the alkali earth 
 metals are best titrated by the residual method, after ignition, 
 because the carbonates formed are insoluble in water. It is 
 likewise best to use standard hydrochloric instead of standard 
 sulphuric acid. The process for lithium citrate here given 
 exemplifies the method. 
 
 One gram of the salt is thoroughly ignited in a pprcelain 
 
82 
 
 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 crucible as described for potassium tartrate. The residue of 
 
 lithium carbonate is then dissolved out of the crucible by add- 
 
 N 
 ing 20 cc. of — hydrochloric V.S. and filtering. The crucible 
 
 and filter are washed with several small quantities of water 
 
 and the washings adde.d to the acid filtrate. Three drops of 
 
 methyl orange are now added, and the solution titrated with 
 
 N 
 
 — sodium hydroxid V.S. until the yellow color appears. 
 
 Assuming that 5.8 cc. of the standard alkali were required, 
 then 20—5.8=14.2 cc, the quantity of normal hydrochloric 
 acid which reacted with the lithium carbonate. This quantity 
 multiplied by the normal factor for lithium citrate, 0.06994, 
 gives the weight of pure salt in the i gm. taken. 
 
 0.06994X14.2 = 0.993+ gm. or 99.3 per cent. 
 
 The other lithium organic salts of the U. S. P. are assayed 
 gravimetrically by conversion to sulphate. 
 
 TABLE SHOWING THE NORMAL FACTORS, ETC., OF THE 
 ORGANIC SALTS OF THE ALKALI METALS. 
 
 Substance. 
 
 Formula. 
 
 Molecular 
 Weight, 
 
 Equivalent 
 Weight in 
 Carbonate. 
 
 Normal 
 Factor, 
 
 Lithium benzoate 
 
 " citrate 
 
 " salicylate 
 
 Sodium acetate 
 
 ' ' benzoate 
 
 ' salicylate 
 
 Potassium acetate 
 
 " bitartrate . . . . 
 
 " citrate 
 
 " tartrate 
 
 " and sodium 
 
 tartrate, . . 
 
 LiCjHjOj 
 
 LijCeHjO, 
 
 LiCjHjOj 
 
 NaC2H,Oj-3H20 
 
 NaCjHjOj 
 
 NaC,H503 
 
 KCHjO, 
 
 KHC\H,0, 
 
 KjCeHjOj-HjO 
 
 2K.AH,0„-H20 
 
 KNaC,H,0e-4H,0 
 
 128 
 
 o.S 
 
 36 
 
 209 82 
 
 92 
 
 144-05 
 
 36 
 
 136 
 
 09 
 
 S3 
 
 144 
 
 05 
 
 53 
 
 160 
 
 °5 
 
 53 
 
 98 
 
 I 
 
 69 
 
 188 
 
 I 
 
 69 
 
 324 
 
 37 
 
 207. 
 
 470 
 
 50 
 
 13S 
 
 282, 22 
 
 ■94 
 ■35 
 ■94 
 
 o , 1 2805 
 o 06994 
 o 14405 
 0.13609 
 o 14405 
 o 16005 
 o.ogSl 
 o . 1881 
 0.108 
 U.1131 
 
 0.105 
 
ANALYSIS BY NEUTRALIZATION 83 
 
 Estimation of AlIcaLi Metals in their Salts 
 
 This may be done by first converting the salt into a 
 sulphate, and then by means of barium hydroxid, forming 
 an alkali hydroxid which is finally converted into an alkali 
 carbonate by means of carbon dioxid. 
 
 (a) K2S04 + Ba(OH)2 = BaS04 + 2KOH;' 
 
 (b) 2KOH + CO2 =K2C03+H20. 
 
 The conversion of the original salt into a sulphate may 
 be done in several ways, depending upon whether the acid 
 in combination is a volatile or a non-volatile one, as described 
 below. 
 
 Alkalies Combined with Volatile Acids. A definite quantity 
 of the salt in solution is treated with an excess of sulphuric 
 acid and evaporated to dryness, and then further heated to 
 drive off some of the excess of sulphuric acid. The residue, 
 which consists of the alkali as a sulphate, is dissolved in water 
 and treated with a slight excess of barium hydroxid solution. 
 The mixture now contains the alkali in solution as hydroxid, 
 and a precipitate of barium sulphate (see equation (a)), also 
 the excess of barium hydroxid in solution. A stream of carbon 
 dioxid (CO2) is now passed through the mixture; this converts 
 the alkali hydroxid into carbonate and at the same time 
 removes the barium hydroxid by precipitating it as barium 
 carbonate (see equation (b)). When this conversion into car- 
 bonate is complete, the free carbon dioxid must be driven 
 off by boiling, because barium carbonate is converted into 
 the soluble barium bicarbonate, in the presence of free carbon 
 dioxid. The mixture now contains the alkali in solution as 
 a carbonate, and a sediment consisting of barium sulphate 
 and barium carbonate. This mixture is now made up to a 
 definite volume and the alkali carbonate titrated in the usual 
 
8-1 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 manner in an aliquot portion, which may be removed by 
 filtration, or by means of a pipette, if the precipitate settles 
 rapidly and leaves a clear supernatant liquid. 
 
 Alkalies Combined with Non-Volatile Acids. In the case 
 of alkali salts of non-volatile acids, as phosphoric, boric, 
 chromic, arsenic, molybdic, etc., the acid is removed by means 
 of lead acetate and the resulting alkali acetate converted into 
 sulphate by means of sulphuric acid. 
 
 (c) K2Cr04+Pb(C2H302)2 = PbCr04 + 2KC2H302; 
 
 (d) 2KC2H3O2+H2SO4 =K2S04 + 2HC2H302; 
 
 (e) Pb(C2H302)2+H2S04 =PbS04 + 2HC2H302. 
 
 To the solution of the salt an excess of lead acetate solu- 
 tion is added; this causes a precipitation of the acid as a 
 lead salt and converts the alkali into an acetate which remains 
 in solution. (See equation (c)). The excess of lead acetate 
 is also in solution. The mixture is filtered and the filtrate 
 treated with a slight excess of sulphuric acid. This converts 
 the alkali acetate into a sulphate (see equation (d)) and 
 removes the lead acetate by precipitation in the form of lead 
 sulphate (see equation (e)) which is filtered out, and the 
 solution of alkali sulphate treated as above described. 
 
 Estimation of tlie Salts of the Alkali Earths 
 
 Standard solution of hydrochloric or of nitric acid is 
 preferred by many operators for the titration of hydroxids 
 or carbonates of the alkali earths. 
 
 These acids possess the advantage over most other acids 
 of forming soluble salts. The hydroxids may be estimated 
 by any of the indicators, but as they readily absorb CO2 out 
 of the air they are generally contaminated with more or less 
 carbonate, and the residual method should be used, i.e., a 
 known excess of standard acid should be added, the mixture 
 
ANALYSIS BY NEUTRALIZATION 85 
 
 boiled to expel any trace of CO2, and titrated with standard 
 alkali. 
 
 The carbonates are of course estimated in the same way, 
 as are also the organic salts of the alkali earths, after ignition. 
 As an example: 
 
 One gram of calcium carbonate is mixed with 5 cc. of 
 water. A measured excess of normal hydrochloric acid V.S. 
 is now added, and the solution boiled to drive off the CO2. 
 Then add a few drops of phenolphthalein, and titrate with 
 
 N 
 
 — alkali \^S. until a faint pink color is obtained. 
 
 N 
 Note the quantity of — alkali used, and deduct this from 
 
 N 
 the quantity of — acid first added, and the remainder will 
 
 represent the amount of acid which combined with the calcium. 
 
 N 
 Each cc. of — acid V.S. represents 0.05 gm. of CaCO:i. 
 
 CaCO.3 + 2HCI = CaCls + H2O + CO3. 
 
 2)100 2)72^ j^ 
 
 50 gms. 36.46 gms, or 1000 cc — acid V.S. 
 
 N 
 Assuming that 30 cc. of — HCl V.S. were added to the i gm. 
 
 N 
 of CaCOs, and that 11 cc. of — KOH V.S. were required 
 
 N 
 to bring the mixture back to neutrality, then 19 cc. of — HCl 
 
 were actually required to saturate the CaCOs. 
 
 Therefore 0.050X19 = 0.950 or 95 per cent. 
 
 The hydroxids and carbonates may also be estimated by 
 direct titration with standard hydrochloric acid (in the cold) 
 using methyl orange as indicator. A better plan, however, 
 would be to add the standard acid in slight excess, and then 
 standard alkali until a distinct yellow color appears; the 
 
86 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 slight excess of alkali is then determined by adding standard 
 hydrochloric acid until the red color reappears. A much more 
 distinct color reaction is thereby obtained. The quantity of 
 the standard alkali used is deducted from the total quantity 
 of standard acid added. 
 
 Soluble salts of calcium, barium and strontium, such as 
 chlorids, nitrates, etc., may be readily estimated as follows: 
 
 A weighed quantity of the salt is dissolved in water, 
 cautiously neutralized if it is acid or alkaline, phenolphthalein 
 is added, the mixture heated to boiling, and standard solution 
 of sodium carbonate delivered in from time to time, with 
 constant boiling until the red color is permanent. 
 
 This process depends upon the fact that sodium carbonate 
 forms with soluble salts of these bases insoluble neutral car- 
 bonates. 
 
 CaCla + NazCOa = CaCOg + 2NaCl. 
 Ba(N03)2 + NajCOs = BaCOs + 2NaN03. 
 
 Magnesium salts cannot be estimated in this way, as 
 magnesium carbonate affects the indicator. 
 
 The alkali earth salts may also be estimated by dissolving 
 them in water, precipitating the base as carbonate, with an 
 excess of ammonium carbonate and some free ammonia. 
 The mixture is then heated for a few minutes, and the car- 
 bonate separated by filtration, thoroughly washed with hot 
 water till all soluble matters are removed, and then titrated 
 with normal acid V.S. as carbonate. 
 
 Normal Sodium Carbonate V.S. (Na2C03 = io6) contains 
 53 gms. in i liter. This solution is made by dissolving 53 
 gms. of pure sodium carbonate (anhydrous) previously ignited 
 and cooled, in distilled water, and diluting to i liter at 15° 
 C. (59° F,). 
 
 If a pure salt is not at hand the solution may be made as 
 follows : 
 
ANALYSIS BY NEUTRALIZATION 87 
 
 About 85 gms. of pure sodium bicarbonate, free from 
 thiosulphate, chlorid, etc., are heated to dull redness (not to 
 fusion) for about fifteen minutes to expel one half of the CO2; 
 it is then cooled under a desiccator. 'WTien cool, weigh off 
 53 gms. and dissolve it in distilled water to make i liter at 
 
 15° C. (59° F.). This solution should neutralize — acid V.S. 
 
 I 
 
 volume for volume. 
 
 As an example of the process: Take of calcium chlorid 
 
 one gram, dissolve it in a small quantity of water, neutralize 
 
 the solution if it is acid or alkaline, heat to boiling, add a 
 
 N 
 few drops of phenolphthalein, and titrate with — sodium 
 
 carbonate, delivered cautiously while boiling until the red 
 color is permanent. 
 
 CaCla + NaoCOs = CaCOs + 2NaCl. 
 
 56.5 gms. 53 gms. or 1000 cc. — V.S. 
 
 N 
 Each cc. of — Na2C03 V.S. represents 0.0565 gm. of 
 
 CaCla- If 17 cc. are used the salt contains 0.0565 gm. X17 
 = 0.96 gm. or 96 per cent. 
 
 In the other method in which an excess of ammonium 
 carbonate is added together with some free ammonia, the 
 calcium is precipitated as carbonate; this is then separated 
 by filtration, thoroughly washed with hot water to remove all 
 soluble matters, and then titrated as directed for carbonate. 
 
 CaBra = CaCOz = H2SO4. 
 
 2 )198-52 2 )q9-.^.s 2 )98 j^ 
 
 99.26 gms. 49-675 gms. 49 gms. or 1000 cc. — V.S. 
 
 N 
 Each cc. of — acid thus represents 0.09926 gm. of CaBr2. 
 
88 
 
 THE ESSENTIALS OF VOLUIVIETRIC ANALYSIS 
 
 The Estimation of Mixed Hydroxids and Carbonates of 
 Alkali Earths. This may be done as described under esti- 
 mation of mixed alkali hydroxids and carbonates, page 73, 
 except that in this case it is unnecessary to precipitate the 
 carbonate by barium chlorid in that the alkali earth carbonates 
 are already insoluble. 
 
 Acidimetry 
 
 The Estimation of Acids by Neutralization. In the preceding 
 pages it has been shown how alkalies are estimated by the 
 
 1,: 
 
 t=--^ 
 
 Fig. 35. Fig. 36. 
 
 use of acid solutions of known neutralizing power. In the 
 estimation of acids, which will now be described, the order 
 is reversed, alkaline solutions of known power being used 
 in determinmg the strength of acids and of acid salts. Thus 
 the procedure is analogous to that of the alkalimetric methods. 
 
ANALYSIS BY NEUTRALIZATION .S9 
 
 The choice of the indicator, whether litmus, phenolphthalein, 
 or methyl orange, depends upon the particular acid to be 
 estimated. Phenolphthalein is employed for the organic acids 
 and boric acid and is preferred for phosphoric acid; while 
 methyl orange and litmus are usually employed in the titration 
 of the inorganic acids. 
 
 The standard alkali used may be either an hydroxid or 
 a carbonate, the former is, however, usually preferred, because 
 the carbonate when brought in contact with an acid gives 
 off carbonic acid gas (CO2) which interferes to a great extent 
 with most indicators. On the other hand, it must be remem- 
 bered that the alkali hydroxids are very prone to absorb carbon 
 dioxid from the atmosphere, therefore their solutions often 
 contain some carbonate, the presence of which even in small 
 quantities will occasion errors when used with most indicators, 
 especially with litmus and phenolphthalein. It is therefore 
 advisable, when using these indicators or others which arc 
 affected by carbon dioxid, to employ gentle heat toward the 
 close of each titration, in order to drive off the liberated gas. 
 Methyl orange is not affected by this gas, and therefore heating 
 is not necessary when this indicator is used. In fact, it is 
 imperative that heat should not be employed with this 
 indicator. 
 
 In acidimetrical operations when methyl orange is used 
 as indicator, residual titrations may be advantageously done, 
 because the change of color from yellow to red which is 
 brought about by the acid is much more readily seen than 
 that from red to yellow. 
 
 In the U. S. P. standard solutions of both potassium and 
 sodium hydroxid are official. The former, however, is pref- 
 erable, because it attacks glass more slowly and less energetically, 
 and also foams much less than does the sodium hydroxid 
 solution. The neutralizing power of each is, however, the 
 
90 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 same. Standard solutions of alkali hydroxid should be pre- 
 served in small vials, provided with well-fitting rubber stoppers, 
 or better still they should be provided with tubes filled with 
 a mixture of soda and lime, which absorbs CO2 and prevents 
 its access to the solution. A vessel of this description is 
 illustrated in Fig. 35. 
 
 An improvement upon this is shown in Fig. 36, since it 
 allows of the burette being filled without removing the stopper, 
 and consequently without any access of CO2 whatever. 
 
 Where a series of titrations of the same kind have to be 
 made with the same alkali standard solution, the arrangement 
 shown in Fig. 9 may be used, both the reservoir and the 
 burette in this case being provided with soda-lime tubes. 
 
 Preparation of Standard Alkali Solutions 
 
 N 
 , Normal Potassium Hydroxid (KOH = 56.i; — V.S. = 56.1 
 
 gms. in 1000 cc). Potassium hydroxid being prone to absorb 
 carbon dioxid out of the air the pure article is not readily 
 obtained in commerce. If pure potassium hydroxid were 
 easily obtained it would only be necessary to dissolve 56.1 
 gms. in sufficient water to make 1000 cc. But since it always 
 contains some CO2 and H2O, it is necessary to take a slight 
 excess and dilute the solution to the proper volume after having 
 determined its strength. 
 
 The standardization may be effected by means of any of 
 the standard acid solutions. 
 
 A satisfactory method for the preparation and standard- 
 ization of this solution is as follows: 
 
 Dissolve 75 gms. of potassium hydroxid in sufficient water 
 to make about 1050 cc. at 15° C. (59° F.), and fill a burette 
 with a portion of this solution. 
 
 Dissolve 0.63 gm. of pure oxalic acid in about 10 cc. of 
 
ANALYSIS BY NEUTRALIZATION 91 
 
 water in a beaker or flask, add a few drops of phenolphthalein 
 T.S., and then carefully add from the burette the potassium 
 hydroxid solution, agitating frequently and regulating the flow 
 to drops towards the end of the operation until a permanent 
 pale pink color is obtained. Note the number of cc. of the 
 alkali solution consumed, and then dilute the remainder so 
 that exactly lo cc. of the diluted liquid will be required to 
 neutralize 0.63 gm. of oxalic acid. Instead of weighing off 
 0.63 gm. of the acid, 10 cc. of its normal solution may be 
 used. 
 
 Example. Assuming that 8 cc. of the stronger potassium 
 hydroxid solution had been consumed in the trial, then each 
 8 cc. must be diluted to 10 cc, or the whole or the remaining 
 solution in the same proportion. Thus if 8 cc. must be diluted 
 to 10 cc, 1000 cc. must be diluted to 1250' cc. 
 
 8:10: : 1000: :*; 31;= 1250 cc. 
 
 It is always advisable to make another trial after diluting. 
 10 cc. should then neutralize 0.63 gm. of pure oxalic acid. 
 
 Standardization by Means of Potassium Bitartvate. 
 This method is based upon the reaction 
 
 KHC4H4O6 + KOH = K2C4H4O6 + H2O. 
 
 188.: 4 56.1 
 
 N 
 1000 cc. of — KOH contains 56.1 gms. of KOH and will 
 
 react with 188.14 gms. of potassium bitartrate. 25 cc. of 
 
 — KOH will therefore react with 4.7035 gms. of potassium 
 I 
 
 bitartrate. 
 
 A solution of potassium hydroxid, 75 gms. in 1050 cc, is 
 
 prepared and titrated against pure potassium bitartrate, 
 
 using phenolphthalein as indicator. 
 
92 TPIE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 4.7035 gms. of purified dry potassium bitartrate * are 
 dissolved in 100 cc. of boiling distilled water, a few drops (3) 
 of phenolphthalein are added, and the solution of potassium 
 hydroxid run into it (the solution being frequently boiled) 
 until a permanent pale pink color appears. Exactly 25 cc. 
 will be required if the alkali solution is normal. If only 23 
 cc. are consumed, then each 23 cc. must be diluted to 25 cc, 
 or the whole of the remaining solution in the same proportion. 
 
 Staiiflardizafioit by Menus of Potasstuni Bl-iodate.\ 
 Potassium bi-iodate is an acid salt and may be directly 
 titrated with potassium hydroxid, using phenolphthalein as 
 indicator. 
 
 One molecule of the bi-iodate is equivalent to one molecule 
 of potassium hydroxid, as shown by the equation, 
 
 KH(I03)2 + KOH = 2KI03 + HaO. 
 
 389.94 56.1 
 
 To standardize a potassium-hydroxid solution, weigh ac- 
 curately 3.8994 gms. of potassium bi-iodate, dissolve it in 
 about 25 cc. of water, add a few drops of phenolphthalein, 
 and then run into this, from a burette, the hydroxid solution 
 
 * Purified potassium bitartrate for standardizing caustic alkali volumetric 
 solutions may be obtained as follows: loo gms. of the salt are placed in a 
 beaker, together with 85 cc. of water and 25 cc. of 10 per cent hydrochloric 
 acid, the beaker is covered and heated on a boiling water bath, stirring 
 occasionally for three hours. The liquid is then quickly cooled, decanted, 
 and the residue washed first by decanlation with 100 cc. of cold water, then 
 again washed after it has been transferred to a plain filler, using cold water 
 until the filtrate ceases to become opalescent when acidified with a few drops 
 of nitric acid upon the addition of silver nitrate solution. The precipitate is 
 then dissolved in the smallest possible quantity of boiling water (about 1500 
 cc.) filtered, and the filtrate stirred constantly while cooling. When cold, 
 the crystalline precipitate is collected upon a filter washed with 300 cc. of 
 cold water, run through it in small portions at a time. It is then allowed 
 to drain, and finally dried in an air oven at 120° C. until its weight is constant. 
 
 t See Meinecke, Chem. Ztg., XIX. 2. Also Caspari, Proc. A. Ph. A., 
 i9°4, 389- 
 
ANALYSIS BY NEUTRALIZATION 93 
 
 which is to be standardized, until a pale pink color appears. 
 Note the number of cc. used and dilute the solution so that 
 exactly lo cc. of it will neutralize 3.8994 gms. of the bi-iodate. 
 
 Example. Assuming that 8.2 cc. had been consumed, 
 then each 8.2 cc. must be diluted to 10 cc, or the whole of 
 the remaining solution in the same proportion. 
 
 The advantages of this salt as an ultimate standard are 
 (i) that it may be procured in the market in a state of absolute 
 purity; * (2) that it is permanent, being neither deliquescent 
 nor efflorescent; (3) that it can be dried at 110° C. without 
 decomposition; (4) that the results obtained with it are quite 
 accurate, and (5) that it may be employed for standardizing 
 most of the volumetric solutions commonly found in the 
 laboratory. 
 
 Stayidartlisatlon by Means of Normal Acid I'.S. 20 CC. 
 of a strictly normal acid V.S. are placed into a beaker, 2 
 drops of phenolphthalein solution are added and the potas- 
 sium hydroxid solution delivered into it until the liquid just 
 turns pink and remains so after boiling. If the alkali hydroxid 
 solution is strictly normal, there will be consumed exactly 
 20 cc. If less is consumed the solution is too strong and 
 must be so diluted with distilled water that equal volumes 
 of it and the normal acid will exactly neutralize each other. 
 Thus if 18 cc. of the alkali are consumed, then each 18 cc. 
 
 must be diluted to 20 cc. 
 
 N 
 Normal Sodium Hydroxid (NaOH = 4o; — V.S. =40 gms. 
 
 * According lo Caspari, the salt may be readily prepared as follows: See 
 A. Ph. A., 1904, 390. Potassium bicarbonate is mixed in solution with an 
 equivalent amount of iodic acid, and to the neutral solution is added an 
 amount of iodic acid equal to the quantity first used. The solution is then 
 evaporated until crystallization begins, and the first crop of crystals rejected. 
 Those which separate after the solution has cooled to 50° C. are almost pure 
 and will be rendered absolutely pure if recrystallized. 
 
94 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 in looo cc). Dissolve 54 gms. of sodium hydroxid in enough 
 water to make about 1050 cc. of solution, fill a burette with 
 a portion of this, and check it with normal acid, or a weighed 
 quantity of oxalic acid or potassium bitartrate, in exactly the 
 same manner as described for normal potassium hydroxid. 
 Other strengths of standard alkali V.S. are Half-normal 
 
 Fifth-normal 1^1, Tenth-normal I — I, Twentieth- 
 
 /N\ /N\ /N 
 
 normal — , Fiftieth-normal — , Hundreth-normal ( — 
 \2o/ Vso/' Vioo, 
 
 These are all prepared by properly diluting the normal V.S. 
 
 and then checking the strength of the product. 
 
 Other standard alkali solutions in frequent use are normal 
 
 sodium carbonate, normal and other strengths of ammonia, 
 
 and decinormal barium hydroxid. 
 
 Estimation of the Inorganic Acids 
 
 To weigh off directly a definite quantity of a fluid acid, is 
 not a very easy matter. It is always a better plan to measure 
 a small quantity of the acid and weigh it accurately in a tared 
 and stoppered weighing flask (Fig. 37), then to add 
 water and titrate with the standard alkali in the 
 presence of a suitable indicator. If the specific gravity 
 of the acid is known or can be easily taken, it is 
 sufficient to measure a certain quantity of it by 
 means of a pipette, and then determine its weight 
 by multiplying the volume in cubic centimeters by 
 the specific gravity. It must be remembered, however, that 
 the liquid must be measured at the same temperature at 
 which the specific gravity was taken. This method is 
 applicable to the diluted acids as well as to the concentrated 
 acids of commerce, as hydrochloric, nitric and sulphuric. 
 
ANALA'SIS BY NEUTRALIZATION 
 
 !)5 
 
 In the case of very volatile acids, i.e., such as evolve acid 
 vapors at ordinary temperatures, the determination of the 
 weight by means of the specific gravity is inadmissable. Such 
 acids should be weighed in a Lunge pipette. Fig. 38, or in a 
 simple bulb pipette provided with a glass stop-cock, Fig. 39, 
 or in a Grethan's pipette. Fig. 40. 
 
 The Lunge pipette is used by producing a vacuum in 
 the bulb (a), the air-tight glass mantle (c) is then removed. 
 
 Fig. 39. 
 
 Fig. 40. 
 
 and the tip of the tube (d) sunk into the acid which is drawn 
 up into the bulb, upon opening the cock (i); when sufficient 
 of the acid has been drawn into the apparatus the cock is 
 closed, the tip of the pipette wiped, the glass mantle put in 
 place, and the whole weighed. The weight of the empty 
 pipette deducted gives the weight of the acid taken up. The 
 pipettes shown in Figs. 39 and 40 are filled by applying direct 
 suction with the lips, the operator protecting himself against 
 inhalation of harmful vapors by attaching an absorption tube 
 containing soda-lime, caustic soda, or similar substance. 
 
96 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The quantity of acid to be taken (in most cases) should 
 be such as will require for neutralization from 20 to 50 cc. 
 of the standard alkali. In the case of concentrated inorganic 
 acids, 2 or 3 gms. may be taken, while in the case of the 
 dilute acids, from 6 to 8 gms. 
 
 Any of the indicators may be employed for the inorganic 
 acids, but because of the usual presence of carbonate in the 
 standard alkali, methyl orange is preferred. 
 
 Hydrochloric Acid (HC1 = 36.468). About 2 cc. of hydro- 
 chloric acid (sp.gr. 1.048) are introduced into a tared weighing 
 flask and its weight accurately taken. (The weight is found 
 to be 2.098 gms.) 20 cc. of water are now added, followed 
 by two drops of methyl orange, and the solution carefully 
 titrated with normal potassium hydroxid until the reddish 
 color of the solution is changed to yellow. 
 
 Assuming that 18.4 cc. were required, then 18.4 cc. X 
 0.036468 gm. =0.671 gm. of absolute hydrochloric acid in the 
 2.098 gms. taken. 
 
 To find the per cent apply the proportion 
 
 2.098 gms. : 0.671 gm. : : 100: x. ^ = 31. 9 per cent. 
 
 The equation is: 
 
 HCl + KOH = KCl + H2O. 
 
 N 
 36.468gms,= 56.i gms. = 1000 cc. — V.S. 
 
 N 
 .036468 gm. = I cc. — V.S. 
 
 Sulphuric Acid (112804 = 98.07). About i cc. of the 
 concentrated acid is weighed in a tared weighing flask and 
 found to weigh 1.8 gms. 20 cc. of water are added and then 
 2 drops of methyl orange, and the titration with normal potas- 
 sium hydroxid begun, and cautiously continued until the 
 yellowish color of the solution indicates the completion of 
 
ANAL\'SIS BY NEUTRALIZATION 97 
 
 the operation. Note the number of cc. of alkali solution used 
 and apply the equation 
 
 H2SO4 + 2KOH = K2SO4 + 2H2O. 
 
 2 )98.07 2)112.2 
 
 49'°3 gms. = 56.1 gms. = 1000 cc. — V.S. 
 
 Thus each cc. of normal KOH V.S. represents 0.04903 
 gm. of pure H2SO4. 
 
 Phosphoric Acid (H3P04= 98.064). In the assay of phos- 
 phoric acid by direct neutralization with standard KOH, the 
 acid is converted into, first, KH2PO4, then K0HPO4, and finally 
 into the normal K3PO4. We have no indicator which reliably 
 shows the completion of the neutralization, i.e., the formation 
 of the tribasic K3PO4. Litmus cannot be used as indicator 
 here for the dipotassic or disodichydric phosphate (K2HPO4 
 or Na2HP04) which is formed is slightly alkaline towards 
 litmus; the same is true of most other indicators. 
 
 It is recommended, therefore, in order to estimate phos- 
 phoric acid alkalimetrically, to prevent the formation of soluble 
 phosphate of the alkali, and to bring the acid into a definite 
 compound with an alkali earth, as follows : 
 
 The free acid in a diluted state is placed in a flask and 
 
 a known volume of normal alkali in excess added in order 
 
 to convert the whole of the acid in a basic salt. A few drops 
 
 of rosolic acid are now added, and sufficient neutral BaCl2 
 
 solution poured in to combine with the phosphoric acid. The 
 
 mixture is heated to boiling, and while hot the excess of alkali 
 
 N 
 is titrated with — acid. 
 
 The suspended basic phosphate, together with the liquid, 
 possesses a rose-red color until the last drop or two of acid, 
 after continuous heating and agitation, gives a permanent 
 white or slightly yellowish milky appearance, when the process 
 is ended. 
 
98 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The volume of normal alkali, less the volume of normal 
 acid, represents the amount of alkali required to convert the 
 phosphoric acid into a normal trisodic or tripotassic phosphate. 
 
 H3PO4 + 3KOH = K3PO4 + 3H2O. 
 
 3 )98-°64 3 )i68-3 j>^ 
 
 32.688 gms. 56.1 gms. = 1000 cc. of — KOH V.S. 
 
 N 
 
 Thus I cc. of — alkali = 0.032688 gm. of H3PO4. 
 
 Thompson, however, has demonstrated that this acid may 
 be accurately titrated by standard alkali when using either 
 methyl orange or phenolphthalein, or both, successively. 
 
 If methyl orange is used the color changes upon the 
 completion of the formation of monobasic phosphate, KH2PO4, 
 as per the following equation : 
 
 H3PO4 + KOH = KH2PO4 + H2O. 
 
 If phenolphthalein is used, the color changes upon the 
 completion of the formation of the dibasic phosphate, K2HPO4, 
 
 H3PO4 + 2KOH = K2HPO4 + 2H2O. 
 
 .45 an example: A weighed quantity of the acid is diluted 
 with water to measure 20 cc. and sufficient pure sodium chlorid 
 added to saturate the solution. Four drops of methyl orange 
 
 N 
 are then introduced and the titration with — KOH begun 
 
 and continued until the red color changes to yellow, indicating 
 the formation of the monobasic phosphate 
 
 H3PO4 + KOH = KH2PO4 + H2O. 
 
 N 
 98.064 gms. 56.1 gms. = 1000 cc. — V.S. 
 
 N 
 Each cc. of - KOH V.S. = 0.098064 gm. of H3PO4. 
 
 The use of sodium chlorid in this assay is to decrease the 
 ionization of the acid salts produced in the reaction. 
 
ANALYSIS BY NEUTRALIZATION 99 
 
 Another portion of the acid is treated in like manner, 
 adding sodium chiorid and titrating, but using phenolphthalein 
 as the indicator. The titration is continued until a faint 
 permanent pink color appears. It is advisable to use heat, 
 or better still, a standard alkali solution which is quite free 
 from CO2. The end-reaction in this case marks the formation 
 of the dibasic phosphate 
 
 H3PO4 + 2KOH = K2HPO4 + 2H2O. 
 
 2)98.064 2)112.2 
 
 49.032 gms. 56.1 gms. = 1000 cc. — V.S. 
 
 N 
 Each cc. of — KOH V.S. =0.049032 gm. of H3PO4. 
 
 N 
 Just twice as much of the — KOH V.S. will be taken 
 
 in this assay as in the foregoing. 
 
 The two assays may be combined as follows: A weighed 
 quantity of the acid ' is diluted with water saturated with 
 sodium chiorid and titrated with the normal alkali V.S. 
 using methyl orange as indicator until the red color of the 
 solution changes to yellow. The number of cc. is noted and 
 multiplied by 0.098064 gm. A few drops of phenolphthalein 
 solution are now added and the titration continued until a 
 pale-red color appears. The total number of cc. of normal 
 alkali used in the double titration is then multiplied by 
 0.049032 gm. 
 
 Hypophosphorous Acid (HPH202 = 66.064). 
 
 HPH2O2 -h KOH = KPH2O2 + H2O. 
 
 N 
 Each cc. of — alkali represents 0.066064 gm. of HPH2O2. 
 
 Nitric Acid (HN03 = 63.01). 
 
 HNO.3 4- KOH = KNO3 + H2O. 
 
 N 
 Each cc. of — alkali represents 0.06301 gm. of HNO3. 
 
100 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Hydrobromic and hydriodic acids may be estimated in 
 the same way as the foregoing, but it is usually preferred to 
 estimate them by precipitation analysis. Sulphurous acid is 
 best assayed by oxidation with iodin. 
 
 Boric Acid (H3B03 = 62). This acid is estimated by 
 
 N 
 neutralization with — NaOH in the presence of a large 
 
 quantity of glycerin. (Thompson's Method, J. S. C. I., XII,, 
 432). The addition of sufficient glycerin to a boric acid 
 solution, so that no less than 30 per cent be present throughout 
 the titration, develops the acidity of boric acid with regard 
 to phenolphthalein to a great degree, and enables one to 
 titrate direct with standard soda solution, i gm. of boric 
 acid is dissolved in 50 cc. of water; to this is added an equal 
 volume of glycerin, then a few drops of phenolphthalein, and 
 the titration with normal sodium hydroxid begun and con- 
 tinued until a pink color appears. 
 
 N 
 Each cc. of — NaOH = 0.062 gm. of H3BO3. 
 
 H3BO3 + NaOH = NaH2B03+H20. 
 
 N 
 62 40 gms. in 1000 cc. — V.S. 
 
 N 
 0.062 gm. = I cc. — V.S. 
 
 Estimation of the Organic Acids 
 
 As the individual organic acids require different indicators, 
 the table on page 27 should be consulted in the selection of 
 an indicator for a particular organic acid. Phenolphthalein 
 is, however, the most suitable for organic acids generally. 
 
 Acetic Acid (HC2H302 = 6o). Mix 3 gms. of the acid 
 with a small quantity of water, add a few drops of phenol- 
 phthalein T.S., and titrate with normal potassium hydroxid 
 
ANALYSIS BY NEUTRALIZATION 101 
 
 V.S. until a permanent pale pink color is obtained, and apply 
 the following equation: 
 
 HC2H3O2 + KOH - KC2H3O3 + H2O. 
 
 60 56.1 
 
 N 
 Thus 1000 cc. of — KOH V.S. will neutralize 60 gms. 
 
 N 
 of acetic acid; therefore each cc. of — KOH V.S. represents 
 
 0.060 gm. of acetic acid. 
 
 If 18 cc. are required to neutralize 3 gms. of the acid, it 
 contains 18X0.06=1.08 gms. of absolute acetic acid. " 
 
 1.08 X TOO 
 
 = 36 per cent. 
 
 Tartaric Acid (H2C4li406=iSo). Dissolve 2 gms. of tar- 
 taric acid in sufficient water to make a solution, add a few 
 drops of phenolphthalein and then pass into the solution 
 
 N 
 from a burette — potassium hydroxid V.S. until a faint pink 
 
 tint is acquired by the solution, and apply the equation 
 H2C4H4O6 + 2KOH = K2C4H4O6 + 2H2O. 
 
 75 gms. = 1000 cc— V.S. 
 
 The other organic acids are assayed in exactly the same 
 manner as that described for the foregoing. 
 Citric Acid (H3C6H507=i92), 
 
 HaCeHsOr + 3KOH = K3C6H5O7 + 4H2O. 
 
 3)192 N 
 
 64 gms. = 1000 cc. — V.S. 
 
 N 
 Each cc. of — KOH represents 0.064 g^- oi citric acid. 
 
102 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 TABLE SHOWING QUANTITY OF SUBSTANCE TO BETAKEN FOR 
 ANALYSIS IN DIRECT PERCENTAGE ESTIMATIONS. 
 
 Molec- 
 ular 
 Weight. 
 
 Quantity to be 
 taken so that 
 
 each cc. of — 
 I 
 V.S. will rep- 
 resent 1%. 
 
 Percentage 
 strength of 
 
 Official 
 Substance. 
 
 Acid, acetic, HC^HjOj 
 
 boric, H3BO3 
 
 citric, HsCoHsOj-FHjO 
 
 hydrobromic, HBr. Dil 
 
 hydrochloric, HCl 
 
 hydriodic, HI. Dil 
 
 hypophosphorous, HPH^Oj 
 
 lactic, HC3II5O3 
 
 nitric, HNO3 
 
 oxalic, HjC204 4- 2H2O 
 
 phosphoric, H3PO4 
 
 phosphoric with methyl orange 
 
 phosphoric with phenolphthalein 
 
 sulphuric, HjSOj 
 
 tartaric, HjC^H^Oj 
 
 trichloracetic, HC2CI3O2 
 
 Ammonium carbonate, N3H11C2O5 
 
 Ammonia water, NH3 
 
 Ammonia water, stronger, NH3 
 
 Lime water, Ca(OH)j 
 
 Lithium carbonate, LijCOs 
 
 " citrate, Li3CeH50,+4H20 
 
 Potassium acetate, KC2H3O2 
 
 bicarbonate, KHCO3 
 
 bitartrate, KHCiH^O,, 
 
 carbonate, K2CO3 
 
 citrate, K3CoH50,+H20 
 
 hydroxid, KOH 
 
 hydroxid, liquor, KOH 
 
 sodium tartrate 
 
 KNaC4H,0„+4H20 
 
 Sodium acetate, NaC2H302+3H20 
 
 " benzoate, NaCjHsOj 
 
 " bicarbonate, NaHC03 
 
 " carbonate, Na2C03 
 
 5o.o 
 62,0 
 
 210.0 
 80.92 
 36.46 
 
 127.92 
 66.04 
 go.u 
 63.01 
 
 r26.o 
 98.04 
 
 98.07 
 150.0 
 163.38 
 157-11 
 17.01 
 17.01 
 74.09 
 73.88 
 281 .92 
 98.1 
 100. 1 
 r88.i 
 138.2 
 
 324-3 
 56.1 
 56.1 
 
 2S2.1 
 136.0 
 144.0 
 84.0 
 106.0 
 
 6.0 
 6.2 
 
 7-" 
 8.092 
 3.646 
 12.792 
 6.604 
 9.0 
 6.301 
 6.3 
 
 gms. 
 
 9.804 
 
 4.902 
 
 4-9°3 
 
 7-5 
 
 1-6338 
 
 5-25 
 
 1-7 
 
 1-7 
 
 3-7°4 
 
 3-69 
 
 9-397 
 
 9.81 
 10.01 
 18.81 
 
 6.gi 
 10.81 
 
 5. 61 
 
 5-61 
 
 14.10 
 
 13.6 
 
 14.4 
 
 8.4 
 
 5-3 
 
 36 
 99 
 99 
 10 
 
 31 
 10 
 
 30 
 75 
 68 
 
 85- 
 
 99 
 99 
 98 
 
 99 
 
 85 
 
 5 
 
 99 
 
 99 
 99 
 99 
 85 
 
 14 
 
 S 
 
 s 
 
ANALYSIS BY NEUTRALIZATION 103 
 
 Oxalic Acid (H2C204.2H20=i26). 
 
 H2C2O4.2H2O + 2KOH = K2C2O4 + 4H2O. 
 
 63 gms. = 1000 cc. — V.S. 
 
 N 
 Each cc. of — KOH represents 0.063 ?>^- of crystallized 
 
 oxalic acid. 
 
 Lactic Acid (HC3H503 = go). 
 
 HC3H5O3 + KOH = KC3H5O3 + H2O. 
 
 N 
 90 56.1 gms. = 1000 cc. — V.S. 
 
 N 
 Each cc. of — KOH represents 0.090 gm. of lactic acid. 
 
 Referring to the table it will be seen that if the quantities 
 indicated are taken for analysis, the amount of standard 
 solution required for substances of high percentage strength 
 will be very large (in some cases over 99 cc), while for sub- 
 stances of low percentage strength, as for instance lime water, 
 so small a volume of standard solution is required as to be 
 unreadable (0.14 cc). It is therefore advisable to take for 
 analysis a smaller quantity of high percentage substances and 
 a larger quantity of such substances as contain a low per- 
 centage. It is usually best to so adjust it that no less than 
 10 nor more than 30 cc. of the standard solution be required. 
 For example: In the case of citric acid, instead of taking 
 for analysis 7 gms. it will be better to take one-fourth of this 
 quantity, then each cc. of the standard solution used will 
 represent 4 per cent, and only one-fourth as much will be 
 required, i.e., 24.9 cc. instead of 99.5 cc. Again, in the case 
 of lime water, if 37.04 gms. are taken instead of 3.704 gms., 
 
 N 
 1.4 cc. of the — standard solution will be required, which 
 
 is better than 0.14 cc, but in this case it will be still better 
 
104 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 to use a decinormal I — I solution, then 37.04 gms. of lime 
 
 N 
 water would require for neutralization just 14 cc. of the — 
 
 acid V.S. If half the quantity indicated in the table is taken, 
 then each cc. of the standard solution will represent 2 per 
 cent. If one-tenth the quantity is taken each cc. will repre- 
 sent TO per cent. If double the quantity is taken each cc. 
 will represent 0.5 per cent", etc. 
 
 ESTIMATION OF ACIDS IN COMBINATION IN NEUTRAL SALTS 
 
 This may be done in the case of a large number of salts, 
 by adding to the solution of the salt a measured excess of 
 alkali or alkali carbonate in the form of normal solution, and 
 then ascertaining the excess by retitration with normal acid. 
 Thus the amount of alkali which went into combination with 
 the acid is obtained. Most bases are precipitated by the 
 hydroxid; some, however, require the addition of carbonate 
 to effect their precipitation. 
 
 The carbonate is required for alkali earth salts, mag- 
 nesium salts, alum, zinc salts, bismuth salts, nickel, cobalt, 
 and lead salts. 
 
 . Example. 2 gms. of barium chlorid are dissolved in water 
 and sufficient normal sodium carbonate added to make the 
 liquid decidedly alkaline (say 20 cc), and the whole diluted 
 to 300 cc. and set aside to settle. 100 cc. of the clear super- 
 natant liquid are then removed with a pipette and titrated 
 for excess of alkali with normal nitric acid or normal hydro- 
 chloric acid, of which say 1.2 cc. are required, making it 3.6 
 cc. for the whole quantity; therefore 20 — 3.6=^16.4 cc. is the 
 measure of the alkali which combined with the acid of the 
 orginal salt. This multiplied by 0.03546, the factor for chlorin, 
 gives 0.57695 gm. of chlorin. 
 
CHAPTER IX' 
 
 ANALYSIS BY PRECIPITATION 
 
 The general principle of this method is that the deter- 
 mination of the quantity of a given substance is effected by 
 the formation of a precipitate, upon the addition of the 
 standard solution to the substance under examination. There 
 are three ways of determining the end-reaction in precipi- 
 tation analyses: 
 
 1. By adding the standard solution until it ceases to produce 
 any more precipitate, as in the estimation of silver by standard 
 sodium chlorid, and the estimation of haloid salts and acids 
 by means of standard silver nitrate. The application of this 
 ending is almost limited to the above estimations, because 
 in these only can accurate results be obtained. The silver 
 halids formed are not only quite insoluble, but they have a 
 tendency to curdle closely upon shaking (especially in acid 
 solutions), and thus leave a clear supernatant liquid in which 
 any further precipitation can readily be seen. Most of the 
 other precipitates, such as barium sulphate, calcium oxalate, 
 etc., although heavy and insoluble, are so finely divided and 
 powdery that they do not readily subside. 
 
 2. By the use of an indicator, as in the estimation of haloid 
 salts by means of standard silver nitrate solution, using neutral 
 potassmm chromate as the indicator. The latter is added to 
 the haloid solution, and the silver nitrate V.S. delivered into 
 the mixture until a permanent rod color (silver chromate) 
 is produced. Silver nitrate reacts by preference with the 
 halogen, and does not react with the chromate until the halogen 
 
 105 
 
106 THE p:ssentials of volumetric analysis 
 
 has been entirely precipitated. Hence the production of a 
 permanent red color in the precipitate marks the completion 
 of the precipitation of the halogen. 
 
 Another illustration is in the estimation of silver by sul- 
 phocyanate solution, using ferric alum as indicator. The 
 sulphocyanate produces with the silver a white precipitate 
 of silver sulphocyanate, but when the precipitation of silver 
 is complete the sulphocyanate reacts with the ferric alum 
 present and a red ferric sulphocyanate appears and marks the 
 end-point. On the other hand, the indicator may be used 
 externally, i.e., alongside of the liquid being analyzed, a drop 
 of the latter being brought in contact with a drop of the indi- 
 cator at frequent intervals in the course of the titration, as in 
 the estimation of phosphoric acid by means of uranium nitrate 
 solution, in which potassium ferrocyanide is used as indicator. 
 
 3. By adding the standard solution until the first appearance 
 
 of a precipitate, as in the estimation of cyanogen by silver 
 
 nitrate solution, and the estimation of chlorin by mercuric 
 
 nitrate V.S. In these estimations the standard solution is 
 
 added to the solution of the substance under analysis until a 
 
 precipitate appears. 
 
 /N\ 
 Preparation of Decinormal I — I Silver Nitrate (AgNOa 
 
 N 
 = 169.89; — V.S. = 16.989 gms. in 1000 cc). Dissolve 16.989 
 
 gms. of pure silver nitrate * in sufficient water to make, at 
 or near 15° C. (59° F.), exactly 1000 cc. One liter of this solu- 
 tion thus contains xV of the molecular weight in grams of 
 silver nitrate. It is therefore a decinormal solution. 
 
 If pure crystals of silver nitrate are not readily obtainable, 
 and pure sodium chlorid is at hand, a solution of the silver 
 
 * This should be pulverized and dried at 120° C. for half an hour in a 
 covered crucible before weighing. 
 
ANALYSIS BY PRECIPITATION 107 
 
 nitrate may be made of approximate strength, a little stronger 
 than necessary, and then standardized by means of the sodium 
 chlorid, as follows: 0.11692 gm. of sodium chlorid is'dissolved 
 in water, and a burette filled with the solution of silver 
 nitrate to be standardized. The silver solution is now slowly 
 added from the burette to the sodium chlorid solution con- 
 tained in a beaker until no more precipitate of silver chlorid 
 is produced. 
 
 If neutral potassium chromate is used as an indicator, 
 the end of the reaction is shown by the appearance of yellowish- 
 red silver chromate. This indication is extremely delicate. 
 The silver nitrate does not act upon the chromate until all 
 of the chlorid is converted into silver chlorid. 
 
 In the above reaction 20. cc of silver nitrate should be 
 required. But since the silver-nitrate solution is too strong, 
 less of it wil! complete the reaction, and the solution must 
 be diluted so that exactly 20 cc. will be required to precipitate 
 the chlorin in o. 11 692 gm. of NaCl. 
 
 Thus if 17 cc. are used, each 17 cc. must be diluted to 
 20 cc, or each 170 cc. to 200 cc, or the entire remaining 
 solution in the same proportion. 
 
 After dilution a fresh trial should always be made. 
 
 Silver nitrate solution should be kept in dark amber- 
 colored, glass-stoppered bottles, carefully protected from dust. 
 
 Titration by decinormal silver nitrate V.S. may be managed 
 in various ways, adapted to the special preparation to be 
 tested. 
 
 1. In most cases it is directed to be used in the presence 
 of a small quantity of potassium chromate T.S. 
 
 2. In some cases it is added until the first appearance of 
 a permanent precipitate, as in potassium cyanid and hydro- 
 cyanic acid assays. 
 
 3. It may be used in all cases without an indicator by 
 
108 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 • 
 
 observing the exact point when no further precipitate occurs. 
 But since this consumes too much time in waiting for the 
 precipitate to subside, so as to render the supernatant hquid 
 sufficiently clear to recognize whether a further precipitate is 
 produced by the addition of the silver solution, it is imprac- 
 ticable. 
 
 4- It may be added in definite amount, known to be in 
 excess of the quantity required, and the excess measured 
 back by titration with decinormal potassium sulphocyanate 
 V.S., or even with decinormal sodium chlorid V.S. (residual 
 titration). 
 
 N N 
 
 Decinormal — Sodium Chlorid (NaCl=?8.46; — V.S. 
 10 ^ -^ lO 
 
 = 5.846 gms. in 1000 cc). Dissolve 5.846 gms. of pure 
 
 sodium chlorid in enough water to make exactly 1000 cc. 
 
 at the standard temperature. 
 
 Check this solution with decinormal silver nitrate. The 
 two solutions should correspond, volume for volume. 
 
 Pure Sodium Chlorid may be prepared by passing into ' 
 a saturated aqueous solution of the purest commercial sodium 
 chlorid a current of dry hydrochloric acid gas. The crystal- 
 line precipitate is then separated and dried at a temperature 
 sufficiently high to expel all traces of free acid. 
 
 N 
 The method of standardizing — NaCl solution is as follows : 
 
 0.33978 gm. of silver nitrate is dissolved in 10 cc. of 
 
 N 
 distilled water, and the solution carefully titrated with — 
 
 •^ ID 
 
 NaCl V.S. until precipitation ceases. 20 cc. of the standard 
 
 solution should be required. 
 
 AgNOg -t- NaCl = AgCl -t- NaNOg. 
 
 10)169.89 io):;8.46 ^ 
 
 16.989 gms. 5.846 gms., or looo cc. — NaCl V.S. 
 
ANALYSIS BY PRECIPITATION 109 
 
 Each cc. of the standard sokition represents 0.016989 gm. 
 of pure AgNOs. 
 
 0.016989X20 = 0.33978 gm. 
 
 0.33978X100 
 
 ^ — =100 per cent. 
 
 0-33978 ^ 
 
 This solution may also be standardized by residual tira- 
 
 tion with Volhard's solution. 
 
 N 
 Decinormal — Potassium Sulphocyanate (Volhard's Solu- 
 
 N 
 tion) (KSCN = 97.18; — V.S. = 9.7i8 gms. in 1000 cc). Dis- 
 solve 10 gms. of pure crystallized potassium sulphocyanate 
 (thiocyanate) in 1000 cc. of water. 
 
 This solution, which is too concentrated, must be adjusted 
 so as to correspond exactly in strength with decinormal silver 
 nitrate V.S. For this purpose introduce into a flask 20 cc. 
 
 N 
 of — AgNOs V.S., 3 cc. of ammonioferric sulphate solution, 
 
 and 5 cc. of diluted nitric acid (10 per cent and free from 
 nitrous compounds). 
 
 Dilute the Uquid with 75 cc. of distilled water, and titrate 
 it with the sulphocyanate solution. 
 
 At first a white precipitate of silver sulphocyanate is 
 produced, giving the fluid a milky appearance, and then as 
 each drop of sulphocyanate falls in it is surrounded by a deep 
 brownish-red cloud of ferric sulphocyanate, which quickly 
 disappears on shaking, as long as any of the silver nitrate 
 remains unchanged. 
 
 When the point of saturation is reached and the silvei 
 has all been precipitated, a single drop of the sulphocyanate 
 solution produces a faint brownish-red color, which does not 
 disappear on shaking. 
 
 Note the number of cc. of the sulphocyanate solution 
 
110 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 used, and dilute the whole of the remaining solution so that 
 equal volumes of this and of the decinormal silver nitrate 
 will be required to produce the permanent brownish-red tint. 
 (The same tint of brown or red to which the volumetric solution 
 is adjusted must be attained when the solution is used in 
 volumetric testing.) 
 
 Assuming that 19 cc. of the sulphocyanate solution were 
 required to produce the reaction, then each 19 cc. must be 
 diluted to make 20 cc, or the whole of the remaining solution 
 in the same proportion. 
 
 Always make a new trial after the dilution to see if the 
 
 N 
 solutions correspond, e.g., 50 cc. of — silver nitrate are taken, 
 
 and 5 cc. of ammonioferric sulphate, 5 cc. of pure nitric acid 
 and 200 cc. of water are added, and there should be required 
 exactly 50 cc. of the potassium sulphocyanate solution. The 
 same depth of reddish-brown tint should be obtained in all 
 assays by this method, as is obtained in standardizing the 
 solution. 
 
 Estimation of Soluble Haloid Salts 
 
 The estimation of these salts is based upon the powerful 
 affinity existing between the halogens and silver, and the 
 ready precipitation of the resulting chlorid, bromid and iodid. 
 Standard solution of silver nitrate is used for this purpose, 
 and for the sake of exactness and convenience, is made of 
 decinormal strength. In some cases it is advisable to use 
 centinormal solutions. 
 
 Mohr's Method with Chromate Indicator. This method 
 is the best to use, if the haloid salts are in neutral solution, 
 and salts of lead, bismuth, barium or iron are absent. If 
 the solution is acid the indicator is inadmissable, in that acids 
 have a solvent action upon silver chromate and thus prevent 
 the end-reaction from being clearly and accurately observed. 
 
ANALYSIS BY PRECIPITATION 111 
 
 If the above-mentioned metals are present, the indicator is 
 likewise useless, as these bases form insoluble, highly colored 
 compounds with the chromate. The neutral potassium chro- 
 mate (yellow chromate) which is used as the indicator must be 
 free from chlorid * and should be used in the form of a lo 
 per cent solution. 
 
 In the volumetric analysis of soluble haloid salts (chlorids, 
 bromids and iodids) 0.5 gm. of the well-dried salt is dissolved 
 in 40 cc. of water in a beaker. This is placed upon a white 
 surface and a few drops of the chromate indicator (or suffi- 
 cient to give the solution a pale yellow tint), added. The 
 
 decinormal — silver nitrate solution is then added cautiously 
 
 from a burette, stirring constantly until a permanent red 
 tint is produced. The red tint is due to the formation of 
 silver chromate, which does not appear permanent until the 
 last trace of halogen has been precipitated. 
 The reactions are as follows: 
 
 NaCl + AgNOs = AgClH- NaNOs 
 and 
 
 K.CrOi + 2AgN03 = Ag2Cr04 + 2KNO3. 
 
 If the solution to be estimated is acid it should be accu- 
 rately neutralized with ammonia, or sodium or calcium car- 
 bonate. If it is alkaline in reaction it should likewise be 
 neutralized, using acetic acid for this purpose. 
 
 In the estimation of bromids and iodids it must not be 
 forgotten to take into account the invariable presence of 
 chlorids as an impurity. 
 
 * The presence of chlorid in the chromate solution may be determined 
 by adding a small quantity of silver nitrate solution, and then some nitric 
 acid. If the red precipitate dissolves completely and leaves a clear solution, 
 chlorid is absent. If it is found to be present it may be removed by the 
 addition of a few drops of silver nitrate solution, and filtering, without using 
 any nitric acid. 
 
112 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The method in detail is exemplified in the following assays: 
 
 Estimation of Sodium Chlorid. One gm. of the well-dried 
 
 sodium chlorid is dissolved in sufficient distilled water to 
 
 measure loo cc. Of this solution lo cc. (representing o.i 
 
 gm. of the salt) is taken, a few drops of neutral potassium 
 
 N 
 chromate solution added, and then the — silver solution 
 
 ID 
 
 delivered from a burette with constant stirring or shaking 
 until the chlorid is entirely precipitated, as evidenced by the 
 formation of a permanent red color (silver chromate). The 
 equation is 
 
 NaCl + AgNOs = AgCl + NaNOg. 
 
 10)58-46 io )i6g.89 j^ 
 
 5.846 lo.gbg gms. = iooo cc. — V.S. 
 
 Thus each cc. of 10 V.S. represents 0.005846 gm. of NaCl. 
 If in the above assay 17 cc. of the silver solution were required, 
 then 17X0.005846 gm. =^0.099382 gm. or 99.382 per cent. 
 
 0.099^82 X 100 
 
 ^ =99.382 per cent. 
 
 Estimation of Ammonium Bromid. Three gms. of the salt 
 are dried at 100° C. (212'' F.) (to remove moisture, which 
 the salt readily absorbs out of the air), and dissolved in suffi- 
 cient water to measure 100 cc. 10 cc. of this solution (repre- 
 senting 0.3 gm. of the salt) are placed in a beaker, a few 
 drops of potassium chromate solution added, and then the 
 
 N 
 
 — silver nitrate V.S. carefully delivered from a burette until 
 
 a permanent red coloration is produced. Apply the equation 
 NH4Br + AgNOa = AgBr + NH4NO3. 
 
 i° )97-93 j^ 
 
 9.793 gms. = 1000 cc. — V.S. 
 
ANALYSIS BY PRECIPITATION 113 
 
 N 
 Thus each cc. of the — V.S. represents 0.009729 gm. of 
 
 N 
 NH4Br. Not more than 31.6 cc. of — AgNOs should be 
 
 required for 0.3 gm. of ammonium bromid. If the salt is 
 absolutely pure only 30.84 cc. would be required for 0.3 gm. 
 The excess is due to the presence in the commercial salt of 
 a certain amount of ammonium chlorid which is precipitated 
 by the silver nitrate as well as the bromid, and which, having 
 a lower molecular weight, requires proportionately more silver 
 nitrate to precipitate it than the bromid does. The presence 
 of chlorids must always be taken into account in the valua- 
 tion of bromids, because the latter usually contain more or 
 less of the former as an impurity. 
 
 The Determination of the Amount of Chlorid Present is 
 calculated as follows: The amount of the salt examined 
 
 . N ., . , . . , 
 
 equivalent to 1000 cc. of — silver nitrate solution is first 
 
 found thus: 
 
 31.6 : 0.3 :. 1000 cc.:x. x =9.493 gms. 
 
 This is then deducted from the quantity of pure ammonium 
 
 N 
 bromid (9.729 gms.) which is equivalent to 1000 cc. of — 
 
 silver nitrate solution. 
 
 9.729-9.493=3/. y=o.236gm. 
 
 y represents the excess of — silver nitrate solution used 
 
 up by the ammonium chlorid, reckoned in terms of ammo- 
 nium bromid, and since 5.311 gms. of NH4CI is equivalent 
 to 9.729 gms. of NH4Br, the excess which NH4CI can consume 
 
 is represented by 
 
 9.729-5.311=4-418 gms. 
 
 therefore, 
 
 4.418: 5.311 :; 0.236: 3. 2 = 0.283 gm. 
 
114 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 z represents the amount of NH4CI present in 9.493 gms. 
 of the sample. Lastly, calculate the percentage 
 
 9.493 : 0.283 : : 100 : p. p = 2.g8 per cent of NH4CI. 
 
 Thus the salt examined contained 97.02 per cent of NH4Br. 
 
 Estimation of Potassium lodid. This is conducted in 
 exactly the same manner as the preceding salts. The pres- 
 ence of chlorid (KCl) as an impurity must likewise be taken 
 into accoimt, and the calculations made to determine its 
 quantity, in the same manner as described under estimation 
 of ammonium bromid. 
 
 0.5 gm. of the well-dried salt is dissolved in 10 cc. of water, 2 
 
 N 
 drops of neutral potassium chromate are added, and then the — 
 
 AgNOa V.S. slowly added from a burette until a permanent 
 red color of silver chromate is produced. Not more than 
 30.5 cc. nor less than 30 cc. of decinormal silver nitrate V.S. 
 should be required. This quantity corresponds to 100 per 
 cent of the pure salt. 
 
 KI + AgNOs = Agl -f KNO3. 
 
 10)166.02 ^^ 
 
 N 
 
 16.602 Kms.= 1000 cc. — V.S. 
 ° 10 
 
 N 
 Each cc. of — AgNOs V.S. thus corresponds to 0.016602 
 
 gm. of KI. 
 
 Thus 0.016602X30.5 = 0.5063 gm. 
 
 To determine the amount of chlorid present as an impurity 
 calculate as follows: 
 
 The amount of the salt under examination equivalent to 
 1000 cc. is first found. 
 
 30.5 cc. :o.5 gm. : : 1000 cc. : x. :!(;= 16.393. 
 
ANALYSIS BY PRECIPITATION 115 
 
 This is deducted from 16.602 gms., the quantity of pure 
 
 N 
 KI equivalent to 1000 cc. of — AgNOa V.S. 
 
 16.602—16.393 = 0.209. 
 
 This represents the excess of standard silver solution used 
 up by the KCl, reckoned in terms of- KI. 
 
 Since 7.45 gms. of KCl is equivalent to 16.602 gms. of 
 KI, the excess which KCl can consume is represented by 
 
 16.602-7.45 = 9.152; 
 therefore 
 
 9.152 : 7.45 :: 0.209 :x. x = o.i8i. 
 
 0.181 gm. is the amount of KCl in 16.393 g™^- o^ the salt 
 examined. 
 
 The percentage is now calculated. 
 
 16.393 ■ o-i8i : : 100 : x. x= 1.104 per cent 
 
 of KCl, which leaves 98.896 per cent of pure KI. 
 
 The same method of assay may be applied to the following 
 haloid salts: LiBr, 86.86; KBr, 119,02; NaBr, 102.92; SrBr2, 
 247.47; ZnBrs, 225.21; Nal, 149-92; NH4CI, 53.50. 
 
 Titration without an Indicator — Gay-Lussac's Method. In 
 this method no indicator is used, the standard solution 
 being added imtil it ceases to produce any further precipi- 
 tation. This method is applicable to acid solution of the 
 haloid salts, and to the haloid acids— hydrochloric, hydro- 
 bromic and hydriodic; also to the estimation of silver by 
 standard solution of sodium chlorid. The method is carried 
 out in hot solutions, slightly acidulated with nitric acid, in 
 order to facilitate the precipitation of the silver halid. The 
 haloid acids are neutralized with an alkali and then slightly 
 acidulated with nitric acid before the titration is begun. The 
 calculations are precisely like those in the foregoing assays. 
 
116 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Estimation of Haloid Acids 
 
 These acids, namely, hydrochloric, hydrobromic and 
 
 hydriodic, may be estimated by Gay-Lussac's method @,bove 
 
 described, or they may be estimated by Mohr's Method. 
 
 using neutral potassium chromate as an indicator. In this 
 
 case it is necessary to carefully neutralize the acid with 
 
 N 
 ammonia and then titrate with — silver nitrate solution, 
 
 lO ' 
 
 using a few drops of chromate as indicator, in the manner 
 described in the foregoing assays. They may also be esti- 
 mated by Volhard's Method, in which an excess of the standard 
 silver nitrate solution is used, in the presence of nitric acid, 
 and the amount of the excess determined by residual titration 
 with potassium sulphocyanate, using ferric alum as the indi- 
 cator. This method is especially useful for iodids and hydri- 
 odic acid, in that the nitric acid need not be added until after 
 an excess of silver nitrate solution is used, and thus liberation 
 of iodin by the nitric acid avoided. This method is more fully 
 described further on. 
 
 The estimation of the haloid acids may also be effected 
 by neutralization with standard alkali, in the same way as 
 other acids, but since hydrobromic and hydriodic acids are 
 now frequently prepared by the method of Fothergill, in 
 which potassium bromid or potassium iodid (according to 
 the acid to be made) is brought in contact with tartaric acid 
 (as shown in the equation), an excess of the latter acid is 
 unavoidably present, and hence the neutralization method is 
 inapplicable. 
 
 KI + H2C4H4O6 = KHC4H4O6 + HI 
 
 Dtassium Tartaric acid Potassium Hydriodi 
 
 Iodid Bitartrate Acid 
 
 KBr " " HBr 
 
ANALYSIS BY PRECIPITATION 117 
 
 Assay of Hydrobromic Acid, Using Chromate as Indicator. 
 
 lo gms. of hydrobromic acid are diluted with sufi&cient dis- 
 tilled water to make loo cc. lo cc. of this solution, repre- 
 senting I gm. of the acid, is exactly neutralized with diluted 
 ammonia water (using litmus solution as indicator); 3 drops 
 of neutral potassium chromate solution are added, and then 
 
 N . . . 
 
 the — silver nitrate run m from a burette until the solution 
 10 
 
 acquires a permanent red tint. The following equation is 
 
 then applied: 
 
 HBr + AgNOs = AgBr + HNO3. 
 
 10)80.02 
 
 ■ — - — N 
 
 8.092 gms. = 1000 cc. — V.S. 
 10 
 
 If the assay is to be made by the direct percentage method, 
 8.092 cc. (8.09 cc.) of the solution (10 gms. in 100 cc.) 
 (representing 0.809 gms. of the acid) should be taken, in 
 which case each cc. of the standard silver solution consumed 
 will at once indicate i per cent. 
 
 Volhard's or Sulphocyanate Method. This method depends 
 upon completely precipitating the halogen in the presence of 
 nitric acid, by a measured excess of standard silver nitrate 
 solution, and then estimating the excess of silver by retitrating 
 with standard sulphocyanate solution, using ferric alum as 
 an indicator. 
 
 The sulphocyanate has a greater affinity for silver than 
 it has for iron, and therefore, so long as any silver is in solu- 
 tion, the sulphocyanate will combine with it and form a pre- 
 cipitate of silver sulphocyanate. 
 
 As soon as the silver is all taken up, the sulphocyanate 
 will combine with the ferric alum and strike a brownish-red 
 color. 
 
 The sulphocyanate solution is to be made of such strength 
 
118 THE ESSENTIALS OF VOLUMETRIC 7VNALYSIS 
 
 that it corresponds with the silver solution, volume for 
 volume. 
 
 The difference between the volume of silver solution origi- 
 nally added and the volume of sulphocyanate solution used, 
 will give the volume of silver solution equivalent to the haloid 
 salt present. 
 
 This method has the advantage over the direct method 
 for haloids with chromate indicator, in that it may be used 
 in the presence of nitric acid. It thus enables one to estimate 
 the haloids in the presence of phosphates or other salts which 
 precipitate silver in neutral but not in acid solutions, and also 
 in that the presence of barium, bismuth, lead, iron and other 
 metals do not interfere, as they do with the chromate in 
 Mohr's method. The presence of mercury, however, exerts 
 a disturbing influence upon the end-reaction. The nitric acid 
 acidulates the solution and thus facilitates the precipitation 
 of silver by the halogens, and prevents its precipitation by 
 other substances. The quantity of nitric acid employed is of 
 no great importance, except in the case of iodids (because 
 silver iodid is slightly soluble in nitric acid). Usually suffi- 
 cient of the acid is added to just remove the color produced 
 by the indicator. A very large excess of the acid would, 
 however, interfere with the proper determination of the end- 
 reaction, in that it to a slight extent prevents the formation 
 of ferric sulphocyanate. In the estimation of iodids by this 
 method, the nitric acid should be added after the standard 
 silver solution, while in the case of the other haloid salts the 
 acid may be added before. 
 
 The indicator also should be added after the standard 
 silver solution, when estimating iodids, because being a ferric 
 salt it is, like nitric acid, capable of liberating iodin. 
 
 The solutions required for this method are: 
 (I) Decinormal Silver Nitrate (page io6); 
 
ANALYSIS BY PRECIPITATION 119 
 
 (II) Decinormal Potassium Sulphocyanate (page 109); 
 
 (III) Ferric Alum Solution. (The indicator.) 
 
 This is a 10 per cent aqueous solution of ferric-ammonium 
 sulphate, Fe2(S04)3- (NH4)2S04 + 24H20. 
 
 (IV) Nitric Acid (C. P.). This must be free from nitrous 
 acid. If it or any of the lower oxids of nitrogen are present 
 they may be removed by diluting with one-fourth part of 
 water and boiling until colorless. 
 
 The process is exemplified in the following assays : 
 Assay of Hydriodic Acid by the Sulphocyanate Method. 
 Introduce into a 200-cc. stoppered flask 2 gms. of the acid, 
 add 50 cc. of distilled water and 25 cc. (accurately measured) 
 of decinormal silver nitrate, shake thoroughly, and then add 
 5 cc. of the ferric alum solution and 3 cc. of nitric acid, 
 C. P. The flask is stoppered and again thoroughly shaken, 
 and finally, the decinormal potassium sulphocyanate run in 
 slowly from a burette, until a permanent reddish-brown tint 
 is produced. Note the number of cc. of sulphocyanate solu- 
 tion employed. 
 
 Deduct this from the 25 cc. of silver solution added, and 
 multiply the remainder by the factor for HI, which is 0.01269. 
 
 I. HI + AgNOa = Agl + HNO3. 
 
 10) 127.92 10)169.89 
 
 12.792 gms. 16.989 gms. = 1000 cc. — V.S. 
 
 N ^° 
 
 0.012792 gm. of HI = i cc. — V.S. 
 10 
 
 II. AgNOa -hKSCN=AgSCN-hKN03. 
 
 III. Fe2(NH4)2(S04)4 + 6KSCN 
 
 = Fe2(SCN)o + (NH4)2S04+3K2S04. 
 
 The reddish-brown color which marks the end-reaction 
 is due to the formation of Fe2(SCN)6 ferric sulphocyanate. 
 
120 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Assuming that in the above titration 10.2 cc. of decinormai 
 sulphocyanate were employed, then 25 cc. — 10.2 = 14.8 cc. 
 
 0.012792X14.8 = 0.1893 gm. 
 
 0.1893X100 
 
 = 940 per cent. 
 
 Assay of Syrup of Hydriodic Acid. Six gms. of the syrup 
 are weighed off carefully in a 200-cc. stoppered flask, 20 cc. 
 of distilled water are added, followed by 10 cc. of decinormai 
 silver nitrate and the mixture thoroughly shaken. Five cc. of 
 diluted nitric acid and 3 cc. of the ferric alum solution are 
 now added, and after again shaking the mixture it is titrated 
 with decinormai potassium sulphocyanate until a permanent 
 reddish-brown tint appears. If 5.1 cc. of the sulphocyanate 
 solution are used, this quantity is deducted from the 10 cc. 
 
 N . . 
 
 of — silver nitrate solution added, which leaves 4.9 cc, the 
 10 
 
 quantity of the latter which reacted with the syrup. 
 
 Assay of Syrup of Ferrous lodid by the Sulphocyanate 
 Method. Take 10 gms. of the syrup, dilute it with distilled 
 water to measure 100 cc. Of this solution 15.58 cc. are 
 mixed with 15 cc. of water, 6 cc. of decinormai silver nitrate 
 and 2 cc. each of diluted nitric acid and ferric alum solution, 
 and then after thoroughly shaking the mixture is titrated 
 with decinormai sulphocyanate until a permanent reddish- 
 brown tint appears. Not more than i cc. of the latter should 
 be used. This i cc. deducted from the 6 cc. of decinormai 
 silver nitrate, leaves 5 cc, the quantity of the latter which 
 reacted with the ferrous iodid. Each cc. represents i per 
 cent. The equation is 
 
 Fel2 + 2AgN03 = 2AgI + Fe(N03)2. 
 2)311.66 
 
 io)2S^83_ j^ 
 
 15.583 gms. = 1000 cc. — AgNOg V.S. 
 
ANALYSIS BY PRECIPITATION 121 
 
 N 
 I cc. of — AgNOs V.S. thus represents 0.015583 gm. 
 
 of ferrous iodid. 
 
 In this (direct percentage) method, a quantity of the syrup 
 is taken which equals the weight of pure Fel2 represented 
 by 100 cc. of the decinormal silver nitrate solution. 
 
 Strontium Iodid (Srl2) and Zinc Iodid (Znl^) may be 
 assayed by the sulphocyanate method above described. 
 
 The sulphocyanate method may be used for the estima- 
 tion of chlorids, and bromids, as well as iodids. 
 
 When used for the estimation of chlorids, however, the 
 precipitated silver chlorid must be removed by filtration, 
 because of the action of ferric sulphocyanate upon silver 
 chlorid, which causes the results of the analysis to be too 
 high. 
 
 In the case of silver bromid no such reaction takes place, 
 or if it does, the reaction is so slow as not to interfere in the 
 least with the getting of accurate results. Therefore, when 
 this method is used for the determination of bromids or iodids, 
 there is no need for filtering to remove the precipitate. 
 
 Estimation of Cyanogen 
 
 Titration with Standard Silver Solution to First Appearance 
 of a Precipitate— Liebig's Method. This gives fairly accurate 
 results. The cyanogen must be in the form of an alkali salt 
 and in an alkaline solution. If hydrocyanic acid is to be 
 estimated, it must be made alkaline by the addition of potas- 
 sium or sodium hydroxid. The standard silver solution is 
 then added cautiously and with constant stirring imtil a per- 
 manent precipitate of silver cyanid is produced. When silver 
 nitrate is added to an alkaline solution of a cyanid, the 
 precipitate which at first forms redissolves on stirring and a 
 soluble double cyanid (AgCN,KCN or AgCN,NaCN, depending 
 
122 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 upon the alkali used) is formed, and when all of the cyanid 
 has been taken up, the further addition of silver nitrate causes 
 a decomposition of this soluble double salt and the formation 
 of a permanent precipitate of silver cyanid. Therefore, the 
 first appearance of this precipitate affords a delicate proof of 
 the completion of the reaction. 
 
 These equations illustrate the reactions: 
 
 2NaCN+AgN03 =AgCN,NaCN+NaN03. 
 
 Double cyanid 
 of silver and sodium 
 
 AgCN,NaCN + AgNOg = 2 AgCN + NaNOs- 
 
 Silver cyanid 
 
 According to these equations it is seen that the end- 
 reaction is reached when two molecules of the alkali cyanid 
 have reacted with one molecule of silver nitrate. The slightest 
 excess of silver nitrate above this quantity brings about a 
 decomposition of the double salt and a precipitation of the 
 silver cyanid, as above stated^ 
 
 This double combination is so firm that if the estimation 
 is done in the presence of a halogen, no permanent precipitate 
 of silver halid is formed until after all of the cyanogen present 
 has been converted into a double salt. This fact is taken 
 advantage of in the processes for hydrocyanic acid and alkali 
 cyanid in which potassium iodid is employed as indicator 
 in the presence of ammonia water. The latter prevents the 
 precipitation of silver cyanids and thus allows the silver iodid 
 to precipitate alone. 
 
 N 
 I cc. of — AgNOs V.S. =0.005202 gm. CN; 
 
 0.005302 gm. HCN; 
 0.0098 gm. NaCN; 
 0.0130 gm. KCN. 
 
ANALYSIS BY PRECIPITATION 123 
 
 Assay of Hydrocyanic Acid (HCN = 26.01). Dilute hydro- 
 cyanic acid may be estimated by weighing out about 5 gms. 
 and adding it without delay (to avoid evaporation) to sufficient 
 sodium or potassium hydroxid solution to convert the acid 
 into sodium or potassium cyanid (NaCN or KCN) and leave 
 the solution strongly alkaline. The mixture is then largely 
 diluted with water (50 to 100 cc); this is to enable one more 
 clearly to observe the end-point. 
 
 The decinormal silver nitrate solution is then delivered 
 in until a permanent turbidity occurs. 
 
 The difficulty experienced in this process is in the con- 
 version of the acid into the cyanid. Sodium cyanid has a 
 strong alkaline reaction, turning litmus blue, when only a 
 small proportion of the acid has been neutralized. If the 
 titration is conducted before the acid is completely neutralized 
 that which is free will not be acted upon. Indeed, cyanid of 
 sodium may be estimated in the presence of hydrocyanic acid 
 in this way. 
 
 According to Senier, the following procedure will answer 
 well: 
 
 To the dilute hydrocyanic acid add sodium hydroxid to 
 strong alkaline reaction, determined by litmus tincture.* 
 
 N 
 Then titrate with — silver nitrate, drop by drop If the 
 
 liquid becomes acid, add a little more soda solution to bring 
 it back to alkalinity, and continue the titration until the tur- 
 bidity indicates the end of the reaction. The liquid must be 
 kept alkaline throughout the process. It is not well to add 
 too much alkali at the beginning as this will use up too much 
 
 * Poirrer Blue C^B is better, in that it is not affected by alkali cyanids, 
 but gives a very sharp indication in the presence of the slightest excess of 
 alkali hydroxid. The amount of alkali used should be as near as possible 
 that which is required to just convert the acid into the alkali cyanid; too much 
 or too little alike affect the accuracy of the result. 
 
V2i THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 of the silver solution and make the reading a trifle too high. 
 The following equations, etc., explain the reactions: 
 
 2HCN + 2NaOH = aNaCN + 2H2O. 
 10)54.02 10)98.02 
 
 5.402 gms. 9.802 gms. 
 
 2NaCN + AgNOs = AgCN.NaCN + NaNOg. 
 10)98.02 10)169.89 „ 
 
 9.802 gms. 16.989 gms. or looo cc. — V.S. 
 
 It is seen that 5.402 gms. of real HCN are equivalent to 
 9.802 gms. of sodium cyanid, and represent 16.989 gms. of 
 
 N 
 silver nitrate or 1000 cc. of the — V.S. That is, 1000 cc. 
 
 10 
 
 N 
 of the — AgNOs V.S. may be added to a solution containing 
 
 9.802 gms. of sodium cyanid and no precipitate will be pro- 
 duced, but if one or two drops more of the standard solution 
 be added, a precipitate is at once formed, the double salt 
 being broken up and silver cyanid produced." 
 
 AgCN,NaCN + AgNOs = 2 A^CN + NaNOs. 
 
 N 
 Each cc. of the — silver solution which fails to produce 
 10 . ^ 
 
 a precipitate represents 0.009802 gm. of NaCN, which is 
 equivalent to 0.005402 gm. of HCN. 
 
 Titration with Standard Silver Solution, Using Chromate 
 Indicator — ^Vielhaber's Method. This method is especially 
 recommended for the assay of weak solutions containing 
 hydrocyanic acid, as bitter almond oil, bitter almond water, 
 cherry laurel water, etc., but it may also be employed for 
 alkaline cyanids. 
 
 A sufficient quantity of an aqueous suspension of mag- 
 nesium hydroxid * to make the solution opaque and distinctly 
 
 * Calcined magnesia triturated with water. 
 
ANALYSIS BY PRECIPITATION 125 
 
 alkaline is added; this is followed by a few drops of potassium 
 
 N 
 chromate indicator and then the — silver nitrate delivered 
 
 lO 
 
 into the mixture from a burette until a permanent red tint 
 appears, as in the titration of haloid salts. The method is 
 a very satisfactory one if chlorids are absent. 
 
 The reactions in this method are the same as in the fore- 
 going, but the end-reaction (the production of silver chromate) 
 does not occur until the double cyanid is completely decom- 
 posed, at which point the addition of another drop of silver 
 solution reacts with the chromate and produces the red pre- 
 cipitate (silver chromate). 
 
 The equations are as follows: Sodium is used in the 
 equations instead of magnesium in order to make the explana- 
 tion clearer. 
 
 (a) 2NaCN-F AgNOg = AgCN,NaCN -HNaNOs = (2HCN) ; 
 
 (b) AgCN,NaCN + AgNOs = 2 AgCN + NaNO.,. 
 
 These equations show that it requires two molecules of 
 
 silver nitrate to completely precipitate two molecules of cyanid. 
 
 169.89 gms. of AgNOg is equivalent to 27.01 gms. of HCN, 
 
 while by Liebig's method 169.89 gms. of AgNOs is equivalent 
 
 to 54.02 gms. of HCN. 
 
 N 
 I cc. — AgNOs V.S. =0.002601 gm. CN; 
 
 0.002701 gm. HCN; 
 
 0.004901 gm. NaCN; 
 
 0.0065 1 1 gm. KCN. 
 
 Example. 1.35 gms. of the diluted acid is mixed with 
 enout^h water and magnesia to make an opaque mixture of 
 about 10 cc. Add to this 2 or 3 drops of potassium chromate 
 solution and then from a burette deliver the decinormal silver 
 
126 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 nitrate V.S. until a red tint is produced which does not 
 disappear by shaking. 
 
 Titration with Standard Silver Solution, Using Potassium 
 lodid as Indicator. This method is recommended by W. J. 
 Sharwood, J. A. C. S., 1897, 400-434, and is a modification 
 of the method proposed by M. Georges Deniges, Ann. chim. 
 phys. (7) 6.381. 
 
 In this method 5 gms. of hydrocyanic acid are diluted 
 with distilled water to measure 50 cc. Then 25 cc. of this 
 solution, after the addition of 5 cc. of ammonia water and 
 3 drops of a 20 per cent potassium iodid solution, are titrated 
 with tenth-normal silver nitrate, until a slight permanent 
 precipitate occurs. The potassium iodid in this process acts 
 as indicator. 
 
 The reactions may be expressed thus: 
 
 (i) HCN + NH4OH-NH4CN + H2O; 
 
 (2) 2NH4CN+AgN03 = NH4Ag(CN)2-^NH4N03; 
 
 (3) NH4Ag(CN)2 + AgNOs = NH4NO3 + 2AgCN; 
 
 (4) KI + AgNOa = KNO3 + Agl. 
 
 The silver nitrate forms with the cyanid a double sale 
 which is soluble, no precipitate occurring until after all of 
 the cyanid has entered into combination as the double salt; 
 then the further addition of silver nitrate decomposes the 
 double salt, and a precipitate of silver cyanid occurs. In 
 the presence of ammonia water, however, as in the above 
 assay, the precipitation of silver cyanid is prevented, but the 
 iodid is now (not before) acted upon by the silver solution 
 and a precipitate of silver iodid occurs, which very delicately 
 indicates the end-reaction. 
 
 Each cc. of the standard silver nitrate solution used repre- 
 sents 0.005302 gm. of absolute HCN. If 26.51 cc. of the above 
 
ANALYSIS BY PRECIPITATION 127 
 
 solution are taken instead of 25 cc, each cc. will represent at 
 once 0.2 per cent. 
 
 Potassium cyanid is assayed in the same way. 
 
 Estimation of Potassium Cyanid (KCN = 65.ii). i gm. 
 of potassium cyanid is dissolved in sufficient distilled water 
 to make 100 cc, then 65.11 cc. of this solution mixed with 
 5 cc. of ammonia water and 3 drops of potassium iodid 
 
 N 
 solution are titrated with — AgNOs V.S. until the appearance 
 
 of a permanent precipitate. Each cc. indicates 2 per cent. 
 
 N 
 I cc. of — AgNOs V.S. =0.0130 gm. KCN. 
 
 Estimation of Silver Salts 
 
 Soluble silver salts are estimated by direct titration with 
 standard sodium chlorid, the process being exactly the converse 
 of the precipitation methods for halogens. The standard 
 sodium chlorid solution is added to the solution of the silver 
 salt until precipitation ceases; or the titration may be done 
 in the presence of chromate indicator, the end-point being 
 then known, to be reached when the red color of the silver 
 chromate disappears. The first of these methods is imprac- 
 ticable. Too much time being consumed in waiting for the 
 precipitate to settle so as to render the supernatant liquid 
 sufficiently clear to recognize whether a precipitate is produced 
 in it by the fiu-ther addition of the standard solution. 
 
 If chromate indicator is used, the end-point is easily over- 
 stepped, because of the slow decomposition of the silver 
 chromate by the chlorid. It is best to add an excess of sodium 
 chlorid solution and then retitrate with standard silver nitrate 
 solution until the red color appears. 
 
 Silver salts may also be titrated by means of standard 
 sulphocyanate solution, using ferric alum as indicator. 
 
128 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 Assay of a Solution of Silver Nitrate by Means of — 
 
 Sodium Chlorid. Ten gms. of the solution are introduced 
 into a beaker and diluted with lo cc. of distilled water. Two 
 drops of yellow potassium chromate solution are added as 
 
 N 
 indicator, and then a measured excess of — sodium chlorid 
 
 10 
 
 (sufficient of this must be added to completely destroy the 
 red color) added slowly from a burette and with constant 
 stirring. Assuming thajt 20 cc. were used, then the excess 
 
 N 
 may be ascertained by titrating back with — silver nitrate 
 
 V. S. until a permanent red tint is produced. Whatever number 
 
 N 
 of cc. of — silver nitrate are used, that number represents 
 
 N 
 the quantity of ^ sodium chlorid which was added in excess, 
 
 and must be deducted from the 20 cc. of the sodium chlorid 
 solution employed. Assuming this number to be 3 cc, then 
 
 N 
 3 from 20 leaves 17 cc, which is the exact quantity of — 
 
 NaCl V.S. which reacted with the silver in the solution examined. 
 
 N 
 The — factor for silver nitrate multiplied by 17 will then 
 
 give the exact weight of silver nitrate in the 10 gms. of solu- 
 tion taken. The following equation illustrates the reaction 
 which occurs: 
 
 AgNOs + NaCl = AgCl + NaNOg. 
 
 10)169.89 io )s8-46 j^ 
 
 16.989 gms. 5.846 gms. = 1000 cc. — V.S. 
 
 0.016989 gm. of silver nitrate is thus represented by each 
 
 N 
 of — sodium chlorid. 
 10 
 
 In the assay of silver nitrate crystals, 0.2 gm. is taken, 
 
ANALYSIS BY PRECIPITATION 129 
 
 dissolved in lo cc. of distilled water, and then treated as in 
 the foregoing assay. In the case of molded silver nitrate about 
 the same quantity is taken for assay. Of mitigated silver 
 nitrate, i gm. may be taken. 
 
 Silver oxid (Ag20) may be converted into nitrate by 
 solution in nitric acid and then tested as above. Free nitric 
 acid is apt to be present in this case and therefore the solu- 
 tion should be neutralized, before it is assayed, if the above- 
 described method is to be employed. The presence of free 
 acid does not interfere, however, if the method of Gay-Lussac 
 or the sulphocyanate method be employed. 
 
 N 
 Assay of Silver Nitrate by Means of — Sulphocyanate. 
 
 This method, as applied to the assay of halogen compounds, 
 is described in the preceding pages. The great advantage 
 which this method presents over the others, is that the presence 
 of most other metals does not interfere. The only metal 
 which does materially interfere with the determination of 
 silver is mercury. 
 
 Example. A weighed quantity (0.2 to 0.5 gm.) of the 
 silver salt is dissolved in water, some diluted nitric acid and 
 ammonium ferric sulphate solution are added, and the mixture 
 
 N 
 then titrated with — potassium sulphocyanate until a per- 
 manent reddish-brown color of ferric sulphocyanate is produced. 
 The following equation explains the reactions: 
 
 AgNOs + KSCN = AgSCN -t- KNO3. 
 10)169. 89 10 )97.18 
 
 16.989 gms. 9-7i8 gms. or 1000 cc. standard V.S. 
 
 Thus each cc. of the standard V.S. represents 0.016989 
 gm. of pure silver nitrate, or 0.010788 gm. of metallic silver. 
 
 Estimation of Metallic Silver and Silver Alloys. A quantity 
 of the metal; weighing about 0.5 gm., is dissolved in 10 cc. . 
 
130 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 of nitric acid, and after complete solution is attained it is 
 
 heated sufficiently to drive off all traces of nitrous acid. The 
 
 solution is then diluted with about loo cc. of distilled water 
 
 and assayed by one of the methods described under the assay 
 
 of silver nitrate. The sulphocyanate method is the preferred 
 
 one. 
 
 Estimation of Alkali lodids by Precipitation with Mercuric 
 
 Chlorid Solution (Personne). Alkali iodids may also be 
 
 N 
 estimated by titration with — mercuric chlorid V.S., the 
 •' lO 
 
 termination of the operation being indicated by the formation 
 of a red precipitate. 
 
 (i) 4KI + HgCla = 2KCI + Hgl2 . 2KI (soluble) 
 
 (2) Hgl2.2KI+HgCl2 = 2KCl + 2Hgl2 
 
 This process originated with M. Personne, and is foiuided 
 on the fact that if a solution of mercuric chlorid be added 
 to one of potassium iodid, in the proportion of one equivalent 
 of mercuric chlorid to four of potassium iodid, red mercuric 
 iodid is formed, which dissolves at once to a colorless solution. 
 The slightest excess of mercuric chlorid will cause a brilliant 
 red precipitate (Hgl2) to make its appearance. 
 
 4KI + HgCla = 2KCI + Hgl2.2KI (soluble). 
 
 20)664.08 20 )270.92 
 
 33.204 gms. 13.546 gms. or 1000 cc. of standard solution. 
 
 Thus each cc. of standard solution of the above strength 
 represents 0.033204 gm. of potassium iodid, which means that 
 I cc. is the largest quantity of this standard solution which 
 can be added to 0.033204 gm. of potassium iodid without 
 producing a permanent precipitate. 
 
 N 
 The above solution of mercuric chlorid is strictly a — V.S. 
 
 10 
 
 The author of this process states that neither chlorids, 
 bromids, nor carbonates interfere with the reaction. 
 
ANALYSIS BY PRECIPITATION 
 
 131 
 
 TABLE OF SUBSTANCES ESTIMATED BY PRECIPITATION 
 
 Name. 
 
 Formula. 
 
 Molecular 
 Weight. 
 
 Standard 
 
 Solution 
 
 Used. 
 
 Factor.* 
 
 Acid, hydrobromic . 
 " hydrochloric . . 
 " hydrocyanic . . 
 
 tt It 
 
 " hydriodic 
 
 Allyl-iso-thiocyanate 
 
 Ammonium bromid 
 
 chlorid . 
 
 " iodid . . 
 
 Calcium bromid . . . 
 
 " chlorid 
 
 Ferrous bromid . . . . 
 
 " iodid 
 
 Lithium bromid . . . 
 
 Potassium bromid . , 
 
 " chlorid . . 
 
 " cyanid . . 
 
 HBr 
 HCl 
 HCN 
 
 HCN 
 
 HCN 
 
 HI 
 
 CSNC3H5 
 NH^Br 
 NH,C1 
 NHJ 
 CaBrj 
 CaClj 
 FeBr^ 
 
 Felj 
 
 LiBr 
 
 KBr 
 
 KCl 
 
 KCN 
 KCN 
 KCN 
 
 80.92 
 36.46 
 
 27. CI 
 27.01 
 27.01 
 
 127.92 
 
 97.08 
 97.96 
 
 53-5° 
 144.96 
 
 199-93 
 III. 01 
 215 .66 
 309.66 
 
 86.86 
 119.06 
 
 74-56 
 
 65.11 
 65.11 
 65.11 
 
 N 
 -AgN03 
 
 N 
 
 — AgNO;i 
 
 N 
 
 - AsNO, 
 10 ^ ' 
 
 without 
 indicator 
 
 N 
 
 - AgN03 
 
 with chromate 
 indicator 
 
 N 
 -AgN03 
 
 with iodid 
 indicator 
 
 N 
 
 - AgNOa 
 
 N 
 -AgN03 
 
 without 
 indicator 
 
 N 
 
 - AgN03 
 
 with chromate 
 indicator 
 
 N 
 -AgN03 
 
 with iodid 
 indicator 
 
 0.008092 
 0.003646 
 0.005402 
 
 0.002701 
 
 0.005402 
 
 0.012792 
 
 0.004854 
 0.009729 
 0.00535 
 0.014496 
 o . 009996 
 
 o-ooSSS 
 
 0.010783 
 
 0.015483 
 
 0.008686 
 
 0.011906 
 
 0.007456 
 
 o 013022 
 0.0065 1 1 
 0.013022 
 
 * This is the coefficient by which the number of cc. used of the decinormal solu- 
 tion is to be multiplied in order to obtain the quantity of pure substance in the 
 sample analyzed. It represents the weight of the substance precipitated by i cc. 
 of the decinormal solution. 
 
132 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 TABLE OF SUBSTANCES ESTIMATED BY PRECIPITATION— 
 
 Continued 
 
 Name. 
 
 Formula. 
 
 Molecular 
 Weight. 
 
 Standard 
 
 Solution 
 
 Used. 
 
 Factor.* 
 
 Potassium iodic! 
 
 KI 
 
 KSCN 
 
 Ag2 
 
 AgN03 
 Ag^O 
 
 NaBr 
 
 NaCl 
 Nal 
 SrBrj 
 Sri, 
 ZnBrj 
 ZnClj 
 Znlj 
 
 166.02 
 97.18 
 2X107.88 
 
 169.89 
 231.76 
 
 102.92 
 
 58.46 
 149.92 
 247-47 
 341-47 
 225.21 
 136.29 
 319.21 
 
 N 
 
 - AgNO, 
 
 It 
 
 N 
 
 - NaC! or 
 10 
 
 N 
 
 — KSCN 
 
 10 
 
 ([ 
 
 N 
 
 - AgNO, 
 
 0.016602 
 
 '* sulphocyanate . . 
 
 0.009653 
 0.010788 
 
 0.016989 
 0.011588 
 
 0.010292 
 
 0.005846 
 0.014992 
 0.012373 
 
 0-017573 
 0. 01 I 260 
 
 
 " oxid 
 
 
 chlorid 
 
 
 Strontium bromid 
 
 ** iodid 
 
 Zinc bromid 
 
 " chlorid 
 
 0.006814 
 
 
 0.015960 
 
 
CHAPTER X 
 
 ANALYSIS BY OXIDATION AND REDUCTION 
 
 An extensive series of analyses are niade by these methods 
 with extremely accurate results; in fact, the results are generally 
 more accurate than those obtained by gravimetric methods. 
 
 The principle involved is exceedingly simple. An oxidizing 
 agent is employed for the estimation of an oxidizable sub- 
 stance, and likewise a reducing agent is employed for the 
 estimation of a reducible substance. Oxidizing agents are 
 always reducible and reducing agents always oxidizable. An 
 oxidation and a reduction take place at the same time, i.e., 
 the oxidizing agent is itself reduced in the operation and the 
 reducing agent is at the same time oxidized. 
 
 Thus substances which are capable of absorbing oxygen 
 or are susceptible of an equivalent action may be accurately 
 estimated by subjecting them to the action of an oxidizing 
 agent of known power, and from the quantity of the latter 
 required for complete oxidation the weight of the oxidizable 
 substance is ascertained. 
 
 Example. Ferrous oxid (FeO), an oxidizable substance, 
 is ever ready to take up oxygen, while potassium permanganate 
 and potassium dichromate are always ready to give up some 
 of their oxygen. When potassium permanganate gives up 
 its oxygen in this way it is reduced and decolorized, while 
 the ferrous oxid in taking up oxygen is oxidized to ferric oxid 
 (Fe203). The decolorization of the permanganate here spoken 
 of is taken advantage of in volumetric analysis for the deter- 
 mination of the completion of the oxidation. The perman- 
 ganate in the form of a standard solution being slowly delivered 
 
 133 
 
134 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 from a burette, until it is no longer decolorized, the iron salt 
 is known to be completely oxidized, when the permanganate 
 is no longer reduced. The reaction is as follows: 
 
 ioFeO + 2KMn04=sFe203 + 2MnO +K2O. 
 
 Ferrous oxid Ferric oxid 
 
 The oxidation of ferrous oxid by potassium dichromate 
 is sho^vn by the following equation : 
 
 6FeO + K2Cr207= sFeaOs + CrgOs + K2O. 
 
 As before stated, an oxidation is always accompanied by 
 a reduction, the oxidizing agent being itself reduced in the 
 operation. ^ As shown in the above equations, the manganic 
 compound is reduced to a manganous, and the chromic to 
 a chromous, while the ferrous salt is oxidized to a ferric 
 condition. 
 
 In the same way any substance which readily yields oxygen 
 in definite quantity or is susceptible of an equivalent action 
 which involves its reduction to a lower quantivalence, may 
 be estimated by ascertaining how much of a reducing agent 
 of known power is required for its complete reduction. 
 
 Example. The available chlorin in bleaching powder may 
 be accurately ascertained by treating it with a standard solu- 
 tion of arsenous oxid, and from the volume of the solution 
 required for the complete reduction of the chlorin, the quantity 
 of the latter present is found, or in other words, from the 
 quantity of arsenous oxid (AS2O3), oxidized to arsenic oxid 
 (AS2O5) the weight of the chlorin present is ascertained. 
 
 The principle substances which are used as oxidizing 
 agents in volumetric analysis, are potassium permanganate, 
 potassium dichromate and iodin. The latter contains no 
 oxygen, but it abstracts hydrogpn from accompanying water 
 and liberates the oxygen which does the oxidizing, hence iodin 
 is known as an indirect oxidizing agent. The other two 
 
ANALYSIS BY OXIDATION AND REDUCTION 135 
 
 contain available oxygen which they readily give up when 
 brought in contact with an oxidizable substance. 
 
 The principal reducing agents or deoxidizers which are 
 used in volumetric analysis are, sodium thiosulphate, sul- 
 phurous acid, oxalic acid, arsenous oxid, stannous chlorid, 
 ferrous oxid, hydriodic acid, hydrosulphuric acid, metallic 
 zinc, and magnesium. 
 
 /N\ 
 Preparation of Decinormal I — ) Potassium Permanganate 
 
 N 
 (2KMn04 = 3i6.o6; — V.S. = 3.1606 gms. in i liter). Abso- 
 lutely pure potassium permanganate cannot be obtained, 
 therefore the preparation of a decinormal solution of this 
 salt cannot be effected by simply dissolving the requisite 
 proportion of the molecular weight in the water. The presence 
 of oxidizable matter in the water used, the contact of dust 
 and exposure to light, have a tendency to decompose the 
 salt and hence weaken the standard solution. It is therefore 
 advisable to use boiling distilled water, and to preserve the solu- 
 tion in amber glass bottles, provided with ground-glass stoppers. 
 It will then retain its strength for several weeks, but should 
 nevertheless be checked by titration immediately before using. 
 It is not necessary, and it is usually undesirable, to make 
 the solution an exact decinormal one. It is preferable to 
 fix the titer of the solution and employ it as it is. 
 
 Place 3.5 gms. of pure crystallized potassium perman- 
 ganate in a flask, add 1000 cc. of distilled water, and boil 
 imtil the crystals are dissolved; put a plug of absorbent 
 cotton in the mouth of the flask and set it aside for two days 
 so that any suspended matter may deposit. After the lapse 
 of this time pour off the clear solution into a glass-stoppered 
 bottle, and when wanted for use standardize by either of the 
 following methods: 
 
136 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Standardisation by Means of Iron. Thin annealed 
 binding- wire, free from rust, is one of the purest forms of iron.* 
 
 O.I gm. of such iron is placed in a flask which is provided 
 with a cork through which a piece of glass tubing passes, 
 to the top of which a piece of rubber tubing is attached, 
 which has a vertical slit about one inch long in its side, and 
 which is closed at its upper end by a piece of glass rod (this 
 arrangement is known as the " Bunsen Valve "). (See Fig. 
 41.) Diluted sulphuric acid is added and gentle heat applied. 
 The iron dissolves and the steam and liberated hydrogen escape 
 through the slit under slight pressure. The air is thus pre- 
 vented from entering and the ferrous solution protected from 
 oxidation. 
 
 A better form of apparatus in which to dissolve the iron 
 
 and avoid oxidation through admission 
 
 Jl of air is shown in Fig. 42. A loo-cc. 
 
 11 flask is fitted with a rubber stopper and 
 
 l| a I I shaped glass tube; into this 
 
 11 flask is placed 20 cc. of diluted sulphuric 
 
 •" 1 acid (i : 5) and then 2 or 3 crystals of 
 
 Fig. 41. Fig. 42. 
 
 pure sodium carbonate; this causes an evolution of carbon 
 dioxid which expels the air from flask. The o.i gm. of iron 
 
 * This contains 99.6 per cent of iron. 
 
ANALYSIS BY OXIDATION AND REDUCTION 137 
 
 wire above described is now introduced, the stopper inserted, 
 and a beaker containing a solution of pure sodium carbonate 
 placed in position so that the tube will dip into the solution. 
 Gentle heat is applied \intil the iron is wholly dissolved, and 
 only a few minute particles of carbon remain (which must 
 not be mistaken for iron). When the fiame is withdrawn 
 the cooling of the flask and contents causes a drawing up 
 of the sodium carbonate solution, but the first drops that 
 enter the flask cause an effervescence with evolution of carbon 
 dioxid, which drives the liquid back and at the same time 
 fills the flask with the gas; this is repeated until the flask 
 and contents are cold. Another useful form of apparatus 
 for this purpose is shown in 
 
 Fig- 43- 
 
 When the iron is completely 
 
 dissolved a small quantity of 
 cold, recently boiled, distilled 
 water should be used to rinse 
 the lower end of the stopper 
 and the neck of the flask, and 
 the titration with potassium per- 
 manganate at once begun and 
 continued imtil a faint permanent 
 pink color is produced. If the 
 solution is decinormal, exactly 
 17.84 cc. will be required to pro- 
 duce this result. 
 
 The iron is converted by 
 the sulphiuic acid into ferrous 
 sulphate, Fe2 + 2H2SO4 = 2FeS04 + 2H2. This ferrous sulphate 
 is easily oxidised by the air, and therefore it is directed that 
 access of air should be prevented, and the distilled water with 
 
 Fig. 43- 
 
138 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 which the solution is diluted previously boiled in order to drive 
 off any dissolved free oxygen. 
 
 loFeSOi + 2KMn04 + 8H2SO4 
 
 10 0)558.2 100 )316.06 ^ 
 
 5.582 gms. 3.1606 gms. or looo cc. — V. S. 
 
 = 5Fe2 (SO4) 3 + K2SO4 + 2MnS04 + 8H3O. 
 
 N 
 
 This equation, etc., shows that each cc. of — perman- 
 ganate represents 0.005582 gm. of metallic iron. 
 
 Standardization by Means of Oxalic Acid. 0.12605 S'^- 
 of the pure crystallized acid is weighed (or 20 cc. of decinormal 
 oxalic acid carefully measured) and placed in a flask, with 
 3 cc. of sulphuric acid C. P. and distilled water to make 100 
 cc. The solution is warmed to 60° C. (140° F.) and the 
 permanganate solution delivered in from a burette. 
 
 The action is in this case less decisive and rapid than 
 in the titration with iron, and more care should be used. 
 The color disappears slowly at first, but afterwards more 
 rapidly. 
 
 Note the number of cc. of the permanganate solution 
 used, and then dilute the remainder so that equal volumes 
 of decinormal oxalic acid and decinormal permanganate solu- 
 tion will exactly correspond. 
 
 Example. Assuming that 18.5 cc. of the permanganate 
 solution are required to produce a permanent pink tint in 
 the above test, then the permanganate solution must be diluted 
 with distilled water in the proportion of 18.5 cc. of the per- 
 manganate solution and 1.5 cc. of v/ater, or 1850 cc. to 150 cc. 
 
 After dilution a new trial should be made, in which 50 
 
 cc. of the diluted permanganate solution should require exactly 
 
 N 
 i;o cc. of — oxalic acid V.S. 
 •^ 10 
 
ANALYSIS BY OXIDATION AND REDUCTION 139 
 
 The reaction between potassium permanganate and oxalic 
 acid is illustrated by the following equation: 
 
 2KMn04 + 5(H2C204.2H20) +3H2SO4 
 
 = K2SO4 + 2MnS04 + 10CO2 + 18H2O. 
 
 Standardization by the lodometric Method, This 
 method, which was proposed by Volhard, is the most accurate 
 and rapid for the standardization permanganate. It is based 
 upon the fact that potassium permanganate reacts with potas- 
 sium iodid in solutions acidulated with either hydrochloric or 
 sulphuric acid, and liberates an equivalent quantity of iodin, 
 which may be estimated by standard solution of sodium 
 thiosulphate. The reactions are illustrated by the equations 
 
 (ff) 2KMn04+8H2S04 + ioKI 
 316.06 
 
 = 2MnS04 + 6K2S04+8H20 + 5l2; 
 
 1269.2 
 
 {b) I2 + 2 (Na2S203 . 5H2O) = 2NaI + Na2S406 + 10H2O. 
 
 253-84 49I5-44 
 
 Thus it is seen that 2KMn04 (316.06 gms.) containing 
 
 five atoms of available oxygen, has the power of liberating 
 
 its equivalent of iodin, i.e., 10 atoms or 1269.2 gms. (see 
 
 equation a) and that 496.44 gms. of sodium thiosulphate 
 
 will reduce 253.84 gms. of iodin (see equation b). Hence 
 
 N 
 1000 cc. of — sodium thiosulphate (contammg 25.384 gms.) 
 10 
 
 will reduce, and therefore be equivalent to 12.692 gms. of 
 iodin, which in turn represents 3.1606 gms. of potassium 
 
 N 
 permanganate. Therefore i cc. of the — thiosulphate repre- 
 sents 0.012692 gm. of iodin and 0.0031606 gm. of potassium 
 permanganate, which latter is the quantity of potassium per- 
 
 N 
 manganate present in i cc. of its — V.S. 
 
140 THE ESSENTIALS OF VOLUiMETRIC ANALYSIS 
 
 The process is conducted as follows: Into a 200-cc. flask 
 place about 0.5 gm. of potassium iodid and 10 cc. of diluted 
 sulphuric acid, add to this (slowly from a burette) exactly 
 10 cc. of the permanganate solution to be standardized and 
 dilute the mixture (which is brown in color, because of the 
 liberated iodin) with distilled water to about 150 cc. Then 
 slowly titrate (with constant stirring) with an accurately stand- 
 
 N 
 ardized — sodium thiosulphate until the color of the solution 
 
 is a faint yellow, then add a few drops of starch solution and 
 continue the titration until the color is discharged. Note the 
 number of cc. consumed' and dilute the permanganate with 
 distilled water so that equal volumes of the two solutions 
 correspond to each other. 
 
 Example. If 13 cc. of the thiosulphate solution were 
 required, then each 10 cc. of the permanganate solution must be 
 
 diluted to 13 cc. 
 
 Standardization with Ferrous Annnoniiim Suljfhate 
 
 {Mohr's salt) (FeSOi- (NH4)2S04.6H20). 392.14 gms. of this 
 salt contains 55.82 gms. of iron (3.512 gms. contain 0.5 gm. of 
 iron). 3.512 gms. of the salt are accurately weighed out 
 and dissolved in sufficient recently boiled distilled water to 
 make 250 cc. Fifty cc. of this solution containing o.i gm. of 
 iron are transferred to a small flask, 10 cc. of diluted sul- 
 phuric acid added, and then the permanganate solution to 
 be standardized is run in slowly until a faint pinkish tint 
 appears. "WTiatever number of cc. is consumed that number 
 represents o.i gm. of iron, and must be diluted to 17.91 cc. 
 to make the solution exactly decinormal. 
 
 Volumetric Analyses by Means of Potassium Permanganate 
 
 When potassium permanganate solution is added to a solu- 
 tion of any readily oxidizable substance strongly acidulated 
 with sulphuric acid, it undergoes reduction, as shown in the 
 
ANALYSIS BY OXIDATION AND REDUCTION 141 
 
 equation below. The molecule (2KMn04) has eight atoms 
 of oxygen which it gives up in the process of oxidation. These 
 eight atoms of oxygen unite with the replaceable hydrogen 
 of an accompanying acid, liberating an equivalent amount 
 of acidulous radical. Three of these atoms of oxygen liberate 
 suflScient acidulous residual to combine with the potassium 
 and manganese of the permanganate, while the other five 
 atoms are a^'ailable for direct oxidation. 
 
 2KMn04+3H2S04=K2S04 + 2MnS04f3H20 + sO, 
 
 or, for combination with the hydrogen of more acid, more 
 acidulous residual being set free, to combine with the salt 
 acted upon. 
 
 2KMn04 + 8H2S04 = K2S04 + 2MnS04 + 8H20 + 5(804). 
 
 5(804) when combined with ioFeS04 forms reio(S04)i5 
 or 5Fe2(S04)3, ferric sulphate. Thus it is seen that one 
 molecule of potassium permanganate (2KMn04) has the 
 power of converting 10 molecules of a ferrous salt to the 
 ferric state. 
 
 The equation in full is 
 
 loFeSOi + 2KMn04 +8H2SO4 
 
 = K28O4 + 2MnS04 + 8H2O + 5Fe2 fS04)3. 
 
 We have seen that 2KMn04 has 5 atoms of oxygen 
 available for oxidizing purposes, and that each of these will 
 combine with 2 atoms of hydrogen. 2KMn04 is consequently 
 chemically equivalent to 10 atoms of hydrogen, and a normal 
 solution of this salt when used as an oxidizing agent is one 
 that contains in one liter one-tenth of the weight of 2KMn04 
 expressed in grams, and a decinormal solution, one which 
 contains one-hundredth of this weight. 
 
 As before stated, when potassium permanganate is brought 
 
142 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 in contact with a ferrous salt or other oxidizable substance, 
 it is decomposed and decolorized. Hence when titrating with 
 a standard solution of this salt it is decolorized so long as an 
 oxidizable substance is present. As soon, however, as the 
 oxidation is completed the standard solution retains its color 
 when added to the substance, and the first appearance of a 
 faint red color is the end-reaction, and the oxidation is laiown 
 to be completed. 
 
 In titrating with potassium permanganate it must be 
 remembered that excess of free acid (preferably sulphuric) 
 should always be present in the solution titrated, in order 
 to keep the resulting manganous and manganic oxids in solu- 
 tion; these, forming a dense brown precipitate, would make 
 it difficult if not quite impossible to recognize the pinkish 
 color of the end-reaction. Sulphuric acid alone, if in large 
 excess, has a reducing effect upon potassium permanganate. 
 
 Nitric and hydrochloric acids are prejudicial and should 
 be avoided; they are, however, frequently present in salts 
 which are to be analyzed, and in such event should be 
 removed by converting them into sulphate. By adding a 
 small excess of sulphuric acid and applying heat, until hydro- 
 chloric acid or nitrous vapors are no longer evolved, the 
 chlorid or the nitrate is converted into sulphate, and the 
 deleterious effect of their presence overcome. Hydrochloric 
 acid, unless present in very small quantities, and the titra- 
 tion conducted at a low temperature, will vitiate the analysis 
 through its action upon the permanganate whereby chlorin 
 is liberated,* thus 
 
 KMn04 + 8HCl = KCl-(-MnCl2-h4H20-H5Cl. 
 
 * This decomposition of the permanganate by hydrochloric acid is due 
 to the presence of ferric salt, which latter seems to act catalytically, for oxalic 
 acid may be accurately titrated with permanganate even in the presence of 
 hydrochloric ifcid, no chlorin being given off. Thus the decomposition of 
 the permanganate is not due to the hydrochloric acid alone. 
 
ANALYSIS BY OXIDATION AND REDUCTION 143 
 
 A very convenient way of obviating the irregularities due 
 to the presence of hydrochloric acid is to add a few grams 
 of manganous sulphate * to the solution before titrating it. 
 
 Mercuric sulphate f and magnesium sulphate may also 
 be used with satisfactory results. 
 
 Potassium permanganate, being so readily decomposed by 
 contact with organic matter, should be protected from such 
 contact. It should never be filtered through paper (glass- 
 wool or gimcotton may be used), nor should it be used in 
 a Mohr's burette or in any other apparatus in which it is 
 in contact with rubber or cork. Furthermore, all substances 
 of an oxidizing or reducing nature, aside from that being 
 analyzed, must be excluded from the solution. Among such 
 substances may be mentioned hydriodic acid, sulphureted 
 hydrogen, nitrous acid and the lower oxids of nitrogen, phos- 
 phorous and hypophosphorous acids, thiosulphuric, sulphurous, 
 and all the other acids of sulphur except sulphuric, also ous 
 salts and the metallic suboxids and peroxids. 
 
 Burettes and other apparatus which have been used for 
 
 * Kessler and Zimmermann suggest using 20 cc. of a solution of man- 
 ganous sulphate (200 gms. per liter). 
 
 X Cady and Ruediger (J. A. C. S., XIX, 575) concluded from the following 
 general principles that it is possible to titrate iron with permanganate in the 
 presence of hydrochloric acid if an excess of mercuric sulphate be added to 
 the solution. Mercuric halids in solution ionize to an extremely slight extent, 
 while the mercuric salts of oxyacids are readily ionized, since compounds of 
 slight ionization always result when their constituent ions meet; mercuric 
 halids are always produced when a mercuric salt of an oxyacid is added to a 
 solution containing halogen ions. Therefore when mercuric sulphate solution 
 and hydrochloric acid are mixed, ionization of both occurs, and the mercuric 
 ions unite with the chlorin ions and produce mercuric chlorid which is only 
 very slightly ionized. In the presence of a large excess of mercuric sulphate, 
 the mercuric ions resulting from its dissociation diminish the ionization of 
 the mercuric chlorid until it is practically nil. Thus no chlorin ions will be 
 present in the solution to induce decomposition of the permanganate. 
 
144 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 permanganate, should be emptied and rinsed immediately 
 after use, and any manganic oxid which may be adhering 
 to the glass should be removed by means of hydrochloric 
 acid and boiled water. 
 
 Not only oxidizable substances but reducible substances 
 may be estimated by means of potassium permanganate. 
 
 In the estimation of oxidizable substances the standard 
 potassium permanganate is added directly to the acidulated 
 solution of the substance being analyzed. The completion of 
 the oxidation being then known by the appearance of a faint 
 pinkish tint. This is the direct method. 
 
 In the estimation of reducible substances (i.e., oxidizing 
 substances) the indirect or the residual method is employed. 
 
 In this an accurately weighed or measured quantity of 
 the substance is brought together with an excess of a third 
 substance having reducing power, and which is similarly 
 effected by the permanganate and by the substance analyzed. 
 After completion of the reaction the excess of the reducing 
 substance is found by titration with standard permanganate. 
 T^he difference between the quantity so found and that originally 
 added gives the quantity which reacted with the salt under 
 analysis, and from this the calculation is made. 
 
 On the Use of Empirical Permanganate Solutions. 
 
 A. If the standardization of the solution is done by means 
 of iron, as described on page 136, o.i gm. of iron wire (repre- 
 senting 0.0996 gm. of pure iron) will require 17.84 cc. of the 
 permanganate solution if the latter is exactly decinormal. 
 If less than this quantity of solution is used (say r7.5 cc.) 
 it indicates that the solution is stronger than decinormal, 
 and may either be diluted so that each 17.5 cc. will measure 
 17.84, or it may be used as it is. This latter is in most cases 
 
ANALYSIS BY OXIDATION AND REDUCTION 145 
 
 preferable. The value of i cc. of the solution in iron is 
 calculated thus: 
 
 17.5 cc. : I cc. : : 0.0996 gm. : x. 51; =0.00569 gm. 
 
 If a solution of this strength is to be used for the estima- 
 tion of iron, simple multiplication of the number of cc. used 
 by 0.00569 gm. gives the weight of Fe present. If, however, 
 this solution is employed for the titration of other oxidizablc 
 substances, the number of cc. consumed is multiplied by 
 0.00569 gm. and then by a fraction in which the numerator 
 represents a quantity of the substance examined, equivalent 
 in grams to an atom of iron in its reaction with permanganate, 
 and the denominator is the atomic weight of iron. 
 
 Example. If in a titration we use 40 cc. of a perman- 
 ganate solution, the titer of which has been found to be i cc. 
 = 0.00569 gm., the calculation would be: 
 
 in the case of ferrous sulphate (FevSOa, 151.89), 
 
 151.89 
 40X0.00569 gm.X =0.6192 gm.; 
 
 in the case of oxalic acid (H2C2O4. 21120=126.05), 
 
 63 
 40X0.00569 gm.X— -g^= 0.2 56 gm.; 
 
 in the case of hydrogen dioxid (H20o = 34), 
 
 17 
 40X0.00569 gm.X—— = 0.0693 gm. 
 
 B. Another way. The solution just mentioned, of which 
 17.50 cc. are consumed in titrating o.i gm. of iron wire, is 
 compared with a true decinormal permanganate solution, of 
 which 17.84 cc. are consumed in the same reaction. Tho 
 
146 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Strength of the former sokition is therefore as compared 
 
 with a decinormal solution. 
 
 In titrating with this solution the number of cc. consumed 
 
 are to be multiplied by -^-;^ and then by the true decinormal 
 
 factor for the substance being analyzed. 
 
 Example. 40 cc. of the solution are consumed. 
 
 In the case of ferrous sulphate (FeS04= 151.89), the deci- 
 normal factor (i.e., the weight of ferrous sulphate represented 
 by I cc. of a true decinormal solution) is 0.015189 gm. 
 
 1984 
 40X — — X0.01S189 gm. =0.0192 gm. 
 
 In the case of oxalic acid (H2C2O4, 2H2O = 126.05) j ^^ deci- 
 normal factor is 0.0063 g™- 
 
 1784 
 40X X0.0063 gm. =0.256 gm. 
 
 In the case of hydrogen dioxid (H2O2 = 34) , the decinormal 
 factor is 0.001688 gm. 
 
 1784 
 40 X X0.0017 gm. =0.0693 gin- 
 
 C. If the standardization is done by means of oxalic acid, 
 as described on page 138, in which 10 cc. of a strictly deci- 
 normal oxalic acid solution are titrated with the permanganate 
 solution which is being standardized, exactly 10 cc. of the 
 latter will be consumed if it is of decinormal strength. If 
 in the trial, however, it is found that only 9.6 cc. are con- 
 sumed it indicates that the solution is stronger than decinormal; 
 
 100 
 its strength being expressed by —7-. If, on the other hand, 
 
ANALYSIS BY OXIDATION AND REDUCTION 147 
 
 more than lo cc. of the solution are consumed (say 10.4 cc.) 
 
 100 
 the solution is below decinormal strength, namely, 
 
 In using a solution of the first strength the number of 
 
 100 
 cc. of it consumed in any titration is to be multiplied by —7- 
 
 and then by the decinormal factor for the substance examined. 
 In the case of the weaker solution the number of cc. con- 
 
 TOO 
 
 sumed is multiplied by and then by the decinormal factor 
 
 for the substance being analyzed. 
 
 Examples. Ferrous sidphate (FeS04= 151.89) is titrated 
 with the stronger solution, 40 cc. of the latter being consumed. 
 
 100 
 Then 40X— 7-X0.015189 gm.=o.628 gm. 
 
 Oxalic acid (H2C204.2H20= 126.05), 40 cc. are consumed. 
 
 100 
 Then 40X— — X0.0063 gm. =0.260 gm. 
 
 Hydrogen dioxid (H202 = 34), 40 cc. are consumed. 
 
 100 
 Then 40X— — X0.0017 gm. = 0.0703 gm. 
 
 D. If the checking of the permanganate solution is done 
 
 by the iodometric method (page 139) and it is found that 
 
 10 cc. of the permanganate requires the use of 13 cc. of 
 
 decinormal thiosulphate solution, the titer of the solution 
 
 13 
 is expressed with reference to decinormal as — . In using 
 
 a solution of this strength, the number of cc. of it consumed 
 
 13 
 in an analysis is multiplied by — and then by the decinormal 
 
 factor for the substance analyzed. 
 
148 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 TYPICAL ANALYSES WITH PERMANGANATE 
 
 A. Direct Titrations 
 
 a. Estimation of Ferrous Sulphate (FeS04 + 71^20 = 278). 
 One gm. of ferrous sulphate is dissolved in 25 cc. of water 
 and the solution strongly acidulated with sulphuric acid. Deci- 
 normal potassium permanganate is then delivered from a 
 burette until a permanent pink tint is obtained, indicating 
 the complete oxidation of the ferrous salt. 
 
 The reaction is as follows: 
 
 (ioFeS04 + 7H2O) 
 
 100)2780 
 
 27.80 gms. = 
 
 + 
 
 2KMn04 
 
 100)316.06 
 3.1606 
 
 + 8H2SO4 
 
 N 
 gm. = 1000 cc. ■^- V.S. 
 10 
 
 2.780 gms. = 
 
 
 
 N 
 100 cc. — V.S. 
 10 
 
 0.0278 gm. = 
 
 
 
 . cc. ^ V.S. 
 
 10 
 
 = 5Fe2(S04)3 + K2SO4 + 2MNSO4 + 78H2O. 
 
 Thus 316.06 gms. of permanganate = 2780 gms. of crys- 
 tallized ferrous sulphate, which equals 55.82 gms. of metallic 
 
 N 
 iron. One cc. of — permanganate solution therefore presents 
 
 0.0278 gm. of FeS04 + 7H20 or 0.005582 gm. of Fe. 
 
 N 
 
 In the analysis 35 cc. of the — permanganate were con- 
 sumed. The I gm. taken then contains 35X0.0278 = 0.973 
 gm. or 97.3 per cent. 
 
 If it is desired that each cc. of the permanganate solution 
 should represent a certain percentage of pure salt, a molecular 
 quantity of the salt should be taken for analysis instead of 
 I gm. For example, if 2.78 gms. be taken, each cc. of the 
 decinormal solution consumed will correspond to i per cent, 
 because 2.78 gms. is the weight of crystallized ferrous sul- 
 
ANALYSIS BY OXIDATION AND REDUCTION U9 
 
 phate which can be oxidized by loo cc. of the decinormal 
 solution. If half of this weight be taken, i.e., 1.39 gms., each 
 cc. of the permanganate solution compound will represent 
 2 per cent of pure salt. 
 
 Granulated Ferrous Sulphate (FeS04 + 7H20) is estimated 
 in the same way as the foregoing, and should correspond 
 with it in strength. 
 
 Exsiccated {Dried) Ferrous Sulphate. This salt is tested 
 in the same manner as the other two sulphates. It contains 
 a larger percentage of ferrous sulphate than the other two, 
 having less water of crystallization. Its composition is approx- 
 imately FeS04+3H20. 
 
 loFeSOi + 2KMn04 + 8H2SO4 
 
 100)1518.9 100)^16.06 
 
 15.189 gms. 3.1606 gms. or looo cc. — standard solulion. 
 
 = 5Fe2(S04)3 + K2SO4 + 2MnS04 + 8H2O. 
 
 Each cc. of the standard solution represents 0.015189 gm. 
 of anhydrous (real) ferrous sulphate. If one gm. of the dried 
 
 N 
 salt, treated as above described, requires 48 cc. of — per- 
 manganate solution, it contains 
 
 0.01518948 = 0.72808 gm., 
 
 or 72.80 per cent of real ferrous sulphate, and 100.00—72.80 
 = 27.20 per cent of water of crystallization. 
 
 Saccharated Ferrous Carbonate. Two gms. of saccharated 
 ferrous carbonate are dissolved in 20 cc. of 10 per cent sulphuric 
 acid, diluted with water to about 100 cc, and then titrated 
 
 N 
 with — potassium permanganate V.S. imtil a pink tint is 
 
 produced in the liquid. This method is not an exact one, 
 especially if heat is applied for solution of the powder, in 
 that permanganate is reduced by the sugar present. 
 
150 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 b. Estimation of Metallic Iron in Ferrum Reductum. Fer- 
 rum reductum (reduced iron) always contains besides metallic 
 iron a varying quantity of oxid. Therefore, in assaying this 
 preparation a method must be employed which will estimate 
 the iron only, which is present as metallic iron. This may 
 be done by means of a solution of mercuric chlorid which 
 reacts with metallic iron only and not with the oxid. 
 
 The method is as follows : 
 
 0.55 gm. of reduced iron is introduced into a glass-stoppered 
 bottle, 50 cc. of mercuric chlorid solution (5 gms. in 100 cc.) 
 are added and the bottle heated on a water bath for one hour, 
 agitating frequently, but keeping the bottle well stoppered. . 
 
 2HgCl2 + Fea = 2FeCl2 + 2Hg, 
 
 then allow it to cool, dilute the contents with water to 100 
 cc, and filter. Take 10 cc. of the filtrate (representing 0.055 
 gm. of reduced iron) add to it 10 cc. of diluted sulphuric acid and 
 10 cc. of a solution of manganous sulphate (1:5), introduce 
 the mixture into a glass-stoppered bottle (having a capacity 
 of 100 cc), and titrate with decinormal permanganate imtil 
 a permanent pink color is obtained. Each cc. of the per- 
 manganate solution represents 0.005582 gm. of metallic iron. 
 
 loFe S04+2KMn04 + 8H2SO4 
 
 100)558.2 100 1316.06 -- 
 
 5.582 gms. = 3.1606 gms. = iooo cc. — V.S. 
 
 N 
 0.005582 gm.-— 0.0031606 giTi. = i cc. — V.S. 
 
 = sFe2(S04)3 + K2SO4 + 2MnS04 + 8H2O. 
 
 N 
 If 9 cc. of the — permanganate are consumed, then 
 
 9X0.005582=0.0502 gm., and since the 10 cc. of the iron 
 
ANALYSIS BY OXIDATION AND REDUCTION 151 
 
 solution taken for analysis represented 0.055 S'^-> th^ P'^'' 
 cent of metallic iron present is 91. 
 
 0.055:100: : 0.0502: Jt. x = 9i. 
 
 The use of manganous sulphate in this process is to prevent 
 decomposition of permanganate by the hydrochloric acid. 
 The quantity is, however, so small, that if the titration be 
 conducted cold, its use is unneccesary. 
 
 Titration with an Empirical Permanganate Solution. A 
 solution of permanganate, which is found upon standardiza- 
 tion to be of a strength in which i cc. is equivalent to 0.00512 
 gm. of Fe, is to be used. 
 
 Each cc. of this solution is equivalent to the following 
 quantities: 
 
 FeS04 0.01393 gm- 
 
 FeS04 + 7H2O 0.02549 " 
 
 FeCOs 0.01062 " 
 
 FeCla 0.01162 " 
 
 Fe 0.00512 " 
 
 c. Estimation of Oxalic Acid and Oxalates with Potassium 
 Permanganate Solution (H2C2O4 + 21120 = 126.05; H2C2O4 
 = 90). The estimation of oxalic acid may be accurately 
 made either by neutralization with a standard alkali or by 
 oxidation with standard permanganate. The latter method 
 is, however, the one to be employed in the case of oxalates. 
 
 The oxidimetric estimation of oxalic acid is carried out 
 as follows: 
 
 One gm. of the acid (accurately weighed) is dissolved in 
 sufficient water to make 100 cc. Of this solution, 10 cc. 
 (representing o.i gm. of the acid) is taken for. analysis. Two 
 cc. of diluted sulphuric acid are added, the solution is heated 
 to between 40° C. and 60° C, and keeping it at about this 
 
152 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 temperature, is titrated with decinormal potassium perman- 
 ganate, agitating constantly, until a faint rose tint marks 
 the completion of the reaction. 
 
 Each cc. of the permanganate solution consumed repre- 
 sents 0.0063 gm. of crystallized oxalic acid. 
 
 The reaction is as follows: 
 
 S(H2C204 + 2H20) +3H2S04 + 2KMn04 
 
 100)6,^0.25 100)316.06 
 
 N 
 
 6.30 gms. = 3.1606 gms. = iooo cc. — V.S. 
 
 10 
 
 = K2S04 + 2MnS04 + 8H20 + 502. 
 
 Direct Percentage Titration. 0.63 gm. of crystallized oxalic 
 
 N 
 acid is oxidized by 100 cc. of — permanganate. Therefore 
 
 N 
 if 0.63 gm. of the acid is taken for analysis, each cc. of — 
 
 permanganate will represent i per cent. 
 
 Titrating with an Empirical Solution. If the permanganate 
 
 is checked with iron, we take into consideration that 2KMn04 
 
 will oxidize 10 atoms of iron (558.2 parts), and on the other 
 
 hand 5 molecules of oxalic acid (630.25 parts). If the titer of 
 
 the permanganate be found on experiment to be i cc. =0.00569, 
 
 whatever number of cc. of this solution is consumed is to 
 
 63 
 be multiplied by 0.00569 and then by — ^. 
 
 Example. 0.3 gm. of oxalic acid require for oxidation 
 40 cc. of a permanganate solution whose titer is i cc. = 0.00569 
 gm. Fe, the calculation is made as follows : 
 
 63 
 40X0.00569 gm.X— -^^ = 0.256 gm. 
 
 0.256 gm. is the quantity of pure crystallized oxalic acid 
 
ANALYSIS BY OXIDATION AND REDUCTION 153 
 
 present in the 0.3 gm. taken for analysis. This is 85.3 per- 
 cent. 
 
 0.21^6X100 
 
 — ^ = 85.3. 
 
 0.3 ^-^ 
 
 If the standardization of the permanganate is done by 
 means of a decinormal oxaHc acid, or by the iodometric method, 
 the calculation is as described on pages 146-147. 
 
 Oxalates are estimated in the same manner; a much 
 larger quantity of sulphuric acid is, however, required. This 
 serves to liberate the oxalic acid from its combination. 
 
 The presence of precipitates of sulphates of calcium, barium, 
 or lead does not interfere with the recognition of the end- 
 point. 
 
 iLach cc. of — potassium permanganate represents 
 
 Oxalic acid anhydrous (H2C2O4) 0.0045 o™- 
 
 Oxalic acid crystallized (H2C2O4 + 2H2O) . . .0.0063 " 
 
 This method may be applied to calcium salts which are 
 soluble in water or in acetic acid, as well as to other metals 
 which are precipitable as oxalates. It is, however, especially 
 applicable to calcium, because of the readiness with which 
 this metal may be separated from others as oxalate. The 
 precipitation may be accomplished in either an ammoniacal 
 or a weak acetic acid solution. If it is necessary to dissolve 
 the calcium salt with the aid of hydrochloric acid, the solution 
 must be rendered strongly alkaline with ammonia, and the 
 precipitation effected with ammonium oxalate. 
 
 Calcium Carbonate may also be estimated by this method. 
 The precipitation is effected either in ammoniacal or weak 
 acetic acid solution. 
 
 Example (a). 0.2 gm. of calcium carbonate is placed in 
 a beaker. 20 cc. of water are added, and then sufficient strong 
 
154 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 hydrochloric acid to effect solution. The beaker should be 
 covered with a watch-glass to prevent loss of contents during 
 effervescence and the liquid should be heated to boiling in 
 order to remove CO2. The liquid is then diluted with water 
 to 150 cc. and made alkaline by the addition of ammonia 
 water. Ammonium oxalate solution is then added, drop by 
 drop, until precipitation is complete, the mixture boiled for 
 about 3 minutes and set aside for several hours. 
 
 The precipitate is then washed with hot water containing 
 some ammonium oxalate, transferred to a filter and the washing 
 continued with cold water until free from soluble oxalate. 
 The washed precipitate, together with the filter, is then trans- 
 ferred to a beaker, some sulphuric acid added, the mixture 
 
 N 
 
 diluted to iKo cc. with distilled water and titrated with — 
 
 ■^ 10 
 
 potassium permanganate V.S. 
 
 N 
 Each cc. of — permanganate V.S. represents 
 
 0.002004 gm. of Ca, 
 0.002804 gm. of CaO, 
 0.005004 gm. of CaCOs. 
 
 Example (b). 0.2 gm. of calcium carbonate is dissolved 
 in the smallest necessary quantity of dilute acetic acid, then 
 sufficient oxalic acid is added to completely precipitate the 
 calcium as oxalate (CaC204). This precipitate is then 
 thoroughly washed on the filter. A hole is then made in 
 the filter and the precipitate washed through the funnel into 
 a flask (about 100 cc. of water being used), a small quantity, 
 say about 2 cc, of dilute sulphvu-ic acid are added, the mixture 
 warmed to between 40° C. and 60° C. and titrated with 
 decinormal potassium permanganate until a permanent rose 
 tint appears. 
 
ANALYSIS BY OXIDATION AND REDUCTION ITj") 
 
 Each cc. of the decinormal permanganate represents 
 0.0045 S^- of oxalic acid (anhydrous) or 0.005004 gm. of cal- 
 cium carbonate. 
 
 d. Estimation of Hydrogen Dioxid and Barium Dioxid with 
 Standard Potassium Permanganate. Hydrogen dioxid {Hydro- 
 gen peroxid) (Hi.02 = 34). Hydrogen dioxid and potassium 
 permanganate, though both oxidizing agents, will, when mixed 
 in an acid solution, reduce each other. The reaction which 
 occurs is probably primarily an oxidation of the H2O2 to a 
 higher oxid (H2O4 (?)) which, however, immediately breaks 
 up with the liberation of oxygen. The method of assaying 
 hydrogen dioxid by means of permanganate is applicable not 
 only to this substance but also to the estimation of barium 
 dioxid and the soluble alkali peroxids. The method is usually 
 carried out by adding the permanganate solution to the dioxid 
 in a solution acidulated with sulphuric acid. Immediate 
 decolorization of the permanganate occurs, as long as any 
 hydrogen dioxid is present. When the latter has been entirely 
 taken up the permanganate is no longer decolorized and a 
 faint pink tint marks the end-point. In the estimation of 
 the pharmacopoeial or commercial dioxid solutions, containing 
 2 or 3 per cent of H2O2, a measured quantity is taken for 
 analysis. The specific gravity of the solution, being nearly 
 that of water, i cc. is taken to represent i gm. In the case 
 of solutions of hydrogen dioxid of high percentage strength, 
 it is advisable to take a weighed quantity for analysis. If 
 hydrochloric acid is present a small quantity of manganese 
 sulphate should be introduced before titrating. 
 
 The assay is conducted as follows : 
 
 Ten cc. of the solution are accurately measured and diluted 
 (in a graduated cylinder) with water to make 100 cc. Ten cc. of 
 this diluted liquid (containing i cc. of hydrogen dioxid solu- 
 tion) are transferred to a beaker, 5 cc. of diluted sulphuric 
 
156 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 acid (i-8) are added and then the — permanganate solution 
 
 run in from a burette, stirring after each addition until a 
 permanent faint pink tint appears. The reaction is as follows: 
 
 5H2O2 + 2KMn04 + 3H2SO4 
 
 100)170 100)316.06 
 
 1.7 gms. 3.1606 gms. = iooo cc. — permanganate V.S. 
 
 0.0017 gm. = I cc. — permanganate VS. 
 
 = K2SO4 + 2MnS04 + 8H2O + SO2. 
 
 N 
 Thus each cc. of — permanganate represents 0.0017 S^^- 
 
 of absolute hydrogen dioxid. Assuming that in the above 
 estimation 19 cc. of the permanganate solution were required, 
 then the i cc. taken for analysis contained 0.0017 g™-^i9) 
 which is 0.0323 of absolute H2O2. This corresponds to 3.23 
 per cent. 
 
 The Direct Percentage Method. Ten cc. of the solution are 
 diluted with water to measure 100 cc. Seventeen cc. of this 
 diluted solution (containing 1.7 cc. of hydrogen dioxid) are 
 acidulated with sulphuric acid and titrated with decinormal per- 
 manganate, as above described. Each cc. of the permanganate 
 solution consumed will represent o.i per cent by weight of H2O2. 
 
 Titration with an Empirical Solution. A permanganate 
 solution is on hand which is found upon standardization with 
 iron to be i cc. =0.00569 gm. Fe. To use this solution as 
 it is, we take into consideration that 2KMn04= (316.06) = 10 
 atoms of Fe (558.2) and also 5 molecules of H2O2 (170). 
 31.606 gms. KMn04, = 55.82 gms. Fe, = 17 gms. H2O2. What- 
 ever number of cc. of this permanganate solution is used, 
 
 17 
 multiplied by 0.00569 gm. and then by — —, will give the 
 
 55-°2 
 
 weight of H2O2 present in the sample analyzed. 
 
ANALYSIS BY OXIDATION AND REDUCTION 157 
 
 Estimation of Volume Strength. Let us look at the above 
 equation in a different light. 
 
 We see that when potassium permanganate and hydro- 
 gen dioxid react, lo atoms of oxygen are liberated. 
 
 The permanganate itself when decomposed liberates five 
 atoms of oxygen. Therefore of the above ten atoms only 
 five come from the hydrogen dioxid. 
 
 5H202 = sH20 + 50; 
 2KMn04-l-3H2S04 = K2S04 + 2MnS04 + 3H20 + 50. 
 
 In order to find the factor for volume of available oxygen, 
 see the following equation, etc. : 
 
 SH2O2 + 2KMn04+3H2S04=K2S04 + 2MnS04 + 8H2O + sO +5O. 
 100)316.06 ^ 100)80 
 
 3.1606 gms. or 1000 cc. — V.S.= 0.80 gm. 
 
 N 
 I cc. — V.S.= 0.0008 gm. 
 
 10 " 
 
 Thus it seen that each cc. of — potassium permanganate 
 
 V.S. represents 0.0008 gm. of oxygen. But we require to 
 find the volume of oxygen, not the weight represented by 
 
 I cc. of — permanganate. 
 
 1000 cc. of oxygen at 0° C. and 760 mm. pressure, weigh 
 1.43 gms. Therefore, if 1.43 gms. measure 1000 cc, 0.0008 
 gm. will measure 0.57 cc. 
 
 The factor, then, for volume of oxygen liberated when 
 
 N 
 hydrogen dioxid is titrated with — potassium permanganate, 
 
 N 
 is 0.57, and the number of cc. of the — potassium permanganate 
 
 consumed in the titration gives the volume of oxygen liberated 
 by the quantity of hydrogen dioxid taken. 
 
158 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 Thus if 19 cc. of the — V.S. were required, 
 
 0.57X19=10.83 cc. of oxygen. 
 
 It is convenient to operate upon i cc. hydrogen dioxid 
 solution. Then each cc. of potassium permanganate V.S. used 
 will represent 0.57 cc. of available oxygen and is necessary 
 only to multiply the number of cc. by this factor to find the 
 volume of available oxygen. 
 
 If any other quantity ' than i cc. of dioxid be taken for 
 analysis, it will be necessary after multiplying by 0.57 to 
 divide the result by the quantity of dioxid solution taken, in 
 order to find volume strength. 
 
 Hydrogen dioxid solution may also be volumetrically 
 assayed by Kingzett's method, which is described under 
 lodometry. 
 
 The gasometric estimation is also described further on. 
 
 Barixun Dioxid {Barium Peroxid) (Ba02 = 169.37). This 
 substance is assayed by treating it with an acid, and then 
 estimating the liberated hydrogen dioxid, as follows: 
 
 Weigh off 2 gms. of the coarse powder, put it in a porcelain 
 capsule, add about 10 cc. of ice-cold water, then 7.5 cc. of 
 phosphoric acid (85 per cent), and sufficient ice-cold water 
 to make 25 cc. Stir and break up the particles with the end 
 of the stirrer xmtil a clear or nearly clear solution is obtained 
 and all that is soluble is dissolved. 
 
 Five cc. of this solution (which corresponds to 0.4 gm. 
 of barium dioxid) is measured off for assay. 
 
 Drop into this from a burette, with constant stirring, 
 decinormal potassium permanganate until a final drop gives 
 the solution a permanent pink tint. 
 
 About 40 cc. of the decinormal permanganate should l?e 
 required to produce this result. 
 
ANALYSIS BY OXIDATION AND REDUCTION 159 
 
 In this process, the first step is the formation of hydrogen 
 dioxid by treating the barium dioxid with phosphoric acid, 
 as illustrated by the following equation : 
 
 BaO^i + H3PO4 = BaHPOi + H2O2. 
 
 169-37 34 
 
 The hydrogen dioxid is then estimated with decinormal 
 permanganate, as described above. 
 
 S(Ba02) = 5H2O2 + 2KMn04 + 3H2SO4 
 
 100)845.85 100)170 100)316.06 
 
 8.4585 gms.= 1.7 gms. 3.1606=1000 cc. — permanganate. 
 
 = K2SO4 + 2MnS04 + 8H2O + 5O2. 
 
 e. Estimation of Ferric Salts by Means of Potassium Per- 
 manganate (after reduction). It is frequently necessary to 
 estimate ferric salts by means of permanganate solution; this 
 is particularly the case in compounds where ferric and ferrous 
 salts are both present. 
 
 The ferric salts must of course be reduced to the ferrous 
 state in order to estimate them with permanganate, or in 
 fact with any oxidizing agent. There are many ways of 
 affecting this reduction, but the best way (where permanganate 
 is to be used) is no doubt by the use of metallic zinc or mag- 
 nesium in sulphuric acid solutions. Hydrochloric acid may 
 be used instead of sulphuric, but in that case the solution 
 must be very dilute and of low temperature, in order to avoid 
 the liberation of chlorin, which would spoil the analysis. 
 
 In concentrated or hot solutions hydrochloric acid acts 
 upon permanganate as a reducing agent, as shown in the 
 equation 
 
 KMn04+8HCl = KCl+MnCl2+4H20 + 5Cl. 
 
 The irregularities due to the liberation of chlorin may 
 be obviated by the addition of an excess of mercuric sulphate 
 
160 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 before titration, as suggested by Cady and Ruediger, or by 
 the use of magnesium or manganous sulphate, as suggetsed by 
 Kessler and Zimmermann (see page 143). 
 
 The reduction is effected by adding to the warm diluted 
 solution of the ferric salt acidulated with sulphuric acid small 
 pieces of pure metallic zinc or coarsely powdered magnesium, 
 and setting aside in a covered vessel imtil the solution is 
 colorless, or until it fails to produce a red color when a drop 
 of it is brought in contact with a drop of sulphocyanate. 
 The zinc used must be free from iron, or if the latter metal 
 is present its quantity must be known. All of the zinc or 
 magnesium must be dissolved before the titration is begun, 
 otherwise the reduction would continue while the titration 
 is being done. When the ferric salt is completely reduced 
 the titration should be carried out without delay in order 
 to avoid reoxidation through exposure to the air. Reoxida- 
 tion takes place more readily in the presence of hydrochloric 
 acid than of sulphuric acid. According to D. J. Carnegie, 
 reduction takes place much more rapidly in neutral than 
 in acid solutions. He suggests neutralization of the solu- 
 tion with ammonia, and after reduction the addition of sul- 
 phuric acid to keep the iron in solution. Other methods 
 for the reduction of ferric salts are described further on. 
 
 The solution of the ferric salt to be estimated should not 
 contain more than 0.15 gm. of metallic iron in 250 cc. To 
 this quantity of solution about 10 gms. of metallic zinc and 
 25 cc. of sulphuric acid are taken. The solution is kept at a 
 temperature between 60° and 80° C. imtil the zinc is entirely 
 dissolved, then the mixture is boiled in a flask provided with 
 a valve stopper, as shown in Fig. 41 (in order to exclude air 
 and prevent reoxidation). It is then rapidly cooled, and 
 titrated with permanganate without delay. 
 
 Example. Estimation of Ferric Chlorid. 0.35 gm. of the 
 
ANALYSIS BY OXIDATION AND REDUCTION 161 
 
 dried salt are dissolved in 250 cc. of water in a flask, 25 cc. 
 of sulphuric acid are added, and then 10 gms. of metallic 
 zinc are introduced. The flask is then gently warmed until 
 the zinc is entirely dissolved and the solution is colorless and 
 fails to give a red color when a small portion of it, removed 
 with a glass rod, is brought in contact with a drop of potas- 
 sium sulphocyanate. The solution is then brought to a boil, 
 and after this rapidly cooled, avoiding entrance of air, and 
 
 N 
 when cool titrated with — potassium permanganate, until a 
 
 faint permanent pink color appears. 
 
 19.5 cc. of the permanganate were required. Each cc. of 
 
 N „ , „. . 
 
 — permanganate represents 0.005582 gm. of metallic iron, 
 
 which is equivalent to 0.01622 gm. of anhydrous ferric chlorid. 
 Then if 19.5 cc. of the permanganate were employed, the 
 quantity of real ferric chlorid present in the sample is 0.01622 
 gm. X19.5, while the quantity of metallic iron present is 
 0.005582 gm.Xi9.5. 
 
 The reactions are represented by the following equations: 
 
 FeaCle + Zn = 2FeCl2 + ZnCb 
 
 loFeCla + 2KMn04 + 16HCI 
 
 = 5Fe2Cl6 + 2KCI + 2MnCl2 + 8H2O. 
 
 Because of the frequent and almost invariable presence in 
 zinc of carbon and iron, which have a reducing action upon 
 permanganate, it is necessary to carry out a blank experiment, 
 to determine the quantity of permanganate solution used up 
 by these impurities in the zinc. This blank experiment must 
 be conducted under the same conditions as the assay, and 
 differs only in that the iron is left out. 
 
 Example. Ten gms. of zinc from the same lot as used for 
 the assay is treated with 250 cc. of water and 25 cc. of sul- 
 
162 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 phuric acid, and when it is completely dissolved, the potassium 
 permanganate solution is added, until a permanent pale pink 
 tint results. The number of cc. consumed is deducted from 
 the quantity employed in the assay; the difference is the 
 quantity of the permanganate solution which was consumed 
 by the ferrous chlorid. Another methjd for the reduction of 
 ferric salts, previous to titration with permanganate, is that 
 of N. Matolcsy. In this the ferric salt is precipitated with 
 ammonium sulphid and the precipitated ferrous sulphid dis- 
 solved in sulphuric acid, which converts it into ferrous sul- 
 phate. This is then titrated with permanganate after the H2S 
 has been driven off. 
 
 /. Estimation of Nitrous Acid and Nitrites. Nitrous acid, 
 when brought in contact with a potassium permanganate solu- 
 tion acidulated with sulphuric acid, is oxidized to nitric acid. 
 Two molecules of KMn04 reacting with 5 molecules of HNO2, 
 as the equation shows, 
 
 5HNO2 + 2KMn04 + 3H2SO4 
 
 --= 5HNO3 -l-K2S04-l-2MnS04-h3H20. 
 
 In the case of nitrites, as for example sodium nitrite, the 
 oxidation takes place in the same manner, and the process 
 may be applied with equally good results to the salts, as well 
 as to free HNO2. At ordinary temperatures the oxidation 
 proceeds very slowly, but at a temperature of 40° C. (104° F.) 
 rapid reaction occurs. But because of the volatility of nitrous 
 acid in acidulated solutions of its salts it is. impossible to 
 accurately estimate them by direct titration with permanganate 
 at a raised temperature. 
 
 It is customary to add the nitrite solution to a measured 
 volume of warm acidulated standard permanganate solution. 
 The nitrite is then oxidized immediately as it comes in con- 
 tact with the permanganate, and each drop added makes the 
 
ANALYSIS BY OXIDATION AND REDUCTION 103 
 
 permanganate lighter in color, and when complete decolor- 
 ization of the permanganate is attained, the reaction is at 
 an end. 
 
 The process in detail is as follows: 
 
 loo cc. of — potassium permanganate are measured mto 
 
 a flask, 5 cc. of diluted sulphuric acid (1:5) are added, and 
 the solution warmed to 40° C. (104° F.). A solution of the 
 nitrite (say sodium nitrite) is now prepared by dissolving i 
 gm. of the salt in sufficient water to make 100 cc. This solu- 
 tion is placed in a burette and delivered slowly into the acidu- 
 lated permanganate solution, with constant shaking, and reduc- 
 ing the flow to drops towards the end of the titration, until 
 the permanganate is completely decolorized. We will assume 
 that 38 cc. of the NaN02 solution were used. 
 
 SNaNOa -h 2KMn04 + 3HaS04 
 
 100 )345 -OS 100 )316.06 j^ 
 
 3.4505 gms. 3.1606 gms. = iooo cc. — permanganate. 
 
 = sNaNO., -I- K2SO4 + 2MnS04 + 3H2O. 
 
 N 
 Thus 100 cc. of — permanganate represent 0.34505 gm. 
 
 of pure NaN02. Therefore if 38 cc. of the sodium nitrite 
 
 N 
 solution decolorized 100 cc. of — permanganate, the 38 cc. 
 
 must contain 0.34505 gm. of NaN02. If 38 cc. contains 
 0.34505 gm., 100 cc. contains 0.902 gm. 
 
 38 : 0.34505 : : 100 : x. :x;=o.9o8 gm., 
 
 and since i gm. of salt was dissolved in 100 cc. of solution 
 the percentage of pure NaN02 is 90.8 per cent. 
 
 Nitrites may also be estimated by adding an excess of 
 acidulated permanganate solution, warming the mixture, and 
 retitrating with standard oxalic acid. 
 
161 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 To 30 cc. of — potassium permanganate add water to 
 
 make about 150 cc. of solution, then add 5 cc. of sulphuric 
 acid and 10 cc. of a solution (i gm. in 100 cc.) of the sodium 
 nitrite to be assayed, warm the mixture to 40° C. (104° F.) 
 
 N 
 and allow to stand for five minutes, then titrate with — oxalic 
 
 10 
 
 acid solution until complete decolorization is effected. Not 
 more than 3.75 cc. of the latter should be required. 
 
 N 
 The volume of — oxalic acid solution, deducted from the 
 10 
 
 30 cc. of — potassium permanganate solution used, gives the 
 
 quantity of the latter which reacted with the one gram of 
 
 sodium nitrite, each cc. of — permanganate represents 0.0034505 
 
 gm. of NaN02. 
 
 B. Residual Titrations. 
 
 a Methods in which an Excess of Standaed Permanganate is Added, 
 AND THE Excess Determined by Residual Titration with Stand- 
 ard Oxalic Acid. 
 
 Estimation of Hypophosphorous Acid and Hypophos- 
 phites. An accurately weighed quantity of the acid or its 
 salt is dissolved in water, the solution is strongly acidulated 
 
 N 
 with sulphuric acid, and then a measured excess of — potas- 
 sium permanganate solution added. The mixture is boiled 
 for fifteen minutes to hasten and facilitate the oxidation, and 
 then the excess of the permanganate solution estimated by 
 
 N . ". 
 residual titration with — oxalic acid solution. 
 
 ID 
 
 Hypophosphorous Acid (HPH202 = 66.04). Three gms. of 
 the acid accurately weighed are diluted with water to make 60 
 cc. Of this solution, 6 cc. (containing 0.3 gm. of the acid) 
 are carefully removed with a pipette, and introduced into a 
 
ANALYSIS BY OXIDATION AND REDUCTION 165 
 
 flask. Three cc. of sulphuric acid are added and then 50 cc. of 
 
 N 
 
 — potassium permanganate soKition, and the mixture boiled 
 
 for fifteen minutes. 
 
 The potassium permanganate, in the presence of sulphuric 
 acid, oxidizes the hypophosphorous acid to phosphoric, as 
 the equation shows: 
 
 5HPH2O2 + 6H2SO4 + 2(2KMn04) 
 
 2 )33°-^ 2)632^ 
 
 100)165.1 100 )316.06 
 
 1. 651 gms. 3.1606 gms. or 1000 cc. — V.S. 
 
 = 5H3PO4 + 6H2O + 2K2SO4 + 4MnS04. 
 
 Each cc. of the decinormal V.S. represents 0.001651 gm. 
 of absolute hypophosphorous acid. The quantity of per- 
 manganate solution directed to be added is slightly in excess. 
 The excess is then ascertained by retitration with decinormal 
 oxalic acid. Each cc. of oxalic acid required corresponds to i 
 cc. of decinormal permanganate which has been added in 
 excess of the quantity actually required for the oxidation. 
 
 The excess of permanganate colors the solution red, and 
 the oxalic acid V.S. is then added until the red color just 
 disappears, which indicates that the excess of permanganate 
 is decomposed. 
 
 If 4.7 cc. of decinormal oxalic acid are required, it indicates 
 that 50 cc — 4.7 cc. = 45.3 cc. of decinormal permanganate 
 were actually used up in oxidizing the hypophosphorous acid; 
 therefore 
 
 0.001651 gm.X45.3=o.o747-|- gm. 
 
 of absolute hyposphorous acid, HPH2O2, or 
 
 0.0747 X 100 
 
 — —^ = 24.7 per cent. 
 
166 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 In the above process, boiling facilitates the oxidation, but 
 if the acid is boiled before sufficient permanganate has been 
 added to completely oxidize it, decomposition will take place. 
 Hence direct titration with the permanganate is impossible, 
 and the residual method must be resorted to. 
 
 Calcium Hypophosphite (Ca(PH202)2 = 170.17). o.i gm. 
 of the salt is dissolved in 10 cc. of water, then 10 cc. of sul- 
 phuric acid and 50 cc. of decinormal potassium permanganate 
 are added, and the mixture boiled for fifteen minutes. 
 
 The excess of permanganate is then found by titrating 
 with decinormal oxalic acid solution. 
 
 The reactions which take place are expressed by the fol- 
 lowing equations: 
 
 (i) sCa(PH202)2 + SH2S04= 5CaS04 + 10HPH2O2; 
 
 (2) ioHPH202-l-i2H2S04-F4(2KMn04) 
 
 = ioH3P04 + i2H20-f-4K2S04+8MnvS04. 
 
 These two reactions may be written together thus : 
 
 5Ca(PH202)2 + 17H2SO4 +4(2KMn04) 
 
 = SCaS04 + 4K2S04-l-8MnS04-fioH3P04-t-i2H20. 
 
 6. Methods Involving a Precipitation by Oxalic Acid and the Titra- 
 tion OE THE Excess of the Latter with Standard Permanganate. 
 
 Estimation of Soluble Calcium Salts. To a weighed quan- 
 tity of the calcium salt dissolved in water, a measured excess 
 of normal oxalic acid is added. The mixture is then made 
 alkaline with ammonia and boiled, to facilitate the separa- 
 tion of the precipitate. It is then cooled and diluted with 
 water to an accm-ately measured volume, and after filtra- 
 tion an aliquot portion removed, acidulated with sulphuric 
 
 N 
 acid, and carefully titrated with — potassium permanganate. 
 
 Example. 0.4 gm. of calcium chlorid is dissolved in water, 
 
ANALYSIS BY OXIDATION AND REDUCTION 167 
 
 lo cc. of normal oxalic acid added, the mixture made alkaline 
 with ammonia water, and boiled for a few minutes. It is 
 then filtered, the residue and filter washed with water, and 
 after cooling the combined filtrate and washings are diluted 
 to make loo cc. 
 
 Of this solution 50 cc are taken for analysis (representing 
 0.2 gm. of the salt), acidulated with sulphuric acid, and then 
 
 , ., N . 
 
 titrated with — potassium permanganate to a faint rose tint. 
 
 The 50 cc. of solution represent 5 cc. of normal oxalic 
 
 acid, which is equivalent to 50 cc. of decinormal oxalic add, 
 
 so that whatever number of cc. of decinormal permanganate 
 
 solution is required in the titration, that quantity is to be 
 
 deducted from 50 cc. and the difference multiplied by the 
 
 N 
 
 — factor for calcium chlorid to find the quantity of pure 
 
 CaCl2 present in 0.2 gm. 
 
 N 
 If 14 cc. of — permanganate are employed, then 14 from 
 
 50 cc. leaves 36 cc, the quantity of decinormal oxalic acid 
 solution which combined with the 0.2 gm. of calcium chlorid. 
 Then 
 
 0-00555 gm.X36 = o.i998 gm., 
 
 the quantity of pure CaCb present in the 0.2 gm., or 99.9 
 per cent. 
 
 Estimation of Lead in the Acetate and Subacetate. Take 
 for assay 0.2 gm. of the salt or 2 gms. of the solution in a 
 beaker and add 20 cc. of recently boiled distilled water. 
 Pour this slowly and with constant shaking into a graduated 
 
 N 
 cylinder containing 50 cc. of — oxalic acid V.S. Wash the 
 
 beaker with small portions of distilled water and add the 
 washings to the contents of the cylinder. Then dilute the 
 mixture to 100 cc. and set aside to allow the precipitate to 
 
168 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 settle. Remove 20 cc. of the clear liquid (representing 0.04 
 gm. of the salt or 0.4 gm. of the solution) for titration. Add 
 
 N 
 5 cc. of (1:10) sulphuric acid, and titrate with — perman- 
 ganate until a final drop imparts a permanent pale pink tint. 
 The reactions are represented by the following equations : 
 
 Pb(C2Hs02)2 + H2C2O4.2H2O 
 
 Lead acetate . 
 
 325.1 126.05 
 
 = PbC.Oi + 2HC2H3O2 + 2H2O. 
 
 Pb20(C2H302)2 + 2H2C2O4.2H2O 
 
 Lead subacetate 
 
 548.2 2X126.05 
 
 = 2PbC204 + 2HC2H3O2 + 5H2O. 
 
 N 
 Each cc. of — oxalic acid represents 0.01625 gm. of 
 
 Pb(C2H302)2 or 0.013705 gm. of Pb20(C2H302)2. 
 
 Calcium salts to be estimated by this method must be 
 tolerably pure, and free at least from impurities which would 
 react with oxalic acid or which would reduce permanganate. 
 
 Many of the less soluble calcium salts may be estimated 
 by this method, but they must be subjected to longer treat- 
 ment with the oxalic acid. 
 
 Gold and lead salts may also be estimated by the same 
 method. 
 
 c. Methods Involving a Reduction by Means of Oxalic Acid, and 
 Retiteation of the Excess of the Latter with Potassium Pek- 
 manganate. 
 
 Estimation of Manganese Dioxid (Mn02). The estima- 
 tion of manganese dioxid depends upon the fact that when 
 it is boiled with oxalic acid in the presence of sulphuric acid 
 definite reaction takes place, as the equation shows: 
 
 Mn02+H2C204+H2S04 = MnS04-h2C02 + 2H20. 
 
ANALYSIS BY OXIDATION AND REDUCTION 169 
 
 In the estimation a measured excess of oxalic acid solu- 
 tion is added, together with some sulphuric acid, and the 
 mixture heated until solution is complete. 
 
 The excess of oxaHc acid is then found by retitration with 
 standard permanganate solution. It is well to use a normal 
 oxalic acid solution and a decinormal permanganate solution. 
 
 0.5 gm. of the dioxid is a convenient quantity to operate 
 upon. Each cc. of decinormal solution represents 0.004346 
 gm. of Mn02. 
 
 Example. 0.5 gm. of Mn02 is treated with sulphuric 
 acid and 10 cc. of normal oxalic acid solution, which is equiv- 
 alent to 100 cc. of decinormal oxalic acid solution, the mix- 
 ture treated as described above, and upon retitrating 25 cc. 
 or decinormal permanganate are required. Thus 
 
 IOC cc — 25 cc. = 75 cc. 
 
 N 
 of — oxalic acid went into reaction with the Mn02. Then 
 10 
 
 75X0.004346 = 0.3259 gm. 
 
 d. Methods Involving a Reduction by Means of a Standardized 
 Solution of a Ferrous Salt, and Titration of the Remaining 
 Unoxidized Ferrous Salt, by Permanganate. 
 
 Estimation of Nitrates (Pelouze). This method consists 
 in adding a weighed quantity of the nitrate to an acidulated 
 solution of a ferrous salt of known strength, and, when reaction 
 is complete, estimating the ferrous salt remaining, by titra- 
 tion with permanganate, or in certain cases by means of 
 dichromate V.S. The principle upon which the method is 
 based is, that when nitric acid or a nitrate is brought in con- 
 tact with a highly acidulated solution of a ferrous salt, the 
 former gives off oxygen, which, passing over to the ferrous 
 
170 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 salt, oxidizes it to the ferric state, while at the same time NO 
 is evolved. The reaction is 
 
 2HNO3 + 6HC1 + 6FeCl2 = 3F2CI6 + 4H2O + 2NO. 
 
 Nitric acid Ferrous chlorid 
 
 126.02 760.44 
 
 Iron 
 334.92 
 
 Thus one molecular weight of nitric acid (63.01) will 
 oxidize three molecular weights of ferrous salt, or three atoms 
 of iron (167.46). 
 
 Either hydrochloric or sulphuric acid may be employed. 
 The former is preferred by most operators, and it is generally 
 agreed that in order to attain results of sufficient precision 
 the estimation should be done in the presence of hydrochloric 
 acid only. In using hydrochloric acid, however, where the 
 titrations are to be made with permanganate, certain precau- 
 tions (previously mentioned) must be observed, because of 
 the evolution of chlorin which will otherwise take place and 
 spoil the analysis. This may be obviated by adding to the 
 solution to be titrated an excess of manganese sulphate. 
 
 The NO which is produced during the reaction must be 
 removed by boiling before titration with permanganate is 
 begun. Air must be absolutely excluded during the entire 
 process to prevent oxidation of ferrous salt by the atmospheric 
 oxygen, as well as to prevent oxidation of NO to HNO3, 
 which will oxidize more ferrous salt. The exclusion of air 
 may be partially affected by the use on the flask of a Bimsen 
 valve stopper (see Fig. 41), but the best method is to employ 
 an apparatus so arranged that a constant stream of CO2 or 
 H gas may be passed through it (see Fig. 44). 
 
 This method, although theoretically perfect, is in practice 
 liable to great irregularities, and will give fairly good results 
 only if the directions, especially those as to exclusion of air. 
 
ANALYSIS BY OXIDATION AND REDUCTION 171 
 
 are faithfully carried out. The method of Kjeldahl is to be 
 preferred. 
 
 To conduct the process, weigh accurately 1.5 gms. of flower 
 wire* free from rust (the iron content of which is known), 
 place it an Erlenmeyer flask which is provided with a double 
 perforated stopper fitted with two glass tubes, one of which 
 should reach just to the surface of the liquid in the flask when 
 in place, and the other, which is the outlet tube, should reach 
 
 Fig. 44. 
 
 no lower than the bottom of the stopper. The first of these 
 tubes is connected with an apparatus generating carbon dioxid 
 or hydrogen, while the outlet tube serves to convey the gas 
 into the air or into another flask containing water or an alka- 
 line solution. 30 to 40 cc. of pure fuming hydrochloric acid 
 are added to the iron wire in the flask, gentle heat is applied, 
 and a stream of either CO2 or H passed through the flask and 
 maintained throughout the entire process. When the iron is 
 completely dissolved, the stopper is raised just long enough 
 to introduce a small glass tube open at one end and con- 
 taining the nitrate to be estimated. The quantity or nitrate 
 taken must be equivalent to not more than 0.2 gm. of HNO3. 
 The stopper is then reinserted, heat applied, and gradually 
 increased until the reaction is complete. The free hydro- 
 chloric acid liberates nitric acid from the nitrate and oxidation 
 
 * Or fine piano-forte wire. 
 
172 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 of a portion of the iron is effected. The ferrous chlorid is 
 oxidized to ferric chlorid, as the equation shows, and the 
 solution becomes at first dark brown through the presence 
 of NO. As the heat is increased, the, dark-brown color of 
 the solution is gradually changed to yellow, as ferric chlorid 
 is formed, and increases in intensity until the reaction is 
 complete, then the color remains stationary and indicates com- 
 pletion of oxidation. The solution is now allowed to cool, 
 but the stream of CO2 or H gas is maintained. Forty cc. of 
 a solution of manganese sulphate are now added (this is not 
 necessary if sulphuric acid is used instead of hydrochloric), 
 
 and titration with — potassium permanganate solution begun, 
 
 in order to determine the quantity of unaltered ferrous salt 
 
 remaining in the solution. Assuming that 89 cc. were required, 
 
 the calculation is made as follows: 
 
 Since one molecule of HNO3 (63.01) reacts with three 
 
 atoms of iron (167.46) the quantity of iron found to have 
 
 63.01 
 been oxidized, if multiplied by —z — 7, will give the quantity 
 
 of nitric acid present. 
 
 Example. 1.5 gms. of iron wire, 99.6 per cent Fe= (1.494 gms. 
 
 of iron), is dissolved in hydrochloric acid, as above described, 
 
 and 0.6 gm. of potassium nitrate, KNO3 (loi.ii), added. 
 
 After oxidation, 98 cc. of decinormal permanganate were 
 
 N 
 required. Each cc. of — KMn04 = 0.005582 gm. of Fe. 
 
 0.005582 gm.X 98 = 0.547 gm. of oxidized iron. 1.494 gms. 
 of iron were originally taken. 
 Therefore, 
 
 1.494 
 
 0.547 
 
 gm. = the quantity of iron oxidized. 
 
 0.947 ^ ^ J 
 
ANALYSIS BY OXIDATION AND REDUCTION 173 
 
 Then 
 
 0.047 X6?.oi 
 
 -^%i— °-357 gm. of HNO3, 
 
 which equals 
 
 0.947X63.01X101.11 , rr^rr^ 
 
 167.40X63.01 0/ & -1' 
 
 or 95.5 per cent pure. 
 
 N N 
 
 It is usually advisable to use an — instead of an ^- KMn04 
 , . 5 10 '^ 
 
 solution. 
 
 Chromic Acid and Chromates. Chromic acid oxidizes 
 ferrous salts in the same manner as nitric acid does. The reac- 
 tion is thus expressed: 
 
 6FeS04+6H2S04 + 2Cr03 = Cr2(S04)3+3Fe2(S04)3+6H20. 
 
 SFe =334.92 200 
 
 To an accurately weighed quantity of ferrous ammonium 
 sulphate (Mohr's salt) FeS04 + (NPl4)2S04 + 6H20 (the per- 
 manganate titer of which is known), which is dissolved in a 
 sufficient quantity of diluted sulphuric acid in an Erlenmeyer 
 flask * add a weighed quantity of the chromate or chromic 
 acid in a concentrated aqueous solution. Warm the mixture 
 on a water-bath, under a constant stream of carbon dioxid 
 until the liquid assumes a clear green color. This occurs in 
 a few minutes, and indicates complete reduction of the chromate. 
 
 Now allow the solution to cool, continuing the passage 
 of carbon dioxid through the flask, and transfer the cold 
 solution to a large beaker, and after diluting it to about 300 
 cc. and strongly acidifying it with sulphuric acid, titrate it 
 
 * This flask should be provided with a stopper having two perforations 
 through which glass tubes are passed, one of these, which serves to convey 
 carbon dioxid gas, should reach close to the surface of the liquid, the other 
 tube should end just below the stopper and serve as the outlet tube. See 
 Fig. 44. 
 
174 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 for iinoxidized ferrous salt by means of — potassiiim per- 
 
 lO 
 
 manganate. 
 
 It is usually sufficient to mix the solutions cold, but it 
 is better to employ heat after mixing. A large excess of 
 ferrous salt is unnecessary. It is imperative to dilute the 
 solution highly before titration, as then only can the end color 
 point be accurately determined in the green solution. The 
 use of an excess of sulphuric acid before titration is likewise 
 demanded. A violet-red color marks the end-point, and unless 
 too great a quantity of chromate be taken, or the solution be 
 insufficiently diluted, it can be easily recognized. This method 
 is applicable not only to free chromic acid and soluble chro- 
 mates, but also to chromates which are insoluble in water.* 
 It can therefore be employed for the indirect estimation of 
 such bases as are precipitable by chromic acid, out of neutral, 
 ammoniacal or acetic acid solutions, as for instance lead, 
 bismuth and barium. 
 
 Finally, the method may be employed for the estimation 
 of chromic oxids. The solution of the latter is treated with 
 an excess of sodium carbonate, bromine water added, and 
 heat applied until a clear solution results. This solution, 
 which contains all of the chromium in the form of sodium 
 chromate, is evaporated, the residue dissolved in dilute acetic 
 acid and the chromium completely precipitated by means of 
 lead acetate. The precipitated lead chromate is then treated 
 as above. 
 
 The calculation is made with reference to the equation, 
 in which it is shown that one molecule, loo of chromic oxid 
 (CrOs), is equivalent to three molecules (167.46) of metallic 
 
 * In the case of insoluble chromates the salt is shaken directly with the 
 ferrous solution, and the mixture more highly diluted, and more strongly 
 heated, than in the case of soluble salts. 
 
ANALYSIS BY OXIDATION AND REDUCTION 175 
 
 lOO 
 
 iron. The quantity of iron oxidized, multiplied by 
 
 167.46' 
 
 gives the weight of chromic oxid present, and from this its 
 equivalent in potassium, sodium, lead, bismuth or barium 
 chromate is calculated. 
 
 In the case of potassium dichromate (K2Cr207) one mole- 
 cule (294.2) is equivalent to six atoms (334.92) of metallic 
 
 294.2 
 iron. The quantity of iron oxidized is multiplied by . 
 
 Example. To 1.5 gms. of ammonio-ferrous sulphate (con- 
 taining 0.2142 gm. Fe) add 0.1241 gm. of K2Cr207 (molecular 
 weight 294.2), and after complete oxidation, titrate the solu- 
 
 N 
 tion with — KMn04 to determine the quantity of unchanged 
 
 ferrous salt. 13 cc. are required. Each cc. represents 0.005582 
 gm. of Fe. 
 
 Thus, 13X0.005582 gm. =0.0725 gm., the quantity of iron 
 which was not oxidized by the dichromate. This, deducted 
 from the quantity of iron originally added (0.2142 — 0.0725 
 = 0.1417 gm.), gives the quantity which was oxidized. 
 
 Then, 
 
 O.I4I7XIOO „ r ^ ^ 
 
 — ^ '. — = 0.08475 -f gm. of CrOs 
 
 or 
 
 0.1417X 294.2 
 
 334-92 
 
 = 0.1247 gm. of K2Cr207. 
 
 Example. To 1.5 gms. of ammonio-ferrous sulphate (con- 
 taining 0.2142 gm. of Fe) add the precipitate of barium 
 chromate obtained from 0.2491 gm. of BaCl2-l-2H20 (molec- 
 ular weight 244.29) and after complete oxidation, titrate with 
 
 N 
 
 — permanganate. 7.8 cc. are consumed, thus 7.8X0.005582 
 
 =0.04353 gm. the quantity of unoxidized iron present. Then 
 
176 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 0.2142 — 0.04353 = 0.17057 gm. of iron oxidized by the barium 
 chromate. 
 
 0.17057X244.20 
 
 — '-— — ^^^*— ^ = 0.2489 gm. BaCl2 + 2H20. 
 
 e. Methods Involving the Oxidation of the Substance Analyzed by 
 Means of a Ferric Salt, and Titration or the Resultant Fer- 
 rous Salt. 
 
 Estimation of Tin (Lowenthal). When stannous chlorid 
 is brought in contact with ferric chlorid it acts as a reducing 
 agent, the ferric chlorid being reduced to ferrous, and the 
 stannous chlorid oxidized to stannic. This reaction, which 
 is an exact quantitative one, takes place according to the 
 following equation: 
 
 SnCla + FeaCle = SnCU + 2FeClo. 
 2)119 2)111.64 
 
 59-5 55-82 
 
 Every 55.82 parts of iron reduced to the ferrous state 
 
 represents 59.5 parts of tin. The quantity of ferrous iron 
 
 N 
 produced is determined by titration with — permanganate. 
 
 Metallic tin, or any protosalt of tin, will dissolve in ferric 
 chlorid solution in the presence of a little hydrochloric acid 
 and act in the manner described. 
 
 About 0.5 gm. of tin or an equivalent of stannous salt 
 is introduced into a flask graduated at 250 cc. Five cc. of 
 tolerably concentrated ferric chlorid solution and a little hydro- 
 chloric acid are then added. It is well to drop into the flask 
 a crystal of sodium carbonate, in order to produce carbon 
 dioxid to replace the air in the flask and thus prevent oxida- 
 tion by the oxygen of the air. The mixture is gently warmed 
 until the tin is wholly dissolved, and then the solution diluted 
 with cold, recently boiled water, to 250 cc. Of this solution 
 50 cc. is withdrawn by means of a pipette and titrated with 
 
ANALYSIS BY OXIDATION AND REDUCTION 177 
 
 N N 
 
 — permanganate. Each cc. of — permanganate = 0.005582 
 
 gm. Fe = 0.00595 gm. Sn. 
 
 It is always advisable, when an exact assay is to be made, 
 to make a blank experiment, using a like quantity of water 
 and ferric chlorid solution, and deducting the quantity of 
 permanganate solution used from the quantity required in 
 the assay. The difference is calculated as tin. 
 
 Estimation of Qopper (Fleitmann). The copper salt is 
 first reduced either to cuprous oxid by means of glucose or 
 to the metallic state by means of pure zinc. The reduced 
 product is then dissolved in a mixture of ferric chlorid and 
 hydrochloric acid, a little sodium carbonate added to expel 
 
 N 
 the air, and the titration with — permanganate begun, as in 
 
 the preceding assay. The reaction is 
 
 or 
 
 CU2O + FeaCle + 2HCI = 2CUCI2 + H2O + 2FeCl2 
 Cu + FeaCls = CuCb + iFeClj. 
 
 2)63.57 aim. 64 
 31-78 55-82 
 
 N 
 Each cc. of — KMn04 = o.oo3i78 gm. Cu. 
 
 Example, o.i gm. of the copper salt is dissolved in water, 
 an excess of a weak sodium carbonate solution is then added, 
 the mixture heated to boiling and a slight excess of glucose 
 solution added, i.e., sufficient to completely precipitate the 
 copper as CU2O and cause the blue color of the solution to 
 disappear. The boiling should be continued several minutes 
 longer and the precipitate separated by filtration through asbes- 
 tos. After thoroughly washing the precipitate with hot water, 
 it is transferred to a flask filled with CO2. Solution of 
 ferric chlorid or ferric sulphate is then added, together with 
 
178 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 some hydrochloric or sulphuric acid, and when the precipitate 
 
 N 
 (CuaO) is dissolved, the solution is titrated with — perman- 
 ganate V.S. 
 
 Volumetric Analysis by Means of Potassium Dichromate 
 
 In some respects the dichromate possesses advantages over 
 permanganate: 
 
 1. It may be obtained in a pure state. . 
 
 2. Its solution does not deteriorate upon standing as does 
 that of permanganate. 
 
 3. It is not decomposed by contact with rubber as the 
 permanganate is, and may therefore be used in Mohr's burette. 
 Its great disadvantage, however, is that when used in the 
 estimation of ferrous salts the end-reaction can only be found 
 by using an external indicator. The indicator which must 
 be used is freshly prepared potassium ferricyanid T.S., a 
 drop of which is brought in contact with a drop of the solu- 
 tion being tested, on a white slab, at intervals during the 
 titration, the end of the reaction being the cessation of the 
 production of the blue color, when the two liquids are brought 
 together. Thus the estimation by potassium dichromate is 
 cumbersome, and very exact results are not as easily obtained 
 as with permanganate. 
 
 Besides ferrous salts, a great many other substances may 
 be estimated by oxidation analysis with dichromate. Among 
 them nitrat-es, sulphates, arsenous acid, barium, lead, ferric 
 salts after reduction by stannous chlorid or an alkaline sul- 
 phite, but not after reduction by means of metallic zinc. 
 The presence of the dissolved zinc salt interferes with the 
 reaction of the ferricyanid indicator. Ferrous salts may be 
 estimated in the presence of hydrochloric acid, by means of 
 dichromate, without the precautions that apply in the case 
 
ANALYSTS BY OXIDATION AND REDUCTION 179 
 
 of permanganate. Chromium as chromate may be indirectly 
 
 estimated; an excess of a solution of a ferrous salt being 
 
 added and then the excess determined by dichromate. lodids, 
 
 thiosulphates and alkalies may also be estimated by means 
 
 of potassium dichromate. 
 
 N 
 Preparation of Decinormal — Potassium Dichromate 
 
 10 
 
 N 
 (K2Cr207 = 294.2; — V.S. =4.903 gms. in 1000 cc). 
 
 4.903 gms. of pure potassium dichromate * which has 
 been pulverized and dried at 120° C. is dissolved in sufK- 
 cient water to make 1000 cc. of solution. 
 
 It will be noticed that Yo of the molecular weight of the 
 dichromate (expressed in grams) is taken in the preparation 
 of 1000 cc. of this solution. The reason for this is that one 
 molecule of potassium dichromate when treated with an acid 
 yields three atoms of nascent oxygen which are available for 
 oxidizing purposes, thus 
 
 K2Cr207+4H2S04 = K2S04 + Cr2(S04)3+4H20 + 03; 
 
 and since each atom of oxygen is equivalent to two atoms 
 of hydrogen, one molecule of the dichromate must be equiv- 
 alent to six atoms of hydrogen. Hence a normal solution of 
 potassium dichromate, when used as an oxidizing agent, should 
 contain one-sixth of the molecular weight, expressed in grams, 
 in 1000 cc. (see definition for normal solution) and its deci 
 normal solution ^. 
 
 If a standard solution of potassium dichromate is to be 
 made for use as precipitant, as in the titration of barium, 
 one-fourth of the molecular weight is to be taken for 1000 
 cc. of the normal solution, as explained in Chapter III. 
 
 * Potassium dichromate for use in volumetric analysis should respond 
 to all the tests for purity given in the U. S. P., or it should be recrystallized 
 several times and then dried. 
 
180 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Standard solution of potassium dichromate is sometimes 
 used as a neutralizing solution for estimating alkalies, phenol- 
 phthalein being used as indicator. 
 
 When used for this purpose the normal solution contains 
 147. 1 gms. in I liter (one-half the molecular weight in grams). 
 It is then the exact equivalent of any normal acid V.S. 
 
 2KOH + KaCr^Or = 2K2Cr04 + H2O. 
 
 2)112.2 2)294.2 
 
 56.1 gms. 147. 1 gms., or looo cc. normal V.S. 
 
 Decinormal potassium dichromate may also be used in 
 conjunction with potassium iodid and sulphuric acid for 
 standardizing sodium thiosulphate. lodin is liberated from 
 potassium iodid in this reaction. The reaction is expressed 
 by the equation 
 
 KaCraO 7 + 6KI + 7H2SO4 = 4K2SO4 + Crg ( SO4 ) 3 + 7H2O + 3I2. 
 
 Thus one molecule of the dichromate will liberate six 
 atoms of iodin, therefore a normal solution should contain 
 one-sixth of the molecular weight, and a decinormal. solution 
 ^V in Tooo cc. The solution is hence of the same strength 
 as that which is used for oxidizing purposes. If the deci- 
 normal solution containing 14.71 gms. in i liter is used, it 
 
 3N 
 has the effect of a — solution. 
 10 
 
 The decinormal solution which is used as an oxidizing 
 agent is chemically equivalent to decinormal potassium per- 
 manganate. When used for the purpose of liberating iodin 
 from potassium iodid, it is the equivalent of an equal volume 
 of decinormal sodium thiosulphate. 
 
 For litraling ferrous salts the decinormal solution of dichro- 
 mate is used in the following manner: 
 
 Make an aqueous solution of the ferrous salt, introduce 
 
ANALYSIS BY OXIDATION AND REDUCTION 181 
 
 it 'into a flask, and acidulate it with sulphuric or hydrochloric 
 acid. Now add gradually from a burette the decinormal 
 potassium dichromate until a drop taken out upon a white 
 slab no longer shows a blue color with a drop of freshly 
 prepared potassium ferricyanid T.S. Note the number of 
 cc. of the standard solution used, multiply this number by 
 the factor, and thus obtain the quantity of pure salt in the 
 sample taken. 
 
 Ferrous salts strike a blue color with potassium ferricyanid, 
 but as the quantity of ferrous salt gradually diminishes during 
 the titration, the blue becomes somewhat turbid, acquiring 
 first a green, then a gray, and lastly, a brown shade. The 
 process is finished when the greenish-blue tint has entirely 
 disappeared. 
 
 The reaction of potassium dichromate with ferrous salts 
 always takes place in the presence of free sulphuric or hydro- 
 chloric acid at ordinary temperatures. Nitric acid should not 
 be used. 
 
 If it is desired to estimate ferric salts by this standard 
 solution it is necessary to first reduce them. 
 
 This may be done by metallic magnesium, sulphurous 
 acid, the alkali sulphites, or by stannous chlorid. 
 
 Standard potassium dichromate may be checked in the 
 same way as standard permanganate, with pure metallic iron. 
 
 ESTIMATION OF FERROUS SALTS WITH POTASSIUM DICHROMATE 
 
 One molecule of potassium dichromate yields, under favor- 
 able circumstances, three atoms of oxygen. This is shown by 
 the following equation: 
 
 KaCraOy = CrgOs + K2O + O3. 
 Here it is seen that the three liberated atoms of oxygen 
 
182 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 combine at once with the ferrous oxid, converting it into 
 ferric oxid: 
 
 6FeO + 03 = Fe609 or sFcaOs. 
 
 In the oxidation of a ferrous salt, the reaction takes place 
 only in the presence of an acid. 
 
 The dichromate then gives up its oxygen. Four of its 
 oxygen atoms combine at once with the replaceable hydrogen 
 of the accompanying acid, the other three being liberated. 
 The three oxygen atoms thus set free are available either for 
 direct oxidation or for combination with the hydrogen of 
 more acid. In the latter case a corresponding quantity of 
 acidulous radicals is set free. 
 
 K2Cr207+4H2S04 = K2S04 + Cro(S04)3+4H20 + 03. 
 
 In this case four of the liberated atoms of oxygen combine 
 with eight of the atoms of hydrogen of sulphuric acid and 
 liberate four SO4 radicals, which at once combine with the 
 K2 and Cr2 of the dichromate. The other three atoms are 
 set free. If seven sulphuric acid molecules are used instead 
 of four molecules, the three free atoms of oxygen will liberate 
 3(S04): 
 
 K2Cr207 + 7H2S04 = K2S04 + Cr2(S04)3 + 7H20 + (804)3. 
 
 If this liberation of 3(804) takes place in the presence of, 
 a ferrous salt, the 3(804) will combine with six molecules of 
 the ferrous salt, converting it into a ferric salt: 
 
 6Fe804+3S04=Fe6(S04)9 = 3Fe2(S04)3; 
 6FeS04 + K2Cr207 + 7H2SO4 
 
 = K2S04 + Cr2(804)3 + 7H20 + [3Fe2(S04)3]. 
 If in the above case hydrochloric acid is used instead of 
 
ANALYSIS BY OXIDATION AND REDUCTION 183 
 
 sulphuric, fourteen molecules of the former must be taken to 
 supply the neccessary hydrogen. 
 
 The seven liberated atoms of oxygen must have fourteen 
 atoms of hydrogen to combine with. 
 
 Three of these atoms of oxygen liberate six univalent or 
 three bivalent acidulous radicals. 
 
 Therefore, since one molecule of K2Cr207 will give up 
 for oxidizing purposes three atoms of oxygen, which are equiva- 
 lent chemically to six atoms of hydrogen, one-sixth of the 
 molecular weight in grams of the dichromate, dissolved in 
 sufl&cient water to make one liter, constitutes a normal solution, 
 and one-tenth of this quantity of K2Cr207 in a liter, a deci- 
 normal solution. 
 
 Thus the estimation of ferrous salts is effected by oxidizing 
 them to ferric with an oxidizing agent of known power, the 
 strength of the ferrous salt being determined by the quantity 
 of _ the oxidizing agent required to convert it to ferric. 
 
 Saccharated Ferrous Carbonate (FeC03= 115.82). One 
 gm. of saccharated ferrous carbonate is dissolved in 10 cc. of 
 diluted sulphuric acid and the solution diluted with water 
 to about 100 cc. The decinormal potassium dichromate is 
 carefully added, until a drop of the solution taken out and 
 brought in contact with a drop of freshly prepared solution 
 of potassium ferricyanid ceases to give a blue color. 
 
 The number of cc. of the dichromate solution is read off 
 and the following equations applied : 
 
 6FeC03-f6H2S04=6FeS04 + 6H20 + 6C02; 
 694.9 911-34 
 
 then 
 6FeC03 or 6FeS04 + K2Cr207 + 7H2S04 
 
 6 )694.9 6 )911. .S4 6) 294.2 
 
 1 9)115.82 10 )151.89 10 )49.03 j^ 
 
 11.582 gms. 15.189 gms. 4.903 gms., or 1000 cc. — KjCrjO, V.S. 
 
 = K2S04 + Cr2(S04)3 + 7H20+3Fe2(S04)3. 
 
181 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 Thus each cc. of — K2Cr207 represents 0.011582 gm. of 
 
 pure ferrous carbonate or 0.005582 gm. of metalhc iron. 
 
 If strong sulphuric acid is added to saccharated ferrous 
 carbonate it will char the sugar, and a black mass of burnt 
 sugar is obtained. This may be prevented by adding water 
 first, and then, slowly, the sulphuric acid. 
 
 Instead of sulphuric acid, hydrochloric acid may be used. 
 
 Fig. 45- 
 
 This will not char the sugar, but the ferrous chlorid which 
 is then formed is too readily oxidized by the air. 
 
 It has also been suggested that as hydrochloric acid so 
 rapidly converts ordinarj' sugar into invert sugar as to render 
 it easily attacked by the dichromate, it should be cautiously 
 used, if at all. Phosphoric acid has none of these disad- 
 vantages, and may be employed with good results. 
 
 In making estimations of ferrous salts with potassium 
 
ANALYSIS BY OXIDATION AND REDUCTION 18.3 
 
 dichromate, care should be taken to avoid atmospheric oxida- 
 tion. It is good practice to calculate approximately how 
 much of the standard solution will probably be required to 
 complete the oxidation, and then add almost enough of the 
 standard solution at once, instead of adding it slowly. 
 
 A white porcelain slab is then got ready, and placed 
 alongside of the flask in which the titration is to be performed. 
 Upon this slab are placed a number of drops of the freshly 
 prepared solution of potassium ferricyanid, and at intervals 
 during the titration a drop is taken from the flask on a glass 
 rod and brought in contact with one of the drops on the 
 slab. The glass rod should always be dipped in clean water 
 after having been brought in contact with a drop of the 
 indicator. See Fig. 45. 
 
 When a drop of the solution ceases to give a blue color 
 on contact with the indicator, the reaction is complete. 
 
 Ferrous Sulphate (FeS04 + 7H2O = 277.89). Dissolve about 
 one gram of crystallized ferrous sulphate in a little water, 
 add a good excess of sulphuric or hydrochloric acid, titrate 
 with the decinormal potassium dichromate, as directed under 
 Ferrous Carbonate, and apply the following equation: 
 
 6(FeS04.7H20) + KaCrgOr + 7H2SO4 
 6)1667.34 
 10 )277.89 j^ 
 
 27.789 gins., or 1000 cc. — K^CruOy V.S. 
 
 = 3Fe2(S04)3 + K2SO4 + Cri(S04)3 + 49H2O. 
 
 N 
 Thus each cc. of the — K2Cr207 V.S. represents 0.027789 
 
 gm. of crystallized ferrous sulphate or 0.015 189 anhydrous. 
 If one gm. of the salt is taken and dissolved as above, it 
 should require about 37 cc. of the standard solution, equiva- 
 lent to about 100 per cent. 
 
1S6 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Anhydrous Ferrous Sulphate 
 6FeS04 + KoCraOr + 7H2SO4 
 
 6 )9"-.U 
 10)15 1-89 ^T 
 
 15.189 gms., or 1000 cc. — KjCrjO, V.S. 
 
 = 3Fe2(s6°4)3 + K2SO4 + Cr,(S04)3 + 7H2O. 
 
 Each cc. of the standard solution represents 0.015189 gm. 
 of real ferrous sulphate or 0.005582 gm. of metallic iron. 
 
 Dried (Exsiccated) Ferrous Sulphate has the approximate 
 composition FeS04 + 3H20. 
 
 It is tested in the same manner as the anhydrous ferrous 
 sulphate. 
 
 TABLE OF SUBSTANCES WHICH MAY BE ESTIMATED BY 
 MEANS OF POTASSIUM PERMANGANATE OR POTASSIUM 
 BICHROMATE. 
 
 Name. 
 
 Fori 
 
 Molecular 
 Weight. 
 
 N 
 
 Factor. 
 
 Acid, chromic 
 
 " hypophosphorous 
 
 " nitric 
 
 " nitrous 
 
 " oxalic (cristallized) 
 
 Barium dioxid 
 
 Calcium chlorid 
 
 " hypophosphitc 
 
 Ferric chlorid 
 
 " hypophosphitc 
 
 " sulphate 
 
 Ferrous carbonate 
 
 " oxid 
 
 " sulphate (anhydrous). 
 " " (crystallized) 
 
 Ferrum (metallic) 
 
 Hydrogen dioxid 
 
 Manganese dioxid 
 
 Potassium hypophosphitc . . . . 
 
 Sodium hypophosphite 
 
 " nitrite 
 
 CrOj 
 
 HPH2O2 
 
 HNO3 
 
 HNOj 
 
 H2C2O4+2H2O 
 
 BaOj 
 
 CaClj 
 
 CaCPH^Oj)^ 
 
 FeCIs 
 
 Fe(PH202)3 
 
 Fe2(SO,)3 
 
 FeCOs 
 
 FeO 
 
 FeSO< 
 
 FeSO, + 7H20 
 
 Fe2 
 
 H2O2 
 
 Mn02 
 
 KPHzOj 
 
 NaPHjOj 
 
 NaNO, 
 
 100 
 66.04 
 
 63-01 
 
 47.01 
 
 126.05 
 
 169-37 
 III. 01 
 1 70. 1 7 - 
 162 . 20 
 250.94 
 
 399-85 
 115.82 
 71.82 
 151.89 
 277.89 
 
 55-82 
 34-0 
 86,93 
 104.15 
 88.05 
 69.01 
 
 0.0033 
 0.00165 I 
 0.0021 
 
 u. 00235 
 
 o . 0063 
 o . 00845 
 
 0-00555 
 
 0.002127 
 
 0.01622 
 
 0.025094 
 
 0.01999 
 
 O.OII582 
 
 0.007182 
 
 0.015 189 
 
 0.027789 
 
 0.005582 
 
 0.0017 
 
 0.004346 
 
 0.0026 
 
 0.0022 
 
 0.00345 
 
ANALYSIS BY OXIDATION AND REDUCTION 187 
 
 Granulated Ferrous Sulphate (FeS04 + 7H20) is tested in 
 the same manner as crystallized ferrous sulphate, with which 
 it should correspond in strength. 
 
 Analysis by Indirect Oxidation 
 
 This method of analysis is based upon the oxidizing power 
 of iodin. 
 
 lodin acts upon the elements of water, forming hydri- 
 odic acid with the hydrogen, and liberating oxygen in a nascent 
 state. 
 
 Nascent oxygen is a very active agent, and readily com- 
 bines with and oxidizes many substances, such as arsenous 
 oxid, sulphurous acid, sulphites, thiosulphates, hydrosulphuric 
 acid, the lower oxids of antimony and their salts. 
 
 As203 + 2H20-|-2l2 = As205+4HI; 
 
 H2SO;h+H20+I2 = 2HI+H2S04. 
 
 Therefore iodin is said to be an indirect oxidizer, and 
 may be used for the estimation of a great variety of substances 
 with extreme accuracy. 
 
 When iodin is brought in contact with certain oxidizable 
 substances it is decolorized. This decolorization occurs as 
 long as some of the oxidizable substance is present, and ceases 
 when oxidation is complete. Hence when the yellow color 
 of iodin shows itself in the solution being analyzed the reaction 
 is known to be at an end. In most cases a more delicate 
 end-reaction is obtained by using starch solution as an indicator. 
 This gives a distinct and unmistakable blue color with the 
 slightest excess of iodin. 
 
 In making an analysis with standard iodin solution, the_ 
 substance imder examination is brought into dilute solution 
 (usually alkaline), the starch solution added, and then the 
 
188 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 iodin, in the form of a standard solution, is delivered in from 
 
 a burette, stirring or shaking constantly, until a final drop 
 
 colors the solution blue — a sign that a slight excess of iodin 
 
 has been added. 
 
 N 
 Preparation of Decinormal Iodin (1=126.92; — V.S. = 
 
 12.692 gms. per liter). Dissolve 12.692 gms. of pure* iodin 
 in 300 cc. of distilled water containing 18 gms. of pure potas- 
 sium iodid. Then add enough water to make the solution 
 measure at 15° C. (59° F.) exactly 1000 cc. 
 
 The solution should be kept in small glass-stoppered vials, 
 in a dark place. 
 
 The potassium iodid used in this solution acts merely as 
 a solvent for the iodin. 
 
 If pure iodin is used in making this solution, there is no 
 necessity for checking (standardizing) it. 
 
 But if desired the solution may be checked against pure 
 arsenous acid or sodium thiosulphate. It there is any doubt 
 as to the purity of the iodin, it is best to take a larger quantity, 
 say 14 gms. instead of the 12.692 gms. directed above, and 
 then dilute the resulting solution to the proper strength after 
 standardizing. 
 
 * If pure iodin be not at hand, it may be prepared from the commercial 
 article as follows: 
 
 Powder the iodin and heat it in a porcelain dish placed over a water- 
 bath, stirring constantly with a glass rod for 20 minutes. Any adhering 
 moisture, together with any cyanogen iodid, and most of the iodin bromid 
 and iodin chlorid, is thus vaporized. 
 
 Then triturate the iodin with abput five per cent of its weight of pure, dry 
 potassium iodid. The iodin bromid and chlorid are thereby decomposed, 
 potassium bromid and chlorid being formed and iodin liberated from the 
 potassium iodid. 
 
 The mixture is then returned to the porcelain dish, covered with a clean 
 glass funnel, and heated on a sand-bath. A pure resublimed iodin is then 
 obtained. 
 
ANALYSIS BY OXIDATION AND REDUCTION 189 
 
 Standardization of lodin V.S. by Means of a Decinor- 
 mal Sodiiiiii TJiiosiilpJiate Solution. 25 cc. of the iodin solu- 
 tion are accurately measured off into a beaker, and then from a 
 
 N 
 burette the — thiosulphate is delivered until the solution is of 
 
 a pale yellow color, two or three drops of starch solution are 
 then added, and the titration with the thiosulphate solution 
 continued until the blue color of starch iodid is discharged. 
 
 If the iodin solution is exactly decinormal, the 25 cc. will 
 require 25 cc. of decinormal sodium thiosulphate to exactly 
 complete the reaction. 
 
 If on the other hand more than 25 cc. of thiosulphate 
 solution is required, it indicates that the iodin solution is too 
 concentrated, and must be diluted so as to correspond with 
 the thiosulphate solution, volume for volume. 
 
 Example. Assuming that in the above titration 27 cc. of 
 the thiosulphate solution were used, then each 25 cc. of the 
 iodin solution must be diluted with water to make 27 cc. in 
 order to convert the iodin solution into a strictly decinormal 
 solution. If, however, the iodin solution is found to be weaker, 
 
 N 
 as evidenced bv its using up less than its own volume of — 
 
 " ^ 10 
 
 thiosulphate, its relative strength should be noted on the label 
 of the container. 
 
 Thus if only 24.8 cc. of the thiosulphate solution are 
 
 25 
 used up, then i cc. of the latter equals — ^ cc. or 1.008 cc. 
 
 of the iodin solution. 
 
 One cc. of this iodin solution is equivalent to 0.992 cc. of 
 
 N 
 
 — thiosulphate, which is the same as saymg i cc. =0.992 
 
 ^° N 
 
 cc. of — iodin, or expressed in another way, i cc. of this iodin 
 10 
 
 solution contains 0.01259+ gm. of iodin. 
 
 Such an iodin solution may be used as an empirical solu- 
 
I!i0 THK ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 tion. and in any assay the quantity of it (in cubic centimeters) 
 
 24.8 
 which is consumed is divided by 1.008 or multiplied by 
 
 or by 0.992, and then multiplied by the decinormal 'factor 
 
 for the substance analyzed. Another way is to multiply the 
 
 cc. of this iodin solution used by the weight of iodin contained 
 
 in each cc, and then by a fraction in which the numerator 
 
 represents the quantity of the substance analyzed equal to 
 
 an atom of iodin, and the denominator is the atomic weight 
 
 of iodin. 
 
 Example, o.i gm. of arsenous acid consumes 20 cc. of 
 
 this empirical solution. How much absolute AS2O3 does it 
 
 N 
 contain ? The — factor for AS2O3 is 0.004948 gm. 
 
 -.M- , , / S 20X0.004948 
 
 Method (a) 3 =0.0981 gm. 
 
 24.8 
 Method (b) 20 X X 0.004948 = 0.0981 gm. 
 
 Method (c) 20X0.992X0.004948 = 0.0981 gm. 
 
 Method (d) 2oXo.oi2c:9X — 7 — =0.0981 gm. 
 ^ -^^ 126.92 ^ * 
 
 It is a good plan to have the factors marked on the labels. 
 In the above case the label may, be marked 
 
 24.8 
 X or X0.992 or ICC =0.01259 g^i- iodin. 
 
 Standardizntion of Iodin V.S. by Means of Arsenous 
 Oxid. 0.2 gm. of pure resublimed vitreous arsenous oxid is 
 weighed off very carefully into a flask 50 cc. of water are 
 added, and then, after the addition of 2 gms. or more of 
 
ANALYSIS BY OXIDATION AND REDUCTION 191 
 
 sodium bicarbonate, the mixture is gently warmed and shaken 
 until the arsenous oxid is completely dissolved.* 
 
 To this solution a few drops of starch indicator are added, 
 and then the iodin solution delivered carefully from a burette 
 until a blue color marks the end of the reaction. 
 
 AS2O3 + I4 + 2H2O = AS2O5 + 4HI. 
 197.92 4x126.92 
 
 49.48 gms. of Ar203 = 126.92 gms. of iodin; 
 
 4.948 " " As203= 12.692" " " or 1000 cc. — V.S. 
 
 0.2 gm. of AS2O3 will require 
 
 1000X0.2 N . 
 
 ^ — = 40.44 cc. of a true — 10dm V.S. 
 
 4.948 10 
 
 Assuming that in the above titration 37.4 cc. of the iodin 
 solution were used, then the iodin solution is too concentrated 
 and must be diluted so that each 37.4 cc. will be made up to 
 40.44 cc. 
 
 After diluting in this way a new trial should be made. 
 
 It is a good plan to make a decinormal solution of the 
 arsenous oxid by dissolving 4.948 gms. of the pure oxid and 
 30 gms. of sodium bicarbonate in sufficient water to make 
 1000 cc. at 15° C. and to titrate this with the iodin solution. 
 25 cc. of this solution should require for complete oxidation 
 exactly 25 cc. of the iodin solution, if the latter is strictly of 
 decinormal strength. 
 
 * Arsenous oxid is much more readily soluble in alkali hydroxid, than in 
 carbonated alkalies, therefore the following method of making the sulution 
 is preferred: 0.2 gm. of arsenous oxid is dissolved in a small quantity of 
 boiling water with the aid of potassium hydroxid (free from sulphur), the 
 solution is then acidified with hydrochloric acid, and then again made alkaline 
 by the addition of sodium bicarbonate. The latter must be added in con- 
 siderable excess, being careful, however, to avoid loss of solution during 
 effervescence. 
 
1<)2 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The Starch Solution. This solution, which is used as an 
 indicator in iodometric determinations, is made as follows: 
 
 One gm. of starch (potato, arrowroot or corn starch), is tri- 
 turated with lo cc. of cold water, until a smooth mixture is 
 obtained, then sufficient boiling water is added, with constant 
 stirring, to make 200 cc. of a thin, translucent fluid. If the 
 solution is not translucent it should be boiled for about three 
 minutes, then allowed to cool, and filtered. This solution 
 does not keep very long, in fact it becomes useless after 
 standing one day, therefore it should be freshly prepared 
 when required. 
 
 This indicator is very sensitive to iodin — it will detect 
 one part of iodin in 3,500,000. If the solution is not clear, 
 or contains flocks of insoluble starch, the characteristic 
 beautiful blue color is not obtained with iodin; instead, a 
 greenish or brownish color is produced, and the insoluble 
 particles are even colored black and are decolorized with 
 difficulty. 
 
 The blue color which starch gives with iodin constitutes 
 a very delicate indication of the slightest excess of iodin. 
 This color is usually regarded as being due to the formation 
 of a compound of starch and iodin, called iodid of starch. 
 It is a compound of very unstable character and of doubtful 
 composition. 
 
 Sodium thiosulphate behaves towards iodid of starch 
 exactly as it does toward free iodin — it takes up the iodin 
 and thus discharges the blue color. 
 
 Iodid of starch dissociates upon heating, but reunites 
 upon cooling, hence it is advisable to avoid heat in estimations 
 where starch is used as an indicator. 
 
 In order to prevent the deterioration of this solution a 
 few drops of chloroform may be added; thi$ will preserve 
 it for a long time. Oil of cassia is also recommended as a 
 
ANALYSIS BY OXIDATION AND REDUCTION 193 
 
 preservative. Moerk adds 2 cc. of the oil to a liter of the 
 cooled starch solution. Zinc chlorid or iodid added to the 
 boiling starch solution will prevent its decomposition for a 
 long time. A starch solution so made, however, should not 
 be used in titrations of sulphids, because zinc reacts with 
 sulphids. 
 
 In the case of solutions containing carbonates, the pre- 
 cipitate of zinc carbonate is so small in amount that it does 
 not interfere in the least with the recognition of the end- 
 reaction tint. Mercuric iodid is also a very valuable preserva- 
 tive. 
 
 o.oi gm. of mercuric iodid in a liter of the starch solution 
 is quite sufficient. A very satisfactory indicator is the com- 
 mercial soluble starch which is made by heating potato starch 
 with glycerin and precipitating the starch by repeated treat- 
 ment with alcohol. This starch dissolves readily in hot water 
 forming a clear solution, which gives a very delicate reaction 
 with iodin. It is best preserved under alcohol, the latter 
 being removed by filtration and evaporation, when the starch 
 is wanted for making a solution. 
 
 In making starch solution for use as an indicator, long 
 continued boiling should be avoided, as this converts some 
 of the starch into dextrin. 
 
 On the Use of Sodium Bicarbonate in Titrations with 
 Iodin. In these titrations an excess of alkali is necessary in 
 order to neutralize the hydriodic acid formed. 
 
 AS203 + 2HoO-|-2l2 = As205-t-4HI. 
 
 If the hydriodic acid is not removed by neutralization it 
 will react with the arsenic oxid (AS2O 5), reducing it to arsenous 
 oxid (AS2O3) and liberate iodine, as shown by the following 
 equation : 
 
 4HI -I- AS2O5 == AS2O3 -f- 2H2O -I- 2I2. 
 
194 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Sodium bicarbonate is usually employed to neutralize the 
 HI and should be used in slight excess. 
 
 Alkali hydroxids or carbonates cannot be used for this 
 purpose, because they react with free iodin or even with 
 starch iodid. Bicarbonates ordinarily have no such action, 
 and therefore sodium bicarbonate is usually directed to be 
 added in excess to the solution to be titrated with iodin. 
 
 It is well known that sodium hydroxid solution reacts 
 with free iodin, with formation of hypoiodite and iodid. 
 
 2NaOH + l2 = NaIO+NaI+H20, 
 
 the hypoiodite quickly forming iodate. 
 
 3NaIO = 2NaI + NaI03. 
 
 It is also now a recognized fact that sodium carbonate 
 is partly hydrolyzed when in solution, with formation of 
 some sodium hydroxid, as per equation, 
 
 NaoCOs + H2O = NaOH + NaHCOs. 
 
 It therefore reacts in much the same way with iodin as 
 the hydroxid, though to a less extent. 
 
 On the other hand, it is generally supposed that bicar- 
 bonate of soda is without effect on iodin, and when, in iodo- 
 metric estimations, addition of sodium bicarbonate is indicated, 
 little attention is given to amount added, as long as it be in 
 excess. 
 
 The experiments of W. A. Puckner, Proc. A. Ph. A., 1904, 
 408, prove that we are entirely wrong in the supposition that 
 sodium bicarbonate has no effect upon iodin. He showed 
 that when using i to 2 gms. of the bicarbonate, an error of 
 1.5 to 4.5 cc. of decinormal iodin may be introduced, even 
 when the sodium bicarbonate used is of exceptional purity. 
 
ANALYSIS BY OXIDATION AND REDUCTION 195 
 
 and especially proven to be free from carbonate, sulphite or 
 thiosulphate. He shows that when sodium bicarbonate is 
 added to a decinormal iodin solution, residual titration with 
 sodium thiosulphate will show a considerable loss of free 
 iodin, which went into combination in some form or other 
 (probably iodid) and that the quantity so lost is proportional 
 to (i) the mass of sodium bicarbonate; (2) the time of the 
 interaction (the reaction is slow); (3) the concentration of 
 the solution; (4) the temperature, and (5) the size of the flask 
 in which reaction occurs. These phenomena are due to the 
 fact that sodium bicarbonate when dissolved in water under- 
 goes hydrolysis, thus 
 
 2NaHC03 = Na2C03 + H2C03 or (H2O + CO2). 
 
 This breaking up of the NaHCOs into Na2C03 and 
 H2CO3, and the latter into H2O and CO2, continues until the 
 pressure of the CO2 above is equal to the pressure of the gas 
 in the solution, i.e., until equilibrium has been reached. In 
 concentrated solutions of NaHCOs the amount hydrolyzed is 
 much greater than in dilute solutions. An elevation of temper- 
 ature materially increases the absorption of iodin. 
 
 Less iodin is lost when smaller flasks are used, provided 
 the glass stopper completely shuts off communication with 
 the atmosphere. The CO2 will escape from the solution 
 until its pressure in the solution is equal to that of the gas 
 above. Thus, since a larger volume of air is contained in 
 a larger flask, more CO2 passes from the liquid before equi- 
 librium is estabhshed, hence more NaHCOs is decomposed, 
 and more iodin in consequence absorbed.* 
 
 Reasoning from the above observations it may be said 
 
 * For further study of equilibrium, see the work of Dr. H. N. McCoy, 
 Am. Ch. J., vol. XXIV, 437- 
 
196 THE ESSENTIALS OF \'OLUAIETRIC ANALYSIS 
 
 that: I, though sufficient sodium bicarbonate be used to 
 more than neutrahze the hydriodic acid formed, the solution 
 titrated should be well diluted; 2, that the titration should 
 be done cold; 3, that the titration should be done in small 
 stoppered flasks, and 4, it should be done quickly. 
 
 Estimation of Arsenous Compounds 
 
 These compounds are estimated by means of iodin in a 
 manner similar to that described under standardization of 
 iodin solution by means of arsenous oxid. The method is 
 as follows: 
 
 Arsenous Oxid {Arsenous Acid, Arseiwus Anhydrid, Arsenic 
 Trioxid){As203='-!g'j.g2). When arsenous acid is brought in 
 contact with iodin in the presence of water and an alkali, 
 it is oxidized into arsenic acid and the iodin is decolorized. 
 The reaction is: 
 
 As20n + 2l2 + 2H20 = As205 + 4HI; 
 
 NaHCOa + HI = NaT + H2O + CO2. 
 
 The alkali should be in sufficient quantity to combine 
 with the hydriodic acid formed, and must be in the form of 
 potassium or sodium bicarbonate. 
 
 The hydroxids or carbonates should not be used. Starch 
 solution is used as the indicator, a blue color being formed 
 as soon as the arsenous acid is entirely oxidized into arsenic 
 acid. 
 
 0.1 gm. of arsenous acid is accurately weighed and dis- 
 solved, together with about i gm. of sodium bicarbonate, in 
 20 cc. of water heated to boiling. Allow the liquid to cool, 
 add a few drops of starch solution, and allow the decinormal 
 iodin to flow in, shaking or stirring the mixture constantly 
 
ANALYSIS BY OXIDATION AND REDUCTION 107 
 
 until a permanent blue color is produced. The following 
 equation illustrates the reaction: 
 
 AS2O3 + 2H2O + 2I2 = 4HT + AS2O5. 
 
 4)197-92 4 )5°7-68 
 
 io)4C).48 10)126.92 
 
 4.948 gms. 12.692 gms. or 1000 cc. — I V.S. 
 
 10 
 
 N 
 Thus each cc. of — I V.S. represents 0.004948 gm. of 
 
 pure AS2O3. 
 
 Solution of Arsenous Acid and Solution of Potassium 
 Arsenite are assayed in the manner above described. Twenty cc. 
 are taken for the assay, i gm. of sodium bicarbonate added, 
 the solution diluted to 100 cc, and titrated with the deci- 
 normal iodin solution. No indicator is required though starch 
 may be used. 
 
 In the case of solution of potassium arsenite, it is advisable 
 to slightly acidify with hydrochloric acid, then to make the 
 solution alkaline with sodium bicarbonate before titrating. 
 The hydrochloric acid is employed here in order to neutralize 
 any potassium hydroxid which may have been formed through 
 hydrolysis of the potassium bicarbonate contained in the 
 solution. 
 
 Arsenous lodid (Asis = 455.72). This salt is estimated in 
 the same way as described for arsenous oxid. The reaction 
 is illustrated by the following equation: 
 
 2ASI3 + 5H2O + 2I2 = AS2O5 + loHI. 
 
 4 )9"-44 
 io)_227^86_ j^^ 
 
 22.786 gms. = 1000 cc. — I V.S. 
 
 The Direct Percentage Assay of Arsenous Compounds. 
 A quantity of arsenous acid is taken, which is equal to the 
 weight of pure As20-f, oxidized by 100 cc. of decinormal iodin, 
 i.e., 0.4948 gm. 
 
198 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 If 0.4948 gm. of the sample be taken then each cc. of — 
 
 10 
 
 I V.S. will represent rJir °^ ^^^^ quantity or i per cent of 
 
 pure AS2O3. In the case of weak solutions of arsenic, as 
 
 liquor acidi arsenosi, liquor potassii arsenitis, etc., which 
 
 contain only one per cent of arsenous acid. A much larger 
 
 quantity should be taken for analysis, otherwise the quantity 
 
 of standard iodin solution used will be so small as to diminish 
 
 the accuracy of the test. 
 
 Thus, if only 0.4948 gm. of either of the above solutions 
 
 be taken, no more than i cc. of the standard solution would 
 
 be required. It is better to take enough of the preparation 
 
 to use up 30 to 50 cc. of standard solution. 
 
 Estimation of Antimony Compounds 
 
 Antimonous oxid (Sb203) or any of its compoxmds may 
 be accurately estimated by means of iodin, in a manner similar 
 to that described for the estimation of arsenous oxid, the 
 antimonous oxid being oxidized to antimonic oxid, as per 
 equation, 
 
 SbaOs + 2H2O + 2I2 = 4HI + SbsOs. 
 
 The antimonous oxid is dissolved and kept in solution by 
 
 the aid of tartaric acid, and then after the addition of an 
 
 N 
 excess of sodium-bicarbonate, the solution is titrated with — 
 
 10 
 
 iodin, using starch as an indicator. Accurate results can only 
 
 be obtained if the solution is sufficiently alkaline to neutralize 
 
 the hydriodic acid formed during the reaction. The titration 
 
 should be conducted without delay after the addition of the 
 
 bicarbonate, otherwise a precipitate of antimonous hydrate 
 
 will be formed, upon which iodin has little effect. The anti- 
 
ANALYSIS BY OXIDATION AND REDUCTION 199 
 
 mony must be in solution to be properly attacked by the 
 iodin. 
 
 To O.I gm. of antimonous oxid 20 cc. of water are added 
 and the mixture heated to boiling; to this tartaric acid is 
 added in small portions at a time until the oxid is completely 
 dissolved. The solution is then • neutralized by means of 
 sodium carbonate, and sufficient of a saturated solution of 
 sodium bicarbonate is added to make the solution distinctly 
 alkaline (about 10 cc. is required for o.i gm. of the antimonous 
 oxid). The mixture is now ready for titration with standard 
 iodin solution. This should be done immediately. The 
 appearance of a permanent blue color marks the end-point, 
 starch being used as indicator. 
 
 SbaOa + 2H2O + 2I2 = 4HI + SbaOs. 
 
 4 )288.4 4 )507.68 
 
 10)72.1 10)126.92 
 
 7.21 gms. 12.692 gms. or 1000 cc. — V.S. 
 
 10 
 
 N 
 One cc. of — iodin represents 0.00721 gm. of SbaOs. 
 
 The solution of the oxid may be made by means of hydro- 
 chloric acid, and after adding a portion of tartaric and diluting 
 with water, sodium bicarbonate is added and the titration 
 conducted as above. 
 
 Other compounds of antimony may be estimated in the 
 same way. Antimonic compounds are reduced to antimonous 
 sulphid (SbaSs) by precipitating with hydrogen sulphid, and 
 after thoroughly washing the precipitate, dissolving it in hydro- 
 chloric acid; thus a solution of antimonous chlorid is obtained 
 from which all traces of hydrogen sulphid are expelled by 
 boiling. This solution is diluted with water, tartaric acid 
 added, and finally, after making alkaline with sodium bicar- 
 bonate, titrated with the standard iodin solution as above 
 described. 
 
200 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 ' Antimony and Potassium Tartrate (Tartar Emetic) 
 
 [2(K[SbO]C4H406)+H20 = 564.7]. i gm. of the salt is 
 
 dissolved in sufficient water to make loo cc. 30 cc. of this 
 
 solution, representing 0.3 gm. of the salt, are taken for assay. 
 
 20 cc. of a cold saturated solution of sodium bicarbonate are 
 
 added, then a little starch .solution, and the mixture titrated 
 
 N 
 with — iodin until a permanent blue color appears. 
 
 The calculation is as follows: 
 2K(SbO)C4H,06 + H2O + 2I2 + 3H2O 
 
 4)664.7 4)5°7-68 
 
 10)166.17 10)126.92 
 
 16.617 gms. 12.692 gms. = 1000 cc. — V.S. 
 
 = 4HI + 2KHC4H4O6 + 2HSbOa. 
 
 N 
 I cc. of — iodin represents 0.016617 g™- o^ 2K(SbO)C4H406 
 
 + H2O (crystallized tartar emetic). 
 
 K(SbO)C4H406 (anhydrous tartar emetic) = 323.34. 
 
 1° ) 161 -67 j^ 
 
 16.167 gms. = 1000 '^'^- — V-S- 
 
 N 
 Thus I cc. of — iodin represents 0.016167 gm. of anhydrous 
 
 tartar emetic. 
 
 Estimation of Sulphurous Acid and Sulphites 
 
 These substances may be accurately estimated by means 
 of a standard solution of iodin. When sulphurous acid or 
 one of its salts is brought in contact with iodin, a complete 
 oxidation takes place. The sulphurous acid is oxidized to 
 
ANALYSIS BY OXIDATION AND RKDUCTION 201 
 
 sulphuric acid and the sulphite to a sulphate, as the equations 
 show: 
 
 H2S03 + H20 + T2 = 2HI+H2S04, 
 
 Na2S03+H20 + l2 = 2HI + Na2S04, 
 
 NaHS03 + H20+l2-=2HI+NaHS04. 
 
 There are two methods which may be employed. In one 
 method the substance is brought into solution in water, an 
 excess of sodium bicarbonate is added, and then the standard 
 iodin solution is run in until a faint yellow color of free iodin 
 marks the end-reaction. If starch solution is used as indicator 
 the end-point is the production of a blue color. The other 
 method is that of Giles and Shearer, who, in a ver}' voluable 
 series of experiments detailed in the J. S. C. I., Ill, 197, and 
 IV, 303, suggest the following modification: 
 
 The weighed sulphurous acid or the sulphite (in fine 
 
 N 
 powder) is added to an accurately measured excess of — 
 
 iodin, without diluting with water. After the mixture has 
 
 been allowed to stand for about one hour, with frequent 
 
 shaking, the oxidation is complete, and the excess of iodin 
 
 N 
 is ascertained by titrating back with — sodium thiosulphate. 
 
 The quantity of the latter deducted from the quantity 
 
 N 
 of — iodin solution added, will give the quantity of the latter 
 10 
 
 which reacted with the sulphite. 
 
 The neutral and acid sulphites of the alkalies, alkali earths, 
 
 and even zinc and aluminum, may be accurately estimated 
 
 in this manner. The less soluble salts requiring, of course, 
 
 more time and shaking, to insure their complete oxidation. 
 
 The latter is the U. S. P. method. 
 
202 THE ESSENTIALS OF VOLU.METRIC ANALYSIS 
 
 Sulphurous Acid. This is an aqueous solution of sulphur 
 dioxid (502 = 64.07). 
 
 Sulphurous acid when brought in contact with iodin is 
 oxidized into sulphuric, the iodin being decolorized because of 
 its union with the hydrogen of the accompanying water, 
 forming hydriodic acid. 
 
 Two grams of sulphiu-ous acid are taken and diluted with 
 distilled water (recently boiled and cooled *) to about 25 
 cc. Two grams of sodium bicarbonate are added, and then 
 the decinormal iodin WS. is delivered into the solution (to 
 which a little starch solution had been previously added) 
 imtil a permanent blue color is produced. 
 
 The following equations, etc., show the reactions that take 
 place: 
 
 H2S03+H20+l2 = 2HI+H2S04. 
 
 Sulphurous acid being, however, looked upon as a solu- 
 tion of SO2 in water, the quantity of this gas is generally 
 estimated in analyses. 
 
 H20,S02 + H2O + I2 = 2HI + H2SO4. 
 
 2 )6407 ^ )253S4 
 
 10)32. o^ 10)126. q2 
 
 3.203 gms. 12.692 gms. 
 
 N 
 Thus each cc. of — iodin consumed before the blue color 
 10 
 
 appears, represents 0.003203 gm. of SO2. 
 
 The Residual Method. Because of the volatile nature of 
 
 this acid the residual method described below is the most 
 
 satisfactory in that loss by volatilization is avoided and com- 
 
 * "Recently boiled" insures absence of air, the oxygen of which would 
 partially oxidize the sulphurous acid, and "cooled" is directed to avoid loss 
 of SO2, which would occur if hot water were used. 
 
ANALYSIS BY OXIDATION AND REDUCTION 
 
 203 
 
 plete oxidation of the acid assured. When the direct method 
 described is used there is more or less loss of SO2 and incom- 
 plete oxidation, with separation of sulphur. 
 
 Measure 2 cc. of the sulphurous acid into a stoppered 
 weighing flask and find its exact weight. Add this to the 
 
 N 
 ■JO cc. of — iodin contained in a titration flask and let the 
 ■^ 10 
 
 solution stand for about five minutes. Then titrate with 
 
 N 
 10 
 
 sodium thiosulphate until the mixture is decolorized. Subtract 
 
 the number of cc. of the thiosulphate used 
 
 N 
 from the <^o cc. of — iodin added, and multi- 
 ^ 10 
 
 N 
 ply the difference bv the — factor for SO2, 
 ^ •' "10 
 
 which is 0.0032035 gm. This will give the 
 
 weight of SO2 in the quantity of acid taken for 
 
 analysis. 
 
 When a solution containing sulphur dioxid 
 is to be measured by means of a pipette, it is 
 never advisable to fill the instrument by 
 suction in the usual manner, as this would 
 cause a loss of the gas. A better plan is to fill 
 the pipette by pressure by the use of an 
 arrangement similar to that shown in Fig. 46. 
 
 The' solution containing sulphur dioxid or 
 other volatile substance is poured into a flask 
 which is provided with a stopper through 
 which two glass tubes pass; one of these 
 tubes reaches nearly to the bottom of the flask 
 and the other projects about one- half an inch 
 below the stopper and is bent outward above. To the upper 
 end of the former the pipette is. attached by means of a piece of 
 
 Fro. 46. 
 
204 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 rubber tubing. By blowing into the flask through the shorter 
 tube the liquid is caused to rise and fill the pipette, which may 
 then be easily pulled out of the rubber tube connection. 
 
 Sodium Sulphite (Na2S03 + 7H20 = 252.2). One gm. of the 
 salt is dissolved in 25 cc. of distilled water recently boiled to 
 expel air, and after the addition of an excess of sodium bicar- 
 bonate a little starch T.S. is added, and then the decinormal 
 iodin V.S. delivered in from a burette, until the blue color 
 of starch iodid appears, which does not disappear upon shaking 
 or stirring. 
 
 The reaction is expressed as follows: 
 
 NazSOa + 7H2O + I2 = 2HI + Na2S04 + 6H2O. 
 
 2)2^2.^ 2)25.^.84 
 
 10)126.1 10)126.0" 
 
 N . 
 
 12.61 gms. 12.692 gms. or looo cc. — lodin V.S. 
 
 Thus each cc. of the standard solution represents 0.01261 
 gm. of crystallized sodium sulphite. 
 
 If I gm. .of the salt is taken, to find the percentage multi- 
 ply the factor by the number of cc. of standard solution con- 
 sumed, and the result by 100. 
 
 For the residual method take 0.5 gm. of the finely powdered 
 
 N . . 
 crystals, add to 50 cc. of — iodin, contained in a loo-cc. glass- 
 stoppered flask, and allow to stand for one hour (shaking 
 
 N 
 frequently); then titrate with — sodium thiosulphate until 
 
 the color is discharged. 
 
 Potassium Sulphite (K2SO3 -1-21120 = 194.37). Operate 
 upon 0.5 gm. in the same manner as for sodium sulphite. 
 
 K2SO3 + 2H2O -H I2 = 2HI -f K2SO4 -h H2O. 
 
 2 )194-37 
 10)97.18 
 
 9.718 gms. or 1000 cc. of standard V.S. 
 
ANALYSIS BY OXIDATION AND REDUCTION 205 
 
 N 
 
 Each cc. of the — iodin represents 0.009718 gm. of crys- 
 tallized potassium sulphite. 
 
 Sodium Bisulphite (NaHS04= 104.08). Operate upon 
 about 0.25 gm. in the same manner as for sodium sulphite, 
 and apply the following equation: 
 
 NaHS03 + l2 + H20 = 2HI+NaHS04. 
 
 Sodium Thiosulphate (Sodium Hyposulphite) (Na2S203 
 + 51^20 = 248.24). This salt, when brought in contact with 
 iodin, is converted into sodium iodid and sodium tetrathionate. 
 The reaction is expressed by the equation 
 
 2Na2S203 + 12 = 2NaI +Na2Sj06. 
 
 It is estimated as follows: i gm. of the salt is dissolved 
 
 in 20 cc. of water, a few drops of starch solution are added, 
 
 N . 
 and then the — iodin is delivered in from a burette, until 
 10 
 
 the appearance of blue starch iodid indicates an excess of 
 iodin. 
 
 Hydrogen Sulphid (H2S = 34.07) . When iodin and hydrogen 
 sulphid are brought together in solution the following reaction 
 occurs : 
 
 H2S4-2l = 2HI + S. 
 
 The reaction is not regular, however, when performed in 
 an acid solution, but in the presence of alkali bicarbonates 
 the results are constant. The method may "be employed for 
 the estimation of alkali sulphates. 
 
 The process may be conducted as follows : 
 
 Into 30 cc. of a cold saturated solution of sodium bicar- 
 
206 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 bonate, contained in a 500-cc. flask, measure a suitable quantity 
 
 of the solution of hydrogen sulphid, stopper the flask and 
 
 mix contents by shaking. Dilute the solution with about 300 
 
 N 
 cc. of water, add starch solution and titrate with — iodin 
 
 10 
 
 V.S. until a distinct and permanent blue color appears. 
 
 N 
 Each cc. of — iodin represents 0.0017035 gm. of H2S. 
 
 The residual method may also be employed. A suitable 
 
 N 
 volume of the sample is added to an excess of — iodin V.S. 
 
 mixed with some sodium bicar- 
 bonate solution, the solution 
 is thoroughly shaken, and then 
 
 N 
 titrated with — thiosulphate; the 
 
 quantity of the latter, deducted 
 
 N 
 from the quantity of — iodin 
 
 added, gives the quantity of 
 
 N . . 
 
 — iodin which reacted with 
 
 10 
 
 the H2S. 
 
 Sulphids. Soluble sulphids 
 
 Fig. 47. may be estimated by either 
 
 of the above methods. The 
 
 solution of sulphid containing about 0.2 gm. being treated 
 
 like an H2S solution. 
 
 Sulphids insoluble in water but decomposable by dilute 
 
 acids, may be estimated as follows : 
 
 A weighed quantity of the sulphid is introduced into a 
 
 flask, provided with a double perforated stopper; through 
 
 one of the perforations the stem of a separatory funnel is 
 
ANALYSIS BY OXIDATION AND REDUCTION 207 
 
 passed, through the other a glass delivery tube (see Fig. 47). 
 The funnel tube extends nearly to the bottom of the flask and 
 is bent to form a hook, the opening of which is under water. 
 The delivery tube begins at the lower end of the stopper and 
 ends in another flask containing sodium bicarbonate solution. 
 The fimnel contains diluted sulphuric acid, which," upon open- 
 ing the glass stop-cock, is allowed to flow into the flask, upon 
 the contained sulphid; the H2S liberated is conducted into 
 the solution of sodium bicarbonate which absorbs it com- 
 pletely. A current of air aspirated through the apparatus 
 insures absorption of the entire H2S developed. The sodium 
 bicarbonate solution of H2S is then titrated with the standard 
 iodin, in the presence of starch. 
 
 Estimation of Metallic Iron in Reduced Iron 
 
 Professor E. Schmidt, of Marburg, recommends * the fol- 
 lowing process: 
 
 Weigh accurately 0.4 gm. of reduced iron, and place in a 
 loo-cc. flask with 10 cc. of water, and add 2 gms. of pure 
 dry iodin. Now rinse down the iodin left in the neck of the 
 flask with some water, and add 2 gms. of potassium iodid; 
 when all of the iodin has dissolved, add sufficient water to make 
 100 cc. Shake the flask and allow to stand for several 
 hours. 
 
 The iodin combines with metallic iron, but does not react 
 with any ferric oxid which may be present. 
 
 Fe2-l-2l2 = 2Fel2. 
 
 Then measure off 50 cc. of the clear liquid and titrate 
 
 * Proc. Soc. German Naturalists and Physicians, Sept., 1897. 
 
208 THK ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 the free iodin with decinormal sodium thiosulphate, using 
 starch as an indicator. The reaction is thus expressed : 
 
 I2 + 2Na2S203, 5H2O = NaaSiOo + 2NaI + 10H2O. 
 2 )253-^4 2 )496-28 
 10)126.92 10)248.24 „ 
 
 12.692 gms. 24.824 gms. or looo cc. — \^.S. 
 0.012692 gm. I cc. " " 
 
 Example. Assuming that 9 cc. of the decinormal solution 
 were employed in titrating the 50 cc, then 18 cc. would be 
 required for the entire quantity. 
 
 As seen in the above equation, each cc. of the decinormal 
 solution represents 0.012692 gm. of iodin; hence if 18 cc. 
 are employed we have 18X0.01269 gm. =0.228456 gm., the 
 quantity of free iodin. 
 
 Then by subtracting this amount from the quantity of 
 iodin taken (2 gms.) we ascertain the quantity which went 
 into combination with the iron, namely, 1.7715 gms. All 
 that is now necessary is to ascertain by calculation the quantity 
 of metallic iron represented by the above weight of iodin. 
 
 Fe + I2 = Felg; 
 
 55.82 253.84 
 
 i.77it;X«.82 
 " ^ o ^ — = 0.389+ gm. 
 253.84 -3 y B 
 
 Thus the 0.40 gm. of reduced iron taken contained 0.389 + 
 |m. of metalHc iron, or 87.7 per cent. 
 
 Many qther substances besides those mentioned in the 
 foregoing pages may be estimated by titration with standard 
 iodin solution. Among them are cyanids, stannous compounds, 
 mercurous compounds, metallic zinc, and aluminum. 
 
ANALYSIS BY OXIDATION AND REDUCTION 
 
 209 
 
 TABLE OF SUBSTANCES WHICH MAY BE ESTIMATED BY MEANS 
 OF STANDARD lODIN SOLUTION 
 
 Name. 
 
 Formula. 
 
 Molecular 
 Weight. 
 
 — Factor. 
 
 Acid, sulphurous 
 
 Antimonous oxid 
 
 Antimony and potas'm tartrate 
 
 Arsenous iodid 
 
 " oxid 
 
 Cyanogen 
 
 Hydrogen sulphid 
 
 Iron (metallic) 
 
 Mercuric chlorid 
 
 Mercurous chlorid 
 
 Potassium cyanid 
 
 " sulphite (anhydrous) 
 " (crystallized) 
 Sodium bisulphite 
 
 " sulphite (anhydrous). 
 
 " " (crystallized) . 
 
 " thiosulphate 
 
 Sulphur dioxid 
 
 Tin in stannous compounds .. . 
 Zinc 
 
 H2SO, 
 
 SbjOa 
 
 ■[K(SbO)C4He08] + H20 
 
 Aslj 
 
 AS2O3 
 
 CN 
 
 H^S 
 
 Fe, 
 
 HgClj 
 
 HgCl 
 
 KCN 
 
 K2SO3 
 
 K2SO3+2H2O 
 
 NaHSOa 
 
 NajSOj 
 
 Na2S03+ 7H2O 
 
 Na2S203+5H20 
 
 SO2 
 
 Suj 
 
 Zn, 
 
 82.07 
 288.4 
 664.7 
 
 455-72 
 
 197.92 
 
 26.01 
 
 34-07 
 I I I . 64 
 270.92 
 235-46 
 
 65.11 
 158.27 
 
 194-37 
 104.08 
 126.07 
 252.2 
 248 . 24 
 64.07 
 238.0 
 130-74 
 
 0.004103 
 
 0.00721 
 
 0.016617 
 
 0.022786 
 
 o . 004948 
 
 0.0013005 
 
 0.0017035 
 
 0.002791 
 
 0.027092 
 
 0.023546 
 
 0-003255 
 
 0.007913 
 
 0.009718 
 
 o . 005 204 
 
 o . 006303 
 
 O.OI26I 
 
 0.024824 
 
 0.003203s 
 0.00595 
 
 0.003268 
 
 Estimation of Substances Readily Reduced. 
 
 Any substance which readily yields oxygen in a definite 
 quantity, or is susceptible of an equivalent action, which 
 involves its reduction to a lower quantivalence, may be quan- 
 titatively tested by ascertaining how much of a reducing agent 
 of known power is required by a given quantity of the sub- 
 stance foi* its complete reduction. 
 
 The principal reducing agents which may be employed 
 in volumetric analysis are sodium thiosulphate, sulphurous acid, 
 arsenous acid, oxalic acid, metallic zinc, and magnesium. 
 
 The sodium thiosulphate is the only one which is employed 
 ofi&cially in th^ U. S. P. in the form of a volumetric solution. 
 
210 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 It is used in the estimation of free iodin, and indirectly of 
 other free halogens, or compounds in which the halogen is 
 easily liberated, as in the hypochlorites, etc. 
 
 Estimations Involving the Use of Sodium Thiosulphate V.S. 
 
 (lodometry) 
 
 When sodium thiosulphate acts upon iodin, sodium tetra- 
 thiohate and sodium iodid are formed, and the solution is 
 decolorized. 
 
 This reaction takes place in definite proportions: one 
 molecular weight of the thiosulphate absorbs one atomic weight 
 of iodin. 
 
 2Na2S203 +l2 = 2NaI +Na2S406. 
 
 Chlorin cannot be directly titrated with the thiosulphate, 
 but by adding to the solution containing free chlorin an excess 
 of potassium iodid, the iodin is liberated in exact proportion 
 to the quantity of chlorin present, atom for atom. 
 
 Cl2 + 2KI = 2KCl+l2. 
 
 Then by estimating the iodin, the quantity of chlorin is 
 ascertained. All bodies which contain available chlorin, or 
 which when treated with hydrochloric acid evolve chlorin, 
 may be estimated by this method. 
 
 Also, bodies which contain available oxygen, and which 
 when boiled with hydrochloric acid evolve chlorin, such as 
 manganates, chromates, peroxids, etc., may be estimated in this 
 way. 
 
 Solutions of ferric salts, when acidulated and boiled with 
 an excess of potassium iodid, liberate iodin in exact propor- 
 tion to the quantity of ferric iron present. 
 
ANALYSIS BY OXIDATION AND REDUCTION 211 
 
 Thus sodium thiosulphate may be used in the estimation 
 of a great variety of substances with extreme accuracy. 
 
 Preparation of Decinormal Sodium Thiosulphate (Hypo- 
 sulphite) (Na2Sa03 + 51120 = 248.48; conta,ins 24.848 gms. in 
 I liter). Sodium thiosulphate is a salt of thiosulphuric acid 
 in which two atoms of hydrogen have been replaced by sodium; 
 it therefore seems that a normal solution of this salt should 
 contain one-half the molecular weight in grams in one liter. 
 
 But this salt is used chiefly for the estimation of iodin, 
 and, as stated before, one full molecular weight reacts with 
 and decolorizes one atomic weight of iodin, and since one 
 atom of iodin is chemically equivalent to one atom of hydrogen, 
 a full molecular weight of sodium thiosulphate must be con- 
 tained in a liter of its normal solution. 
 
 Sodium thiosulphate is easily obtained in a pure state, 
 and therefore the proper weight of the salt, reduced to powder 
 and dried between sheets of blotting-paper, nSay be dissolved 
 directly in water, and made up to one liter. 
 
 A stronger solution than decinormal is usually made, its 
 titer found, and then the solution diluted to the proper measure. 
 
 Thirty gms. of selected crystals of the salt are dissolved in 
 enough water to make, at or near 15° C. (59° F.), 1000 cc. 
 
 This concentrated solution is then standardized by one of 
 the following methods: 
 
 JV 
 a. HUitulardiMition hy Meanfi of jTi Iodin. Transfer 10 
 
 cc. of this solution into a flask or beaker, add a few drops of 
 starch T.S., and then gradually deliver into it from a burette 
 decinormal iodin solution, in small portions at a time, shaking 
 the flask after each addition, and regulating the flow to drops 
 toward the end of the operation. As soon as a blue color is 
 produced which does not disappear upon shaking, but is not 
 deeper than pale blue, the reaction is completed. Note the 
 
212 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 number of cc. of iodin solution used, and then dilute the thio- 
 sulphate solution so that equal volumes of it and the deci- 
 normal iodin will exactly correspond to each other, under the 
 above-mentioned conditions. 
 
 Example. The to cc. of sodium thiosulphate, we will 
 assume, require 10.7 cc of decinormal iodin. 
 
 The sodium-thiosulphate solution must then be diluted 
 in the proportion of 10 cc. to 10.7 cc, or 1000 cc. to 1070 cc. 
 
 After the solution is thus diluted a new trial should be 
 made, in the manner above described, in which 50 cc. of the 
 thiosulphate solution should require exactly 50 cc. of the 
 decinormal iodin to produce a faint blue color. 
 
 The solution should be kept in small dark amber-colored, 
 glass-stoppered bottles, carefully protected from dust, air, 
 and light. 
 
 One cc. of this solution is the equivalent of: 
 
 Iodin 0.012692 gram. 
 
 Bromin 0.007992 " 
 
 Chlorin .' 0.003546 " 
 
 Iron in ferric salts 0.005582 " 
 
 i. Standai-fJizatiou bjj 3Iean<i of Potassium DU-hromate, 
 
 The potassium dichromate should be pure, if not it should be 
 purified by triple recrystallization and then heated in a porce- 
 lain crucible until the entire mass is just fused; then set 
 aside to cool in a desiccator over calcium chlorid or siilphuric 
 acid. The mass falls to a crystalline powder. Of this purified 
 potassium dichromate, weigh off a definite quantity, say 0.2 
 gm., dissolve it in a small quantity of water, and pour the 
 solution into a beaker containing 2 gms. of pure potassium 
 iodid (free from iodate) and 100 cc. of water. Acidulate the 
 solution with 5 cc. of concentrated hydrochloric or sulphuric 
 acid, cover the beaker, and let stand for about five minutes, 
 
ANALYSIS BY OXIDATION AND R1:DUCTI0N 213 
 
 then titrate with the thiosulphate solution to be standardized 
 (using starch as an indicator) until the blue color is just 
 discharged. The calculation is then made as follows : 
 
 KaCr^Oy + 6KI + i4HCl = 2CrCl3 + 8KC1 + 7H2O + 3I2. 
 294.4 996.12 761.52 
 
 Thus 294.4 gms. of potassium dichromate oxidizes 996.12 
 gms. of potassium iodid and liberates therefrom 761.52 gms. 
 of iodin. 
 
 The potassium iodid must be in excess; in fact for each 
 atom of iodin liberated one molecule of potassium iodid must 
 be present at the completion of the reaction in order to keep 
 the iodin in solution, and thus prevent loss by volatilization. 
 
 If 294.4 gms. of potassium dichromate hberate 761.52 gms. 
 of iodin, 0.2 gm. will liberate 
 
 761.52X0.2 . . ,. 
 
 — ^^ — = 0-5173 gm. of 10dm. 
 
 N 
 0.012692 gm. of iodin=i cc. of — thiosulphate. 
 
 N 
 0.5173 " " " =40.172 cc. of — thiosulphate. 
 
 Therefore if in the above assay 37.60 cc. of the thiosul- 
 phate V.S. were consumed, it must be diluted so that each 
 37.60 cc. will measure 40.72 cc. in order to convert the thio- 
 sulphate solution into a true decinormal solution. A new 
 trial should then be made with the diluted solution to see if 
 its strength is correct. 
 
 It is usually more convenient to use a decinormal dichro- 
 mate V.S. for standardizing decinormal thiosulphate V.S. 
 This may be done as follows: 
 
 Dissolve one gram of pure potassium iodid in a small 
 
21-i THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 quantity (lo to 15 cc.) of sulphuric acid (i-io). Place this 
 into a 500-cc. flask and add to it ver}^ slowly 25 cc. (accurately 
 measured) of decinormal potassium dichromate V.S., mix 
 well and let stand for five minutes, the flask being kept closed. 
 The thiosulphate solution to be standardized is then run in 
 from a burette in small portions, shaking after each addition 
 until the solution is a pale yellow color. Two cc. of starch solu- 
 tion are then added and the titration continued drop by drop 
 until the blue color of starch iodid is just discharged. 
 
 Several trials are made, and from an average of three or 
 four closely agreeing results the quantity of water required 
 for dilution to decinormal strength is readily calculated. 
 
 Example. Assuming that in the above trial 23.8 cc. of 
 the thiosulphate solution were required to react with the 
 iodin liberated by 25 cc. of decinormal dichromate, then the 
 thiosulphate must be diluted with distilled water so that each 
 23.8 cc. will measure 25 cc. After dilution a new trial should 
 be made, in which 25 cc. of decinormal dichromate should 
 require when treated as above described exactly 25 cc. of the 
 thiosulphate. 
 
 c. Standardisation by Means of Potassium Bi-iodate. 
 This method depends upon the fact that when potassium iodid 
 and bi-iodate are brought together in the presence of a small 
 quantity of an acid, an equivalent amount of iodin is set free. 
 The reaction is illustrated by the equation: 
 
 KH(I03)2 + loKI + iiHCl = 12I + iiKCl + 6H2O. 
 
 12 )389.94 12 )1523.04 
 
 10 )32.495 1 0)126.92 
 
 3-2495 gros. 12.692 gms. or looo cc. — V.S. 
 
 10 
 
 Thus it is seen that one molecule of the bi-iodate causes 
 the liberation of 12 atoms of iodin, against which the thio- 
 sulphate solution is standardized. 
 
ANALYSIS BY OXIDATION AND REDUCTION 215 
 
 In order not to use up too large a quantity of the thio- 
 sulphate solution in the titration, a very small quantity of 
 the bi-iodate must be taken, and since a small error in weighing 
 this would entail a relatively large error in the results it is 
 best to use the bi-iodate in the form of a solution of known 
 strength and thus obviate the difficulty. 
 
 A decinormal solution of the bi-iodate, i.e., one containing 
 in looo cc. 3.2245 gms. of the salt may be used to advantage. 
 Such a solution will keep unchanged for years, and may be 
 employed as follows for the standardization of sodium thio- 
 sulphate solution: 
 
 Into a glass-stoppered flask of 250-cc. capacity introduce 
 10 cc. of a 5 per cent solution of potassium iodid,* one cc. of 
 diluted hydrochloric acid and exactly 25 cc. of the above 
 potassium bi-iodate solution. 
 
 This brownish-yellow solution is now titrated with the 
 sodium thiosulphate solution which is slowly delivered from 
 a burette until the solution becomes pale yellow in color; a 
 few drops of starch solution are now added, and the titration 
 continued (the flow being reduced to drops) until the blue 
 color is just discharged. During the titration, the flask should 
 be frequently stoppered and vigorously shaken. 
 
 WTien the blue color is discharged, note the number of 
 CCS used. If an exactly decinormal thiosulphate solution is 
 taken, 25 cc. will be required to react with the 25 cc. of 
 potassium bi-iodate solution, under the above conditions. If, 
 on the other hand, only 22 cc. of the thiosulphate solution 
 are used in the titration, then the latter is too strong, and 
 must be diluted so that each 22 cc. will measure 25 cc. in 
 order to make it strictly decinormal. 
 
 * The potassium iodid must be in sufficient quantity not only to react 
 with the bi-iodate quantitatively, as shown in the equation, but also to dissolve 
 the iodin which is liberated. 
 
216 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 d. Standardization by Means of Potassium Permanga- 
 nate. Sodium thiosulphate solution may be accurately 
 standardized by means of potassium permanganate, if the 
 other substances used for this purpose are not at hand. The 
 method is easily understood by referring to the iodometric 
 standardization of potassium permanganate solution. See 
 page 139. 
 
 Estimation of Free lodin (1=126.92). lodin in the dry 
 state or in the form . of a tincture is brought into aqueous 
 solution by means of pure potassium iodid and then titrated 
 with standard sodium thiosulphate. The potassium iodid is 
 used here to dissolve the iodin; it must be free from iodate 
 (KIO3) because the presence of this salt would cause a libera- 
 tion of iodin from the potassium io,did. 
 
 Dry iodin is assayed as follows: 
 
 About 0.5 gm. of iodin are placed in a tightly stoppered 
 weighing bottle and accurately weighed. One gram of potas- 
 sium iodid and 50 cc. of water are added, and when the iodin 
 
 N 
 is dissolved, — sodium thiosulphate is delivered from a burette 
 
 in small portions at a time, shaking after each addition until 
 the reaction is nearly completed, and the solution is of a faint 
 yellow color. A few drops of starch indicator are now added, 
 and the titration with the thiosulphate continued drop by 
 drop vmtil a final drop just discharges the blue color. The 
 number of cc. of the thiosulphate solution used is noted. 
 This number, multiplied by the decinormal factor for iodin, 
 gives the weight of the latter present in the sample 
 assayed. 
 
 I2 -t- 2(NaS203 -I- 5H2O) = Na2S406 -t- 2NaI + 10H2O. 
 2 )253.84 2 )496.96 
 
 10)126.92 10 )248.48 
 
 12.692 gms. 24.848 gms. or looo cc. — V.S. 
 
ANALYSIS BY OXIDATION AND REDUCTION 217 
 
 N 
 Each cc. of — thiosulphate represents 0.012692 gm. of 
 
 iodin. 
 
 N 
 If in the above assay 39 cc. of — sodium thiosulphate 
 
 were consumed, then 
 
 0.012692X39 = 0.495 gm. 
 
 0.495 + 1°° 
 
 o-S 
 
 = 99 per cent. 
 
 Lugol's Solution. This is an aqueous solution of iodin 
 and potassium iodid. 
 
 It is estimated for iodin in the same way as the foregoing. 
 The potassium iodid acts merely as a solvent for free iodin, 
 and does not enter into the reaction. 
 
 Ten or twelve grams of the solution is a convenient quantity 
 to operate upon. Starch solution is the indicator. 
 
 Tincture of Iodin. An alcoholic solution of free iodin 
 must be diluted with a solution of potassium iodid, before 
 titration, in order to prevent the precipitation of iodin, which 
 would result upon the addition of the aqueous standard solu- 
 tion. 
 
 Indirect lodometric Estimations. The titration methods 
 previously described, in which iodin is used as a standard 
 solution and the estimation of free iodin by means of sodium 
 thiosulphate, are classed as direct iodometric methods. The 
 following methods, in which the strength of the substance 
 under analysis is determined by the quantity of iodin which 
 it hberates from an iodid, are known as indirect iodometric 
 methods. Potassium iodid is added in excess * to an acidulated 
 solution of the substance, and the liberated iodin estimated 
 
 * The iodid should be in sufficient excess to keep the liberated iodin in 
 solution as KI.I. 
 
218 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 by means of standard thiosulphate. These methods are 
 among the most accurate of all volumetric analyses, and 
 take in a very large class of substances. Among the sub- 
 stances which may be analyzed by this method are chlorin 
 and bromin and all substances which readily liberate these 
 elements: ferric salts, manganates, chromates, metallic perox- 
 ids, and other substances from which oxygen can be easily 
 liberated. 
 
 Free Chlorin or Bromin. Free chlorin acts upon potas- 
 sium iodid, liberating iodin, as per the equation 
 
 Cl2+2KI=2KC1-H2. 
 
 Thus it is seen that each atom of chlorin will liberate one 
 
 atom of iodin, hence by determining the quantity of iodin 
 
 by means of a standard thiosulphate solution the quantity 
 
 of chlorin present is easily ascertained (126.92 gms. of 
 
 1 = 35.46 gms. of CI). The same applies to free bromin, 
 
 one atom of bromin (79.92) will liberate one atom of iodin 
 
 N 
 (126.92) Br2 + 2KI = 2KBr-|-l2. 1000 cc. of — sodium thio- 
 
 sulphate is equivalent to 12.692 gms. of iodin, and hence to 
 
 3.546 gms. of chlorin or 7.992 gms. of bromin. 
 
 N 
 Thus- I cc. of — sodium thiosulphate is equivalent to 
 
 0.012692 gm. of iodin; 0.003546 gm. of chlorin; 0.007992 
 gm. of bromin. 
 
 Chlorin cannot be directly, titrated with sodium thiosul- 
 phate because, instead of the tetrathionate being formed as 
 with iodin, sulphuric acid is produced; furthermore there is 
 no readily observable end-point as there is with iodin. 
 
 Chlorin Water. This is an aqueous solution of chlorin, 
 CI = 35.46, containing at least 0.4 per cent of the gas. 
 
 The estimation of chlorin is effected in an indirect way, 
 
ANALYSIS BY OXIDATION AND REDUCTION 219 
 
 namely, by determining the quantity of iodin which it liberates 
 from potassium iodid. 
 
 A definite quantity of chlorin will liberate a definite quan- 
 tity of iodin from an iodid; these quantities are in exact 
 proportion to their atomic weights, as the equation shows: 
 
 Ch + 2KI = 2KCI + I2. 
 
 2 )io.92 2 )253.84 
 
 10)35.46 10)126.92 
 
 3.546 12.692 gms. 
 
 Thus it IS seen that by estimating the liberated iodin the 
 quantity of chlorin may be determined with accuracy. 
 
 Ten grams is a convenient quantity to operate upon. To 
 this about half a gram of potassium iodid is added. A little 
 starch solution is then introduced, and the titration is begun 
 with decinormal sodium thiosulphate. 
 
 When the blue color of starch iodid has entirely disappeared 
 the reaction is finished. 
 
 The reaction between iodin and sodium thiosulphate is 
 illustrated by the following equation: 
 
 I2 + 2(Na2S20n + 5H2O) = 2NaI+ NazSiOe + 10H2O. 
 2 )253.84 2 )496.96 
 
 1 0)126.92 10 )248.48 j^ 
 
 12.692 gms. 24.848 gms. or 1000 cc. — V.S. 
 
 N 
 Thus we see that 1000 cc. of — Na2S203.5H20 represent 
 
 12.692 gms. of iodin, which are equivalent to 3.546 gms. of 
 
 chlorin. 
 
 Each cc. therefore is equivalent to 0.003518 gm. of chlorin. 
 
 This number is the factor which, when multiplied by the 
 
 N 
 number of cc. of — thiosulphate used, gives the weight in 
 
220 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 grams of chlorin contained in tfie quantity of chlorin water 
 acted upon. 
 
 Chlorinated Lime (Calx Chlorinata, Chlorid of Lime, 
 Bleaching-powder). This substance was formerly supposed 
 to be a compound of lime and chlorin, CaOCl2, and hence 
 the name chlorid of lime. It is now generally considered to 
 be a mixture principally of calcium chlorid and calcium hypo- 
 chlorite, CaCl2 + Ca(C10)2 or Ca(OCl)CI. The hypochlorite 
 is the active constituent. This is a very unstable salt, and 
 is readily decomposed even by carbonic acid. When treated 
 with hydrochloric acid it gives o£E chlorin. 
 
 The value of chlorinated lime as a bleaching or disinfecting 
 agent depends upon its available chlorin, that is, the chlorin 
 which the hypochlorite yields when treated with an acid. 
 
 In estimating the available chlorin, the latter is liberated 
 with hydrochloric acid. This liberated gas, then, acting upon 
 potassium iodid, sets free an equivalent amount of iodin. 
 The quantity of iodin is then determined, and thus the amount 
 of available chlorin found. 0.2 to 0.4 gm. are convenient 
 quantities to operate upon. 
 
 Introduce into a stoppered weighing bottle between 3 
 and 4 gms. of chlorinated lime and weigh accurately. (In 
 order to make the descriptions simpler we will assume that 
 3.5 gms. is the weight taken.) This is triturated thoroughly 
 with 50 cc. of water, and the mixture transferred to a graduated 
 vessel, together with the rinsings, and made up to 1000 cc. 
 with water. This is thoroughly shaken, 100 cc. of it (repre- 
 senting 0.35 gm. of the sample) are removed by means of 
 a pipette and treated with i gm. of potassium iodid * and 
 5 cc. of diluted hydrochloric acid, and into the resulting 
 
 * In order to assume a sufficient excess of potassium iodid, take twice 
 as much of it as of the bleaching-powder. 
 
ANALYSIS BY OXIDATION AND REDUCTION 221 
 
 N 
 reddish-brown liquid the — sodium thiosulphate is delivered 
 
 from a burette. Towards the end of the titration, when 
 the brownish color of the liquid is very faint, a few drops 
 of starch solution are added and the titration continued until 
 the bluish or greenish color produced by the starch has 
 entirely disappeared. Not less than 30 cc. of the volumetric 
 solution should be required to produce this result. 
 
 The reactions which take place in this process are illus- 
 trated by the following equations: 
 
 ' Ca(OCl)Cl-F2HCl = CaCl2-hH20 + Cl2 
 or Ca(OCl)Cl+H2S04 = CaS04-fH20-hCl2, 
 
 Cl2-h2KI = 2KCI-M2. 
 
 2)70.92 2)253.84 
 
 10)35-46 10)126.92 
 
 3.546 gms. 12.692 gms. 
 
 I2 + 2(Na2S203 -f SH2O) = 2NaI + Na2S406 + 10H2O. 
 10 )126.92 j^^ 
 
 12.692 gms. =1000 cc. — thiosulphate V.S. 
 
 It is thus seen that i cc. of the decinormal sodium thio- 
 sulphate represents 0.012692 gm. of iodin, which in turn is 
 equivalent to 0.003546 gm. of chlorin. 
 Then 
 
 0.003546X30 = 0.1063 gm. 
 0.1063 X 100 
 
 0-35 
 
 = 30.37 per cent of available chlorin. 
 
 This is a very rapid method for estimating chlorin; but 
 when calcium chlorate is present in the bleaching-powder 
 (and it often is, through imperfect manufacture) the chlorin 
 from it is recorded, as well as that from the hypochlorite, 
 
222 THE KSSENTIALS OF VOLUMETRIC ANALYSIS 
 
 the chlorate being decomposed into chlorin, etc., by hydro- 
 chloric acid. The chlorate, however, is of no value in bleaching; 
 its chlorin is not available. Hence, unless the powder is known 
 to be free from chlorate, the analysis should be made by means 
 of arsenous-acid solution, or by using acetic acid instead of 
 hydrochloric, and thus avoid liberating chlorin from the 
 chlorate which may be present. 
 
 The various bleaching preparations of the market which 
 depend upon their available chlorin are all salts of hypochlorous 
 acid (HCIO) or solutions of such salts. 
 
 Eau de Javelle (Javelle's Water) is a solution of potas- 
 sium h)rpochIorite and potassium chlorid. A solution of mag- 
 nesium hypochlorite is known in commerce as Ramsay's or 
 Grouvelle's Bleaching Fluid. The solution known as Wilson's 
 Bleaching Fluid contains aluminum hypochlorite. 
 
 Solution of Chlorinated Soda (Labarraque's Solution). 
 This is an aqueous solution of several chlorin compounds 
 of sodium, principally sodium chlorid and hypochlorite, con- 
 taining at least 2.4 per cent by weight of available chlorin. 
 
 In this solution-, as in chlorinated lime, it is the available 
 chlorin which it estimated. The chlorin is first liberated with 
 hydrochloric or sulphuric acid; this then liberates iodin from 
 potassium iodid, and the free iodin is then determined by 
 standard sodium thiosulphate. 
 
 Seven grams of chlorinated soda solution are mixed with 
 50 cc. of water, 2 gms. of potassium iodid, and 10 cc. of hydro- 
 chloric acid are then added, together with a few drops of starch 
 solution. Into this mixture the decinormal sodium thio- 
 sulphate is delivered from a burette until the blue or greenish 
 
 N 
 tint of the liquid is just discharged. Each cc. of — thiosul- 
 
 10 
 
 phate used up represents 0.003546 gm. of available chlorin. 
 
 The potassium iodid should always be added before the hydro- 
 
ANALYSIS BY OXIDATION AND REDUCTION 223 
 
 chloric acid, so that the chlorin has potassium iodid to act upon 
 
 as it is Hberated, and thus loss of chlorin is obviated. 
 
 Bromin Water, or any substance containing free bromin, 
 
 may be assayed in exactly the same manner as that described 
 
 for chlorin water. Free chlorin must, however, be absent. 
 
 N 
 Each cc. of — thiosulphate solution represents 0.007992 gm. 
 
 of bromin. 
 
 Assay of Hydrogen Dioxid (H202 = 34). The iodometric 
 method, which originated with Kingzett,* is based upon the 
 fact that iodin is liberated from potassium iodid by hydrogen 
 dioxid, in the presence of sulphuric acid, and that this libera- 
 tion of iodin is in direct proportion to the available oxygen 
 contained in the dioxid. 
 
 Then by determining the amount of iodin liberated, the 
 available oxygen is readily found. 
 
 H2O2 + H2SO4 + 2KI = K2SO4 + 2H2O + I2. 
 
 2)34 2)i_6 2 )253.84 
 
 17 = 1 available 0= 8 126.92 
 
 This shows that 126.92 gms. of iodin are liberated by 17 
 gms. of absolute dioxid, which are equivalent to 8 gms. of 
 available oxygen. 
 
 N 
 Thus 1000 cc. of — sodium thiosulphate V.S., which absorb 
 
 and consequently represent 12.692 gms. of iodin, are equivalent 
 
 to 1.7 gms. of H2O2 or 0.8 gm. of available oxygen. 
 
 N 
 Each cc. of this — V.S., then, represents 0.0017 S"^- of 
 
 H2O2, and 0.0008 gm. of available oxygen. 
 
 The coefficients for weight of H2O2 and of oxygen, it is seen, 
 are identical with those used in the permanganate process. 
 
 * J. Chem. Sor., 1880, \'oi. 37, p. 792. 
 
224 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Therefore the coefficient for volume is also the same in this 
 
 method as in the other if i cc. be taken for assay. 
 
 The process is carried out as follows: Take 2 or 3 cc. of 
 
 sulphuric acid, dilute it with about 30 cc. of water, add an 
 
 £xcess of potassium iodid (about i gm.), and then i cc. of 
 
 hydrogen oxid. After the mixture has been allowed to stand 
 
 five minutes, starch solution * is added, and the titration with 
 
 N 
 
 — sodium thiosulphate begun. 
 
 Note the number of cc. required to discharge the blue color, 
 and multiply this number: by 0.0017 gm. to find the quantity, 
 by weight, of H2O2; by 0.0008 gm. to find the weight of avail- 
 able oxygen; by 0.57 cc. to find the volume of available 
 oxygen. 
 
 If 18 cc. are required, the solution is of 0.57X18=10.26 
 volume strength. 
 
 0.0017X18=^0.0306 or 3.06 per cent H202. 
 0.0008X18 = 0.0144 or 1.44 per cent of oxygen. 
 
 With this method the author has always obtained satisfactory 
 results. The lack of uniformity in the reaction, which is 
 frequently reported, is doubtless due to the use of insufficient 
 acid or to taking a too concentrated solution of the dioxid. 
 
 The best results are obtained if the solution is not more 
 than two volumes strength. 
 
 The sulphuric acid used in this assay must be free from 
 sulphurous acid, arsenous acid and nitric acid, and the potas- 
 sium iodid must contain no iodate. 
 
 Distillation Methods. Manganates, chromates, metallic 
 peroxids, and a great variety of substances containing oxygen, 
 
 * Starch solution may be omitted, as the decolorization of the iodin is 
 distinctly seen if the beaker is placed upon a white surface. 
 
ANALYSIS BY OXIDATION AND REDUCTION 225 
 
 will, when heated with concentrated hydrochloric, liberate 
 an equivalent amount of chlorin. This is illustrated by the 
 following equation: 
 
 MnOa +4HCI = MnCla + 2H2O + CI2. 
 
 The chlorin which is evolved, is passed into a solution of 
 potassium iodid and liberates an equivalent of iodin, which 
 latter substance is then estimated by titration with sodium 
 thiosulphate solution. The quantity so found is therefore 
 a measure of the original substance and of its oxygen content. 
 The process may be carried out by means of the apparatus 
 
 
 Fig. 48. 
 
 devised by Bunsen, Fig. 48, or by that of Fresenius, Fig. 49, 
 or Mohr, Fig. 50. 
 
 An accurately weighed quantity of the substance to be 
 analyzed is introduced into the round-bottomed flask a, Fig. 
 48. The flask is then filled to about two-thirds its capacity 
 with concentrated hydrochloric acid, and quickly connected 
 by means of a short rubber tube with a long-bulbed delivery 
 tube, h, which is introduced into and extends to the bottom of 
 an inverted bulbed retort, c. The larger bulb of the retort is 
 filled to two-thirds of its capacity with a 10 per cent solution 
 of potassium iodid. Heat is applied to the flask, and the chlorin 
 distils over into the potassium iodid solution, which becomes 
 
226 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 brownish-red through liberation of iodin. The distillation 
 is continued until about one-third of the acid fluid has passed 
 over or until a peculiar cracking sound indicates the absorption 
 of hot hydrochloric acid vapor. 
 
 The flask, together with its delivery tube, is then slowly 
 removed, the heating, however, is continued until the tube 
 
 Fig. 49. 
 
 is entirely withdrawn, in order to prevent the iodid solution 
 being drawn over into the flask. The retort is then shaken 
 so that any traces of chlorin which may have escaped absorp- 
 tion, are taken up and the contents of the retort poured into 
 a beaker; the retort and the delivery tube are then rinsed with 
 water and the rinsings added to the fluid in the beaker, and 
 titration with standard sodiiun thiosulphate begun immediately. 
 It is important that the quantity of potassium iodid is suffi- 
 ci«nt to keep the liberated iodin in solution, and the potas- 
 
ANALYSIS BY OXIDATION AND REDUCTION 227 
 
 sium iodid free from iodate, and the titration started with- 
 out delay to avoid Uberation of iodin through action of the 
 air upon the strongly acid potassium iodid solution. When 
 all the chlorin has passed over and hydrochloric acid gas 
 begins to distil, the liquid in the retort is apt to be drawn back 
 into the flask because of the great affinity which hydrochloric 
 acid gas has for water, and the resultant condensation in the 
 flask. This regurgitation may be avoided by introducing 
 into the generating flask a small piece of magnesite, which 
 slowly dissolves in the acid solution and so keeps up a constant 
 flow of carbon dioxid, which by its pressure prevents back-flow 
 of the fluid. The bulbs in the retort and delivery tube are 
 also calculated to prevent this regurgitation. 
 
 The Fresenius apparatus is illustrated in Fig. 49. In this 
 the potassium iodid solution is contained in two joined U-shaped 
 tubes. The delivery tube from the distilling flask enters one 
 of the U-tubes through a paraffin-soaked cork (which fits 
 tightly), and terminates just above the potassium iodid solu- 
 tion. In operation the U-tubes should be kept in ice water, 
 and all the fittings should be air-tight. Paraffin-covered cork 
 stoppers only should be used. 
 
 After all the chlorin has passed over or when about one-third 
 of the acid has distilled over, the apparatus is allowed to stand 
 for a few minutes, to permit all traces of chlorin to become 
 absorbed; the application of a suction pump to the rear outlet 
 tube will help to bring about this result. 
 
 Mohr's apparatus, shown in Fig. 50, is of very simple con- 
 struction and easy to use. 
 
 The distilling flask is fitted with a paraffin-soaked cork, 
 through which a delivery tube containing one bulb passes; 
 this delivery tube again passes through a common cork which 
 loosely fits a stout, large test tube, containing the potassium 
 iodid solution. The delivery tube is drawn out to a fine point 
 
22S THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 and reaches to near the bottom of the test tube. The latter 
 is placed, when in operation, in a hydrometer jar containing 
 cold water. 
 
 Estimation of Manganese Dioxid (Mn02 = 86.93). o-4 
 gm. of pulverized manganese dioxid is placed into the distil- 
 
 FiG. 50. 
 
 ling flask (a. Fig. 48 or Fig. 49) and the latter filled to two- 
 thirds of its capacity with concentrated hydrochloric acid and 
 connected without delay with the vessel containing the potas- 
 sium iodid solution. The flask is then gradually heated so 
 that a steady current of chlorin passes over into the potassium 
 iodid solution. When the evolution of chlorin gas begins to 
 diminish, the heat is slowly raised to boiling, and continued 
 
ANALYSIS BY OXIDATION AND REDUCTION 229 
 
 at this point until about one-third of the acid liquid has dis- 
 tilled over. The delivery tube is then removed and rinsed, as 
 previously described, and the liberated iodin titrated by means 
 
 N . . 
 
 of — sodium thiosulphate solution, of which we will assume 
 
 60 cc. were consumed. 
 
 The reactions are as follows: 
 
 (a) Mn02-l-4HCl = MnCl2-l-2H20-t-Cl2. 
 
 86.93 70.92 
 
 (J)) Cl2 + 2KI = 2KCl+l2. 
 
 70.92 253.84 
 
 (c) I2 -I- 2Na2S203-5H20 = sNal + Na2S406 + 10H2O. 
 2)253-84 2 )496.96 
 
 10)126.92 10)248.48 
 
 12.692 gms. 24.848 gms. = iooo cc. — V.S. 
 
 N 
 0.012692 gm. =1 cc. — V.S. 
 
 These equations show that 494.96 gms. of sodium thiosul- 
 phate will decolorize 253.84 gms. of iodin, which quantity is 
 liberated by 70.92 gms. of chlorin, which is itself liberated 
 from hydrochloric acid by 86.93 g™^- °^ manganese dioxid. 
 
 Therefore i cc. of a decinormal solution of sodium thio- 
 sulphate (containing 24.848 gms. in looo cc.) is equivalent 
 to 0.012692 gm. of I; 0.003546 gm. of CI; 0.0043465 gm. of 
 Mn02; 0.0008 gm. of O (available). 
 
 The 60 cc. of the thiosulphate solution used in this assay 
 will therefore represent 0.0043465X60 = 0.2608 gm. of pure 
 Mn02, or 65.2 per cent. 
 
 0.2608X100 
 
 =65.2 per cent. 
 
 0.4 ^ ^ 
 
 This is the method which should be used for the assay of 
 native manganese dio.xid. The freshly precipitated manganese 
 
230 THE ESSENTIALS OF VOLUISIETRIC ANALYSIS 
 
 dioxid of the Pharmacopoeia may be assayed by the more 
 easily performed digestion method described on page 233. 
 
 Estimation of Chromic Acid and Chromates. Chromic 
 anhydrid, chromium trioxid (Cr03=ioo), when heated with 
 concentrated hydrochloric acid, liberates chlorin as per the 
 equation, 
 
 CrOa + 6HC1 = CrCls + 3H2O + CI3. 
 
 100 3X35.46 
 
 One hundred parts of CrOs liberates 3X35.46 parts of CI, 
 hence one atomic weight of chlorin, 35.46 parts, represents 33.3 
 
 N 
 + parts of CrOs. Or i cc. of — sodium thiosulphate repre- 
 sents 0.00333 g™- of CrOs- 
 
 Estimation of Potassium Bichromate (K2Cr207 = 294.20). 
 This salt, as explained in a previous chapter, has three atoms 
 of oxygen available for oxidation. A molecule of this salt is 
 therefore equivalent to six atoms of chlorin, and when boiled 
 with hydrochloric acid will liberate six atoms of chlorin, as 
 the equation shows. 
 
 KaCraOy + 14HCI = 2KCI + aCrClg + 7H2O + 3CI2. 
 294.2 6X35.46 
 
 Thus one atom of liberated chlorin will represent one-sixth 
 
 of 294.2, which is 49.03 + parts of potassium dichromate. Then 
 
 N . . 
 
 I cc. of — sodium thiosulphate will represent 0.004903 gm. 
 
 of K2Cr207. In the same way all other chromates may be 
 treated, but these compounds will, when treated with hydro- 
 chloric acid, liberate chlorin at once and without the applica- 
 tion of much heat, hence some chlorin is apt to be lost before 
 the distillation flask can be connected with the apparatus, 
 and therefore it is more convenient to employ the digestion 
 method later described. 
 
ANALYSIS BY OXIDATION AND REDUCTION 231 
 
 The reaction in the case of neutral potassium chromate is 
 as follows: 
 
 KaCrOi + 8HC1 = 2KCI +4H2O + CrCls + CI3. 
 
 194-- 
 
 Lead peroxid, Pb02; cobaltic oxid, C02O3; nickel oxid, 
 Ni203, as well as many other substances, may be assayed by 
 this distillation method. 
 
 Estimation of Alkali lodids by the Distillation Method. 
 This method is based upon the fact that metallic iodids, when 
 
 Fig. 51. 
 
 treated with ferric salts in acidulated solution, yield up all of 
 their iodm. As shown in the equation 
 
 Fe2(S04)3 + 2KI = K2S04 + 2FeS04 + l2. 
 
 The iodin thus set free is distilled into a solution of potassium 
 iodid and its quantity determined by titration with sodium 
 thiosulphate in the usual manner. 
 
 For the reaction, ferric sulphate or ammonio-ferric alum 
 may be used; the latter is, however, preferred because of its 
 paler color. Ferric sulphate and ferric chlorid are so dark 
 in color that the determination of the end-reaction is quite a 
 
232 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 difficult matter; furthermore, these salts frequently contain 
 traces of nitrates which, if present, liberate chlorin from the 
 chlorids or distil over and liberate iodin from the potassium 
 iodid solution in the recei\'ing vessel. Ferric chlorid is par- 
 ticularly objectionable because of the tenacity with which it 
 holds the last portions of iodin. 
 
 The distillation may be done in the Fresenius apparatus, 
 Fig. 49, or better in that shown in Fig. 51. The latter consists 
 of a loo-cc. distilling flask {a), connected by means of a glass 
 tube with a nitrogen flask (6), which contains a 10 per cent 
 potassium iodid solution, and which is kept in a vessel of ice 
 water when in use. The stoppers used are cork, well soaked 
 in paraffin. The construction of the flask is particularly suit- 
 able, because it presents a large surface to the vapor of iodin 
 which distils over, and because the titration can be done directly 
 in it, thus avoiding the necessity of transferring its contents 
 to a beaker or other vessel. 
 
 The glass tube which conveys the iodin vapor must not be 
 carried into the solution of potassium iodid and must not be 
 drawn to a fine point. The reason for this is that the iodin 
 condensing at the point would soon choke up the tube and 
 prevent the further passage of iodin vapor. Any iodin which 
 condenses in the tube is washed down into the potassium iodid 
 solution by the steam during the distillation. 
 
 The Process. Into the flask (a) is introduced about 5 gms. 
 of ammonio-ferric alum, 50 cc. of water, 20 cc. of diluted 
 sulphuric acid (i : 10), and the iodid to be examined, accurately 
 weighed. Take about 0.5 gm. The flask is then connected 
 with the receiving vessel Qf), which is about half filled with a 
 10 per cent potassium iodid solution, and after connections 
 are made tight, heat is gradually applied to the distilling flask. 
 After most of the iodin has passed over, the heat is raised to 
 boiling, and continued at this temperature until about one- 
 
ANALYSIS BY OXIDATION AND REDUCTION 233 
 
 fourth of the hquid has passed over, and the solution in the 
 distilUng flask is no longer of a brown color. 
 
 When the receiving vessel has sufficiently cooled, it is dis- 
 connected and its contents titrated with decinormal sodium 
 thiosulphate, using starch as indicator. Before beginning 
 titration, however, it is necessary to rinse the lower extremity 
 of the tube and the stopper into the solution in the receiving 
 vessel, in order that every trace of iodin be collected. 
 
 The calculation is then made as follows • 
 
 (a) Fe2(S04)3.(NH4)2S04 + 2KI 
 
 332.04 
 
 = K2SO4 + (NH4) 2SO4 + 2FeS04 + 12. 
 
 253-84 
 
 (b) I2 + 2Na2S203 + SH2O = 2NaI + Na2S406 + 10H2O. 
 2 )253-84 
 
 12.692 gms. =1000 cc. — V.S. 
 0.012092 gm.= I cc. — V.b. 
 
 Referring to the above equations we see that 253.84 gms. 
 of iodin are liberated from 332.04 gms. of potassium iodid; 
 thus 126.92 gms. of iodin represent 166.02 gms. of potassium 
 iodid, and therefore i cc. of decinormal sodium thiosulphate 
 solution, representing 0.012692 gm. of iodin, will at the same 
 time represent 0.016602 gm. of KI. 
 
 N 
 If in the above assay 29 cc. of — sodium thiosulphate are 
 
 consumed, we multiply the factor for KI = 0.016602 gm. by 29, 
 this gives 0.4814+ gm., the quantity of pure KI in the 0.5 gm. 
 taken, which is about 96 per cent. 
 
 Digestion Methods. The distillation methods above 
 described may be avoided in many cases and the more easily 
 performed digestion process used. For instance, freshly pre- 
 
234 THE ESSENTIALS OF \'OLUMETRIC ANALYSIS 
 
 cipitated manganese dioxid, lead peroxid, chromic acid, chlor- 
 ates, bromates, iodates, ferric salts, and a great many other 
 substances, may be assayed by mere digestion with hydro- 
 chloric acid at a slightly elevated temperature. 
 
 The digestion is performed in a strong glass bottle, provided 
 vi^ith an accurately fitting ground-glass stopper which is tied 
 down by means of wire or secured by a 
 clamp. See Fig. 52. 
 
 Before using the bottle for this opera- 
 tion it should be tested by securely tying 
 down the stopper and immersing the 
 bottle entirely in hot water to see if the 
 stopper fits sufJGiciently tight. If it does 
 not, bubbles of air will escape from 
 inside, and the bottle is useless for the 
 purpose intended. In that event the 
 stopper must be reground into the neck 
 of the bottle with a little very fine emery 
 and water. The capacity of the bottle may vary from 50 to 
 150 cc. 
 
 The Process. The substance is accurately weighed and 
 introduced into the bottle together with a small quantity of 
 coarsely powdered glass or small pure flint pebbles (to prevent 
 caking, especially in the case of insoluble powders). A suffi- 
 cient excess of potassium iodid solution is then added, followed 
 by some pure concentrated hydrochloric acid. The stopper 
 is then quickly inserted, firmly secured by wire or a clamp, 
 and the bottle placed in a water bath, and the water gradually 
 heated to boiling; this temperature being continued until 
 decomposition is complete, which is usually in about half an 
 hour. The bottle is then allowed to cool slowly and its 
 contents emptied into a beaker. Then, after washing the 
 bottle and adding the washings to the contents of the beaker, 
 
 Fig. 52. 
 
ANALYSIS BY OXIDATION AND REDUCTION 235 
 
 the liberated iodin is estimated by titration with sodium thio- 
 sulphate. 
 
 The potassium iodid used in this process must be abso- 
 lutely free from iodate. 
 
 N 
 One cc. of — sodium thiosulphate is equivalent to 
 
 KCIO3 0.002042 gm. 
 
 NaClOs 0.001784 " 
 
 KBrOs 0.0027836 " 
 
 KIO3 0-003567 " 
 
 Mn02 0.0043465 " 
 
 Pb02 0.011955 " 
 
 CrOs 0.00333 " 
 
 Estimation of Chlorates, Bromates and lodates. The 
 
 estimation of these salts is based upon the fact that in each 
 case one equivalent of the acid or its monobasic salt liberates 
 six equivalents of chlorin and consequently six equivalents 
 of iodin when decomposed by the digestion method. 
 This is illustrated by the equations: 
 
 (a) KCIO3 + 6HCU3H2O+KCI + CI6; 
 
 (b) KBr03+6HCl = 3H20 + KBr + Cl6; 
 
 (c) KI03+6HCI = 3H20 + KH-Cl6. 
 and 
 
 (d) Cle + 6KI = 6KCl + l6. 
 
 In the distillation process, however, bromates and iodates 
 liberate only four equivalents of iodin, while bromous chlorid 
 and iodous chlorid remain in the retort, therefore in these 
 cases the digestion is preferable to the distillation method. 
 
 If the bromate or iodate to be assayed contains any bromid 
 
236 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 or iodid, bromin or iodin respectively will be liberated upon 
 the addition of the acid, according to the equations 
 
 (a) sKBr + KBrOs + 6CH1 = 6KaCl + 3H2O + Brg; 
 (6) 5KI + KI03+6CHl=-6KaCl+3H20+l6; 
 
 therefore the method is not applicable for the assay of such 
 mixtures. 
 
 The presence of either of these salts may be ascertained 
 by moistening a small quantity of the salt mth dilute sulphuric 
 acid, when if a yellow or brown coloration results, either a 
 bromid or an iodid respectively is present. 
 
 Example. Estimation of Potassium Chlorate. 0.2 gm. of 
 the salt is introduced into the digestion bottle of about loo-cc. 
 capacity, 10 cc. of water added, and about 4 gms. of potas- 
 sium iodid (or sufficient of its saturated solution). This is 
 followed by 10 cc. of concentrated hydrochloric acid, the 
 stopper quickly inserted, firmly secured by wiring or a clamp, 
 and the flask placed, stopper downward, in a water bath. 
 The water is gradually raised to boiling and kept at this 
 temperature for about half an hour. It is then allowed to 
 cool slowly, and the contents of the bottle washed into a 
 beaker and titrated with decinormal sodium thiosulphate, 
 using starch as indicator. The number of cubic centimeters 
 
 N 
 of — thiosulphate solution used, multiplied by 0.002042 gm., 
 
 gives the weight of pure KCIO3 present in the sample. 
 
 In the assay of bromates and iodates a smaller quantity 
 of hydrochloric acid may be used, and a lower temperature, 
 say 50° C, is sufficient for decomposition. 
 
 Example. Estimation of Potassium Br ornate. 0.2 gm. of 
 the salt is dissolved in 15 cc. of water, 4 gms. of potassium 
 iodid are added, followed by 4 cc. of concentrated hydro- 
 
ANALYSIS BY OXIDATION AND REDUCTION 237 
 
 chloric acid. The bottle is securely closed, as in the fore- 
 going assay, and heated for half an hour at 50° C. Then, 
 
 N 
 after decomposition, titration with — sodium thiosulphate 
 
 solution is begun, each cc. of which represents 0.2836 gm. of 
 
 pure KBrOs. 
 
 If iVo of one sixth of the molecular weight of either salt 
 
 N 
 be taken for assay, each cc. of the — thiosulphate solution 
 
 used will indicate i per cent purity. 
 
 Estimation of Ferric Salts. 'WTien a ferric salt in an 
 acidulated solution is digested with an excess of potassium 
 iodid_the salt is reduced to the ferrous state, and iodin is set free. 
 
 Fe2Cl6 + 2KI=2FeCl2 + 2KCl+l2. 
 
 One atom of iodin is liberated for each atom of iron in 
 
 the ferric state. The liberated iodin is then determined by 
 
 sodium thiosulphate in the usual way. 126.92 gms. of iodin 
 
 = 55.82 gms. of metallic iron. 
 
 The method is exemplified in the following assay: 
 
 Ferric Chlorid (Fe2Cl6 or FeCls). 0.5 of the dry salt 
 
 accurately weighed is put into a loo-cc. glass-stoppered bottle, 
 
 15 cc. of water are added to dissolve the salt, and then 3 cc. 
 
 of concentrated hydrochloric acid and 2 gms. of pure potassium 
 
 iodid are introduced. The bottle is securely closed and its 
 
 contents heated to 40° C. (104° F.) for half an hour. A 
 
 higher temperature must be avoided, otherwise iodin will be 
 
 volatilized. The liquid is then cooled and the liberated iodin 
 
 N . 
 
 titrated with — sodium thiosulphate until the color is just 
 
 discharged. Starch mav be used as indicator. 
 
 N 
 Each cc. of — thiosulphate solution used represents 0.005582 
 
 gm. of metallic iron or 0.01622 gm. of pure ferric chlorid. 
 
238 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The following equations illustrate the reactions: 
 
 FeaCle + 2KI = 2FeCl2 + 2KCI + I2. 
 2)324-4 2)253-84 
 
 162.2 126.92 
 
 Then 
 
 I2 + 2Na2S203.sH20 =2NaI + Na2S406 + 10H2O. 
 2)253-84 
 10 )126.92 j^ 
 
 12.692 gms. = 1000 cc. — V.S. 
 
 The assay of other ferric salts is practically the same as 
 
 that described above. 
 
 N . . 
 
 Each cc. of — sodium thiosulphate is equivalent to 
 
 Ferrum, Fe 0.005582 gm. 
 
 Ferric ammonium sulphate,Fe2 (SO4) 3. (NH4) 2SO4 o. 0266 
 
 Ferric chlorid, Fe2Cl6 0.01622 
 
 " (cryst.), Fe2Cl6 + i2H20 0.02702 
 
 " nitrate, Fe2 (NO 3) 6 0.024T85 
 
 ". oxid, Fe20a 0.007982 
 
 " sulphate, Fe2(S04)3 0.019992 
 
 In the case of the scale salts of iron, which are mostly of 
 indefinite and variable composition, it is the quantity of metallic 
 iron present which is determined. 
 
 The solutions of ferric salts are estimated in the same man- 
 ner. It is the rule to take i to 2 gms. of the solution for assay. 
 
 Reduction Methods Involving the Use of Standard Arsenous 
 Acid Solution (Chlorometry) 
 
 As previously described, arsenous oxid when brought in 
 contact with iodin in an alkaline solution, results in an oxida- 
 tion of the former to arsenic oxid, and a conversion of the 
 iodin to hydriodic acid, as shown in the equation 
 
 As203 + 2H20-H4=As205+4HL 
 
ANALYSIS BY OXIDATION AND REDUCTION 239 
 
 Advantage is taken of this reaction for tlie estimation, not 
 only of arsenous and antimonous compounds, but also of iodin 
 and the other halogens, chlorin and bromin, as well as of all 
 those bodies which, when heated with hydrochloric acid, evolve 
 chlorin, as for instance the peroxids. 
 
 The reaction with chlorin is as follows: 
 
 Cl4+2H20+As203 = 4HCl+As205. 
 
 This reaction is really an oxidation, so far as the formation 
 of arsenic oxid (AS2O5) is concerned, but there is no accom- 
 panying reduction. The conversion of the halogen to an 
 haloid acid is not strictly 'a reduction in the accepted sense 
 of the word. Nevertheless, for obvious reasons, we speak 
 of analyses done by means of arsenous acid as reduction 
 methods. 
 
 The chief value of this method is foimd in the estimation 
 of free chlorin, as in chlorin water, and the available chlorin 
 existing in hypochlorites or that evolved from hydrochloric 
 acid by heating with peroxids. Hence the designation " chlor- 
 ometry." 
 
 In carrying out this method free alkali must be present 
 to combine with the haloid acid which is formed. The alkali 
 must be in the form of bicarbonate. Normal carbonates or 
 hydroxids are not suitable (see page 193). 
 
 The solutions required are: 
 
 Decinormal iodin (see page 188); 
 
 Decinormal arsenous acid; 
 
 Starch solution (see page 192) ; or iodized starch test paper. 
 
 N 
 Preparation of Decinormal — Arsenous Acid (As203 = 
 
 107 02- — V.S. = 4.948 gms. in i liter.) 4.948 gms. of the 
 purest sublimed arsenous anhydrid (As-^Os) are dissolved in 
 
240 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 about 250 cc. of distilled water with the aid of about 20 gms. 
 of pure sodium bicarbonate. The anhydrid should be in fine 
 powder, and the mixture heated to effect complete solution. 
 This is then diluted with some water, cooled, and made up 
 to 1000 cc. It is then standardized with decinormal iodin, 
 using starch as indicator. Decinormal arsenous acid solution 
 should correspond, volume for volume, with decinormal iodin 
 solution. 
 
 If this solution is made from pure arsenous acid, it will 
 hold its titer for years, but if any sulphur is present there 
 will be an absorption of oxygen from the air and a consequent 
 oxidation to arsenic oxid. If the presence of sulphur is sus- 
 pected, the solution should be tested with silver nitrate, when 
 its presence will be indicated by the formation of a reddish 
 precipitate. 
 
 Iodized Starch Test Paper. A portion of starch solution 
 is m.ixed with a few drops of potassium iodid solution and 
 in this are soaked strips of pure white filtering paper. This 
 test paper is used in the damp state; it is then far more 
 sensitive. 
 
 Estimation of Free Halogens. The estimation of chlorin, 
 bromin or iodin by the chlorometric method depends, as 
 before stated, upon their power of oxidizing arsenous acid. 
 WTien a free halogen is brought in contact in alkaline solution 
 with arsenous acid, the latter is oxidized to arsenic acid, while 
 the halogen is transformed into a haloid acid, as per equations 
 
 CI4 ] f 4HCI 
 
 Br4 I- +2H20+As203 = As20.5+ j 4HBr. 
 
 I4 J i 4HI 
 
 The estimation may be carried out in two ways: ist, by 
 direct titration with a standard arsenous oxid solution, using 
 iodized starch test paper as indicator; 2d, by residual titra- 
 
ANALYSIS BY OXIDATION AND RICDUCTION 241 
 
 tion, an excess of the standard arsenous oxid being taken, and 
 then retitrating with standard iodin solution, using starch as 
 indicator. The residual titration method need not be employed 
 for free iodin, as this can be titrated direct with the arsenous 
 oxid solution, using starch as indicator. Furthermore, iodin 
 need not be brought into solution to be titrated by this method. 
 
 The estimation of free halogens by the direct chlorometric 
 method is as follows: 
 
 An acciurately weighed quantity of substance made alka- 
 line by the addition of sodium bicarbonate is titrated with 
 decinormal arsenous acid solution, and from time to time 
 during the titration a drop of the solution is removed on the 
 end of a pointed glass rod and brought in contact with a piece 
 of iodized starch test paper. So long as free chlorin or bromin 
 is present the liquid will cause a blue stain on the test paper, 
 but when the halogen is all taken up no blue color is produced. 
 
 If the exact point is overstepped the residual method must 
 be used. A little additional excess of the arsenous acid solu- 
 tion may be added, together with a few drops of starch solu- 
 tion, and the excess then titrated by means of decinormal 
 iodin solution imtil the blue color is produced. The volume 
 of decinormal iodin solution so used, deducted from the total 
 volume of arsenous acid solution taken, gives the exact quan- 
 tity which was oxidized by the halogen, and from this the 
 percentage of chlorin or bromin may be calculated. 
 
 Example i. Estimation of Chlorin in Chlorin Water. 
 Twenty cc. of chlorin water (sp.gr. i.o) titrated by the direct 
 method required 22 cc. of decinormal arsenous acid solution 
 before the iodized starch test paper indicated the completion 
 of the reaction. 
 
 By referring to the equation we see that each cc. of the 
 arsenous acid solution represents 0.003546 gm. of chlorin. 
 Therefore, if 22 cc. were used, the 20 cc. of chlorin water 
 
242 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 must have contained 22X0.003546 gm. of chlorin, which is 
 0.078012. 
 
 The 20 cc. of chlorin water (sp.gr. i.o) weigh 20 gms. Hence 
 
 0.078012 X 100 
 
 20 
 
 = 0.39+ per cent. 
 
 Example 2. The 20 cc. of chlorin water weighing 20 gms. 
 
 were treated with 26 cc. of the arsenous acid solution, starch 
 
 solution was then added, and the excess of arsenous acid 
 
 solution titrated by means of decinormal iodin. Four cc. of 
 
 the latter were required, then 4 from 26 cc. leaves 22 cc, the 
 
 N 
 quantity of the — arsenous acid solution which reacted with 
 
 the chlorin. The calculation is the same as in Example i. 
 The reaction is as follows: 
 
 CI4 + 2H2O + AS2O3 = 4HCI + AS2O5. 
 
 4)141-84 4)197-9^ 
 
 i° )35-46 i° )49-48 j^ 
 
 3.546 gms.= 4-948 gms. = 1000 cc. — V.S. 
 
 N 
 0.003546 gms. = I cc. — V.S. 
 
 Estimation of Available Chlorin in Bleaching Powder. 
 
 Three and one-half gms. of the bleaching powder (chlorin- 
 ated lime) are triturated thoroughly with 50 cc. of water, 
 and the mixture transferred to a graduated vessel, together 
 with the rinsings, and made up to 1000 cc. with water. 
 This is thoroughly shaken. 100 cc. of it (representing 0.35 
 gm. of the sample) -is removed by means of a pipette and 
 titrated with decinormal arsenous acid solution, as described 
 in the foregoing assay, using either the iodized starch test 
 
 paper as mdicator or retitratmg the excess of — arsenous acid 
 
 N-. 
 solution added by means of — iodin solution. 
 
 ■' 10 
 
 10 
 
ANALYSIS BY OXIDATION AND REDUCTION 243 
 
 N 
 Each cc. of — AS2O3 V.S. represents 0.003546 gm. of 
 
 available chlorin. 
 
 Ca(OCl)Cl + AS2O3 = AS2O5 + 2CaCl2. 
 4 )141-84 4 )197-92 
 
 10 )35-46 10 )49-48 j^ 
 
 3.546 gms.= 4-948 gms. or 1000 cc. — V.S. 
 
 As seen by referring to the above equation this process 
 determines the value of the chlorinated lime by measuring 
 the amount of arsenous acid which the oxygen present in 
 the active constituent (Ca(OCl)Cl) is capable of oxidizing. 
 In the formula of this compound there are two atoms of 
 chlorin and one atom of oxygen. Therefore the quantity of 
 bleaching powder which yields 35.46 parts of available chlorin 
 will also supply 8 parts of oxygen; this may therefore be 
 taken as the measure of the chlorin. The same method may 
 be employed for the assay of all other solutions containing 
 available chlorin. 
 
 Assay of Manganese Dioxid (Chlorometric). The chloro- 
 metric assay of manganese dioxid, as well as that of all other 
 bodies which liberate chlorin when heated with hydrochloric 
 acid, may be made in similar manner to that described for 
 the iodometric assays of these substances, the same apparatus, 
 etc., being used. 
 
 N 
 The liberated chlorin is, however, titrated with — arsenous 
 
 ID 
 
 acid solution. 
 
 The chlorin may be distilled into a solution of sodium 
 
 N 
 carbonate, and this solution is then titrated with — arsenous 
 
 acid or the chlorin may be distilled directly into a measured 
 
 N 
 volume of — arsenous acid solution and the latter then titrated 
 
 ID 
 
244 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 with — iodin solution, using starch as indicator, the difference 
 
 10 
 
 between the volume of iodin solution used and that of the 
 arsenous acid solution taken is the measure of the latter which 
 reacted with the chlorin. 
 
 It is a good plan in each case to divide the solution into 
 two or three equal parts and to titrate each separately. 
 
 Reduction Methods Involvng the Use of Stannous Chlorid 
 
 Stannous chlorid (SnCl2) is a very powerful reducing 
 agent. Its action in this respect depends npon its aflBnity 
 for chlorin which it readily abstracts from most other chlorids. 
 In it's action upon mercuric chlorid a portion of the latter 
 is always reduced to the metallic state. 
 
 This reducing action of stannous chlorid is utilized in 
 certain volumetric processes, especially in the estimation of 
 iron. In this case it possesses an advantage over perman- 
 ganate, in that the iron must be in the ferric state, in which 
 condition it is most usually found, while if permanganate 
 is used, a preliminary reduction to the ferrous state is necessary 
 before titrating. The great disadvantage, however, is in the 
 fact that even short contact with air \vill quickly oxidize it, 
 and thus spoil its titer. In consequence of this it must be 
 frequently tested, and can be used only in the form of empirical 
 solutions. 
 
 It is particularly useful in the titration of ferric salts, 
 which salts can be accurately estimated by direct titration 
 with it, the end-point being recognized by the disappearance 
 of the yellow color of the ferric solution. These salts may 
 also be estimated residually by adding an excess of stannous 
 chlorid solution of known strength and retitrating the excess 
 by means of standard iodin, using starch as indicator. 
 
ANALYSIS BY OXIDATION AND REDUCTION 
 
 24,- 
 
 The reactions are: 
 
 FeaCle + SnCl, = 2FeCl2 + SnCLi 
 
 and 
 
 SnCl, + 2HCl + I, = SnCl4 + 2HI. 
 
 The Estimation of Iron by Means of Stannous Chlorid 
 Solutions may be accurately affected by the following pro- 
 cedure, as suggested by Fresenius. The solutions necessary 
 are: 
 
 (a) A solution of ferric chlorid containing lo gms. of pure 
 iron in a liter. 
 
 This is made by dissolving 10.04 gnis. of thin annealed 
 binding vme (which contains 99.6 
 per cent of pure iron) in a sufficient 
 quantity of pure hydrochloric acid. 
 A small quantity of potassiiun 
 chlorate is then added to effect 
 complete oxidation of the iron, and 
 the excess of chlorin expelled by 
 boiling. This solution is then 
 cooled and diluted to i liter. 
 
 {b) A solution of stannous chlo- 
 rid made by dissolving about 10 
 gms. of pure tin in 200 cc. of 
 strong, pure hydrochloric acid. 
 This may be done by heating the 
 tin in small pieces with the acid in 
 a flask, and introducing a few 
 pieces of platinum foil to excite 
 galvanic action. The solution so 
 obtained is diluted to about one liter with distilled water and 
 should be preserved in a bottle, such as shown in Fig. 53, 
 to which air can only gain access through a strongly alkaline 
 
 Fro- 53- 
 
246 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 solution of pyrogallic acid. WTien so kept the strength of 
 the solution can be preserved for several weeks. 
 
 (c) A solution of iodin in potassium iodid. This may be 
 approximately or exactly decinormal. 
 
 The procedure is as follows: 
 
 ist. The relation between the tin solution and the iodin 
 solution is found. 
 
 2d. The relation between the tin solution and the iron 
 solution is determined. 
 
 3d. The assay. 
 
 The relation between the tin solution and the iodin solu- 
 tion is found as follows: 
 
 Two cubic centimeters of the tin solution are put into a 
 beaker, a little starch solution added, and the iodin solution 
 then delivered in from a burette until the blue color occurs. 
 If 4 cc. are used, then each 2 cc. of iodin solution represents 
 I cc. of tin solution. 
 
 The relation between the tin solution and the iron solution 
 is found as follows: 
 
 Fifty cvibic centimeters of the iron solution (representing 
 0.5 gm. of iron) are put into a small flask together with a little 
 hydrochloric acid and heated to gentle boiling. The tin solu- 
 tion is then delivered from a burette until the yellow color 
 of the iron solution is nearly discharged. It is then added 
 continuously, drop by drop, until the color is entirely gone. 
 Assuming that 35 cc. were required, then each 35 cc. of tin 
 solution are equivalent to 0.5 gm. of pure iron. If the end- 
 point is not clearly recognized and an excess of the tin solution 
 was added, the solution should be quickly cooled, a few drops 
 of starch solution added, and the excess estimated by titrating 
 with the iodin solution, each cc. of which represents 0.5 cc. 
 of the tin solution. The excess so found, deducted from the 
 
ANALYSIS BY OXIDATION AND REDUCTION 247 
 
 total quantity of tin solution added, gives the quantity of the 
 latter, which corresponds to 0.5 gm. of iron. 
 
 Having determined these data, the analyst can readily 
 estimate any unknown quantity of iron in solution in the ferric 
 state. 
 
 If the iron is partly or wholly in the ferrous state it may 
 be oxidized by adding some potassium chlorate and boiling 
 to expel excess of chlorin. 
 
 The Assay. A solution of iron taken for analysis, required 
 24 cc. of the tin solution. The quantity of iron present is 
 calculated by proportion as follows: 
 
 35 cc. : 0.5 gm. : : 24 cc. : x; -^ = 0.34 gm. 
 
 To secure accurate results the iron solution assayed must 
 be fairly concentrated, because then the end-reaction is more 
 readily seen, and also because the greater the dilution the 
 larger the amount of tin solution will be required. It is good 
 policy to use very little excess of the tin solution, so that only 
 a very small quantity of iodin solution is required. 
 
 Estimation of Mercuric Salts (Laborde). This depends 
 upon the fact that stannous chlorid solution added to a solution 
 of a mercuric salt reduces the latter first to mercurous chlorid 
 (calomel) and finally the calomel to metallic mercury. The 
 reduction to calomel results in the formation of a white pre- 
 cipitate, and when the mercuric salt is completely reduced 
 the stannous chlorid acts upon and reduces the calomel to 
 metallic mercury, which results in the production of a char- 
 acteristic brownish color. 
 
 The reactions are as follows: 
 
 SnCla + 2HgCl2 = SnCl 1 + 2HgCl ; 
 SnCla + 2HgCl = SnCl.i + Hg.. 
 
248 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 According to Laborde the tin solution is made by dissolving 
 8 gms. of pure tinfoil by means of heat in loo cc. of pure 
 hydrochloric acid, and diluting to 2 liters. 
 
 This tin solution is checked against a solution of mercuric 
 chlorid containing 10 gms. per liter. To counteract the hin- 
 dering effect of the hydrochloric acid the solution under analysis, 
 containing o.i gm. of mercuric chlorid, is mixed with 5 cc. 
 of a solution containing 100 gms. of ammonium acetate and 
 100 cc. of acetic acid to the liter. The acetic acid promotes 
 the disappearance of the brown color which occurs at the 
 point where the tin solution is in excess, but before reduction 
 is complete. The titration with the tin solution is continued 
 until a permanent brown color occurs. 
 
 If the brown color is too dark from overstepping of the 
 end-point, the addition of i cc. of the mercuric chlorid solution 
 will render the solution white again, and the titration can 
 then be carried further. 
 
 This method, which is convenient, rapid, and very accurate, 
 can be employed in many cases. If the mercuric solution 
 contains any free mineral acids, the latter must be neutralized 
 with ammonia (in the presence of ammonium acetate, to 
 prevent formation of ammoniated mercury) . 
 
 The presence of alkali and alkali earth salts or most salts 
 of other metals (except iron, gold, and platinum) do not 
 in the least interfere vnth the accuracy of the results. The 
 same is true of organic acids, either free or in combinatiop 
 with alkali. 
 
PART II 
 
 CHAPTER XI 
 ESTIMATION OF ALKALOIDS 
 
 In malung alkaloidal assays of drugs it has long been the 
 custom to evaporate the final ethereal or chloroformic extract, 
 and to v/eigh the residue as alkaloid. This residue seldom, 
 if ever, consists of the pure alkaloid, and the impurities — i.e., 
 non-alkaloidal matter — is variable in amount and difficult to 
 entirely remove, consequently gravimetric results are in many 
 cases very wide of the truth, and hence unreliable. 
 
 The volumetric methods are in most cases much more 
 satisfactory. 
 
 While the results of the titration of the total alkaloids of 
 drugs cannot be called absolutely accurate, nevertheless experi- 
 ence has shown that they are nearer the truth than those 
 obtained by the gravimetric method. 
 
 In estimating an alkaloid by titration it is essential to 
 know the formula and molecular weight of the alkaloid, as 
 well as the equivalent of acid with which it will combine. 
 
 The quantity of alkaloid present is easily calculated when 
 we know that a molecular weight of a monobasic acid or half 
 a molecular weight of a dibasic acid will combine with and 
 neutralize a molecular weight of an alkaloid, provided the 
 alkaloid is a monacid base. If the alkaloid is a diacid base, 
 one molecular weight will combine with two molecules of a 
 monobasic acid or one molecular weight of a dibasic acid. 
 
 249 
 
250 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Sparteine and emetine (?) are diacid alkaloids; most of 
 the others are monacid bases. 
 Examples. Monacid alkaloids: 
 
 C20H24N2O2 + HCl = C20H24N2O2 . HCl ; 
 
 Quinine 
 
 (Ci7Hi9N03)2 +H2S04= (Ci7Hi9N03)2.H2S04; 
 
 Morphine 
 
 (C21H22N2O2) 2 + H2SO4 = (C21H22N2O2) 2 . H2SO4 ; 
 
 Strychnine 
 
 Diacid alkaloids: 
 
 C15H26N2 + 2PICI = CisH26N2(HCl)2; 
 
 Sparteine 
 
 C30H40N2O5+H2SO4=C30H40N2O5.H2SO4. 
 
 Emetine (Kunz) 
 
 The quantity of alkaloid present in the substance is easily 
 calculated, as illustrated by this equation: 
 
 C20H24N2O2 + HCl = C20H24N2O2 . HCl. 
 
 Quinine 
 
 N 
 324.2 gms. 36.46 gms. = 1000 cc. — V.S.; 
 
 N 
 32.42 " 3.646 " = 1000 cc. — V.S. 
 
 N 
 Thus I cc. of — V.S. =0.03242 gm. of qumme. 
 
 C16H26N2 + 2HCI = Ci5H26N2(HCl)2. 
 
 Sparteine 
 
 2)234-2 2)72.92 j^ 
 
 117. 1 gms. 36.46 gms. = 1000 cc. — V. S.; 
 
 N 
 II. 71 " 3-646 " = 1000 cc— V.S.; 
 
 N 
 I cc. of — V.S. hence = 0.01 171 gm. of sparteme. 
 
 N 
 Thus 1000 cc. of — hydrochloric acid will combine with 
 
ESTIMATION OF ALKALOIDS 251 
 
 iV of the molecular weight in grams of a monacid alkaloid, 
 or tjV of the molecular weight of a diacid alkaloid. 
 
 In the case of drugs where two or more alkaloids are 
 present, accurate results can only be obtained by determining 
 how much of each alkaloid is present by a separate assay. 
 But often it is assumed that the alkaloids are present in equal 
 quantities, and the mean of their molecular weights is taken 
 as the basis for the calculation. 
 
 It must be borne in mind, however, that in titrating alka- 
 loids the greatest care must be exercised and all precautions 
 closely observed in order to attain any degree of accuracy. 
 The volumetric solutions must be prepared with the greatest 
 care and must be absolutely accurate. The eye must be 
 trained (as it can only be through practice), to distinguish 
 the end-colors of the indicators employed, a matter of some 
 difficulty. Furthermore, all measuring instruments used must 
 be accurate, or they should be carefully calibrated in order 
 to find the necessary factor for correction. 
 
 The Volumetric Solutions usually employed in titrating 
 
 N N N N N 
 
 alkaloids are — , — , — , — , and — . The weaker solutions 
 lo 20 25 50 100 
 
 give more accurate results, as will be understood if we remem- 
 
 N N 
 
 ber that a — is 10 X stronger than a — and 100 X stronger 
 I 10 
 
 N N 
 
 than — V.S. Then if — solution be used, one drop may 
 100 I 
 
 overstep the neutral point, while if the same solution were 
 
 N 
 treated with — solution 5 drops would be required to neutralize, 
 
 N 
 which is equivalent to using one half a drop of — solution. A 
 
 — solution will of course be capable of even njore delicate 
 100 
 
252 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 work. In the case just mentioned 45 drops of the — may 
 
 exactly neutralize the solution, hence less than half a drop 
 
 N N 
 
 of — and between 4 and 5 drops of — are represented. 
 
 Therefore in all delicate alkaloidal titrations weak standard 
 
 solutions should be used — in fact, in all titrations where great 
 
 accuracy is required. 
 
 If the alkaloid he from a recent extraction and is in the 
 
 form of a free alkaloid, it is dissolved in a measured volume 
 
 N 
 of — acid solution and the excess of acid solution then deter- 
 10 
 
 N 
 
 mined by residual titration with alkali solution. 
 
 ■' 100 
 
 N 
 In this the — sulphuric acid solution is preferred, except 
 
 N 
 in the case of quinine or cinchonine, in which — hydrochloric 
 
 acid gives better results. 
 
 The process in detail is as follows: Place 2 gms. of the 
 
 N 
 alkaloid into a beaker, add !< cc. of — sulphuric acid solution, 
 
 and warm on a water-bath until the alkaloid is completely 
 
 dissolved. The solution is then allowed to cool and diluted 
 
 to 100 cc. 
 
 Ten cc. of the solution (containing 0.2 gm. of the alkaloid 
 
 N 
 and 7.5 cc. of — sulphuric acid solution) are removed by 
 
 N 
 means of a pipette and retitrated with potassium hydroxid 
 
 N 
 
 solution. One- tenth of the quantity of the alkali used 
 
 100 
 
 N 
 is deducted from the 7.5 cc. of the — acid solution, and the 
 
ESTIMATION OF ALKALOIDS 253 
 
 remainder is the quantity of the latter which combined with 
 and hence represents the alkaloid present. This, if multiplied 
 by the factor, gives the weight of the alkaloid. 
 
 Either haematoxylin solution or Brazil-wood T.S. may be 
 employed as the indicator. 
 
 If the alkaloid is soluble in alcohol, as are quinine and 
 codeine, it may be treated as follows: 
 
 Place 2 gms. of the alkaloid in a graduated cylinder, dis- 
 solve in alcohol, and dilute the solution up to loo cc. with 
 alcohol. Remove lo cc. of this solution (containing 0.2 gm. 
 of the alkaloid) and place in a beaker, add the indicator and 
 run in the decinormal acid solution to slight excess. Rotate 
 the beaker several times, let stand for a few minutes, wash 
 down the sides of the beaker with distilled water, using about 
 
 N 
 
 40 cc, and retitrate the excess of acid with potassium 
 
 ^ 100 ^ 
 
 hydroxid solution, until end-color is given by the indicator. 
 
 N 
 Deduct one-tenth of the quantity of the — solution used 
 
 ^ ■' 100 
 
 N 
 from the quantity of — acid solution added, and the remainder 
 
 is the quantity of the latter, which combined with and hence 
 represents the alkaloid present. 
 
 The indicators best suited for most alkaloids are Hcema- 
 toxylin, Brazil-wood, Cochineal, lodeosin and Litmus. 
 
 The color changes produced by the principal indicators 
 used in alkaloidal titrations are as follows: 
 
254 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 H3ematoxy!in . . 
 Brazil-wooti . . . 
 
 Cochineal , 
 
 Litmus 
 
 lodeosin 
 
 Lacmoid 
 
 Fluorescein 
 
 Congo red 
 
 Melhyl-orange . 
 Phenolphthalein 
 
 Acid. 
 
 Alkali. 
 
 Yellow 
 
 Blue 
 
 (( 
 
 Purplish-red 
 
 Yellowish-red 
 
 Purplish 
 
 Red 
 
 Blue 
 
 Yellow 
 
 Rose-red 
 
 Red 
 
 Blue 
 
 Green fluorescence 
 
 No fluoresc, yellowish 
 
 Blue 
 
 Red 
 
 Red 
 
 Straw-yellow 
 
 Colorless 
 
 Red 
 
 TABLE SHOWING THE FACTOR FOR VARIOUS ALKALOIDS 
 
 N 
 WHEN TITRATING WITH — ACID OR ALKALI 
 
 lo 
 
 Name. 
 
 Formula. 
 
 Molecular 
 Weight. 
 
 Factor. 
 
 Aconilinc 
 
 Atropine 
 
 Brucine 
 
 Cephaeline 
 
 Cinchona alkaloids (combined) , 
 
 Cinchonine 
 
 Cinchonidine 
 
 Cocaine 
 
 Codeine 
 
 Coniine 
 
 Emetine . 
 
 Hydrastine 
 
 Hyoscine , 
 
 Hyoscyamine 
 
 Ipecac alkaloids (combined) , 
 
 Morphine 
 
 Nicotine 
 
 Physostigmine 
 
 Pilocarpine 
 
 Quinine 
 
 Sparteine 
 
 Strychnine 
 
 C3,H„NO,i 
 CijHjjNOs 
 
 C23H29N204 
 CuH.gNOj 
 
 Ci^jjNjO 
 C.sH^jNjO 
 Ci,H2,N04 
 CisH^NOj 
 CjHijN 
 CsoHuNjO, (Glenard) 
 CaoH^oNjOs (Kunz) 
 CisH^iNOj (U. S. P.) 
 CjiHjiNOe 
 C„H„NO, 
 
 C5H,N 
 
 C^HjjNjOj 
 
 Ci.H.eN^O, 
 
 CjoHj^NjOj 
 
 Ci6H2eN2 
 
 C21H22N2O2 
 
 645.48 
 289.24 
 394-28 
 233.20 
 
 294.24 
 294.24 
 303.22 
 299.22 
 127.18 
 496.32 
 
 5°8-:i4 
 247.22 
 
 383-22 
 3°3-i8 
 289.24 
 
 285 . 20 
 81.06 
 
 275 ■ 24 
 208.18 
 324.26 
 234-28 
 334-24 
 
 0.064548 
 
 0.028924 
 
 0.039428 
 
 0.023320 
 
 0.03069 
 
 0.029424 
 
 0.029424 
 
 0.030322 
 
 0.029922 
 
 0.012718 
 
 0.024818 
 
 0.025417 
 
 0.024722 
 
 0.038322 
 
 0.030318 
 
 0.028924 
 
 0.024021 
 
 0.02852 
 
 0.008106 
 
 1.027524 
 
 0.020818 
 
 0.032426 
 
 0.011714 
 
 u. 033424 
 
ESTIMATION OF ALIG\LOIDS 255 
 
 gordin's modified alkalimetric method, using phenol- 
 phthalein as indicator 
 
 The alkaloidal residue obtained by any of the extraction 
 
 N 
 methods in use is dissolved in a measured excess of — hydro- 
 
 20 ■^ 
 
 chloric acid solution. Wagner's or Mayer's reagent is then 
 added, a little at a time, with frequent shaking, until the 
 alkaloids are completely precipitated. The mixture is then 
 diluted with water to zoo cc. and shaken until the double 
 salt of the alkaloid and reagent completely separate. When 
 allowed to stand a few minutes the supernatant liquid 
 should be clear, and if Wagner's reagent has been used, this 
 will be of a dark-red color. The liquid is now filtered, and 
 50 cc. of the filtrate (representing one-half of the alkaloid) 
 is treated with a 10 per cent solution of sodium thiosulphate 
 added, drop by drop, until the color of the free iodin dis- 
 appears. This discolorization is not needed if Mayer's reagent 
 has been used. A few drops of the indicator phenolphthalein 
 are now introduced, and the excess of acid estimated by reti- 
 
 N 
 tration with — potassium hydroxid solution. This, deducted 
 
 from one-half of the volume of the standard acid solution 
 
 N 
 employed, indicates the number of cc. of — hydrochloric 
 
 acid solution, which combined with the alkaloid in 50 cc. 
 
 of the solution. This number, multiplied by two and then 
 
 by the factor for the alkaloid present, gives the total quantity 
 
 of alkaloid. 
 
 Example. An alkaloidal residue consisting of morphine 
 
 N 
 was dissolved in 30 cc. of — HCl V.S. Then Wagner's reagent 
 
 was added in excess as described, and the mixture made qp 
 
256 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 to 100 cc. with water. Fifty cc. of this were filtered off, decolor- 
 
 N 
 ized as directed, and titrated with — KOH V.S. Ten cc. were 
 
 20 
 
 required, which corresponds to 20 cc. for the entire quantity. 
 
 N 
 Then 20 cc. deducted from the ^o cc. of — HCl V.S. added, 
 
 ■^ 20 
 
 leaves 10 cc, the quantity required to neutralize the morphine. 
 
 N 
 The — factor for morphine (0.0137 gm., Gordin), multiplied 
 
 by 10 = 0.137 gm., the quantity of alkaloid in the residue 
 examined. 
 
 N 
 The — factor for morphine here given is somewhat lower 
 
 than the theoretical equivalent. It was ascertained by exper- 
 ment with a sample of the anhydrous alkaloid. 
 
 The following factors (by Gordin) were obtained by com- 
 paring their molecular weights with that of morphine, which 
 factor was determined by experiment: 
 
 N 
 Morphine 0.0137 S^- = i cc. — HCl V.S. 
 
 Hydrastine 0.0184 " = 
 
 Strychnine 0.0160 " = " 
 
 Caffeine, cryst 0.0102 " = " 
 
 Cocaine 0.0146 " = " 
 
 Atropine 0.0139 " = " 
 
CHAPTER XTI 
 
 VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 
 AND THEIR PREPARATIONS 
 
 EXTRACTION OF THE ALKALOIDS 
 
 Selection of the Sample. Care must be taken to secure 
 a fairly representative sample. 
 
 If the drug is in small pieces or consists of seeds or leaves, 
 mix it well, take a portion, pulverize it, and of the powder 
 take a sufhcient quantity for the assay. If the drug be in 
 large lumps, which vary in quality, select a few representative 
 lumps and cut from each a fairly representative section; pul- 
 verize these, mix well, and weigh off a sufficient quantity for 
 the assay. 
 
 If drying is necessary, the loss of weight in drying must 
 be made note of. 
 
 The Exhaustion of the Drug is usually effected by macer- 
 ation in a suitable menstruum, although percolation, boiling, 
 and hot repercolation must be employed in some cases. 
 
 The Choice of Solvent depends upon the nature of the 
 drug. Water dissolves, besides the alkaloids, so much inert 
 matter that the subsequent steps in the assay are liable to 
 be interfered with. Alcohol dissolves too much of the resinous 
 matter, and besides does not penetrate the drug very well. 
 Acidulated water has been much used, but chloroform and 
 ether, separately and in various combinations, are now most 
 generally employed for exhausting drugs in conjunction with 
 
 257 
 
25S THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 alcohol and ammonia. Petroleum benzin has of late been 
 recommended. 
 
 Prollius' fluid, or some modification of it, is very satis- 
 factory. 
 
 Prollius' Fluid consists of ether 325 cc, alcohol 25 cc, 
 and concentrated ammonia water 10 cc. 
 
 Modified Prollius^ Fluid consists of ether 250 cc, chloro- 
 form 80 to ICO cc, alcohol 25 cc, concentrated ammonia 
 water 10 cc 
 
 The Separation of the Alkaloids and the Use of Immiscible 
 Solvents. " The Shaking-Out " Process. Most alkaloids are 
 insoluble in water but are readily soluble in certain " volatile 
 solvents " such as alcohol, ether, chloroform, amyl alcohol, 
 benzene, petroleum benzin, acetic ether, etc., but not in 
 carbon tetrachlorid. Those alkaloidal solvents which do not 
 mix readily with water are known as " the immiscible sol- 
 vents." Salts of alkaloids on the other hand are as a rule 
 soluble in water, but practically insoluble in the immiscible 
 solvents. This property of the alkaloids makes it possible 
 to separate them from their natural sources. 
 
 The alkaloid in solution in an immiscible solvent may be 
 separated by " shaking out " with a dilute acid, the alkaloid 
 is thereby converted into an alkaloidal salt which readily 
 dissolves in the aqueous acid solution. Conversely, an aqueous 
 solution of an alkaloidal salt when treated with an alkali to 
 set free the alkaloid and shaken with an immiscible solvent 
 gives up its alkaloid in a pure state to the latter. The process 
 of assay by this " shaking-out " process is carried out by 
 treating the liquid extracts (after having been freed from 
 alcohol) with an immiscible solvent and a slight excess of 
 an alkali. The alkaloid is thus dissolved in the immiscible 
 solvent, which is then separated from the aqueous liquid, 
 and transferred to another vessel in which it is shaken with an 
 
VOLUMETRIC ASSAYING OF VEGICTABLK DRUGS 259 
 
 excess of acid largely diluted with water. The acid combines 
 with the alkaloid and forms an alkaloidal salt, which leaves 
 the immiscible solvent, and enters the aqueous solution. If 
 the alkaloidal solution so obtained is still colored, the shaking 
 out is repeated until a pure alkaloidal solution in the immis- 
 cible solvent is obtained. The ap- 
 paratus used in this shaking-out 
 operation is known as a separator. 
 See Fig. 54. 
 
 Separators are conical or pear- 
 shaped; the neck is provided with 
 a well-ground glass stopper and the 
 outlet tube or stem at the bottom 
 with an accurately-fitting glass stop- 
 cock. The solvents usually used 
 in alkaloidal drug assaying are 
 alcohol, chloroform, ether and vari- 
 ous mixtures of these containing at 
 least 75 per cent of ether. When 
 chloroform is used it will collect at 
 the bottom of the separator, and can 
 be easily drawn off; but if ether, or ether-chloroform mixture is 
 used, it will form the upper layer in the separator, and may be 
 syphoned off, or the lower aqueous layer drawn off, and the 
 ethereal layer then transferred to another separator. Violent 
 shaking of the contents of the separator is to be avoided, 
 gentle shaking or rotating is quite sufficient to bring about the 
 desired transfer of alkaloid. Excessive shaking will cause 
 the formation of an emulsion of the water and solvent, which 
 is often hard to break up. 
 
 The final operation should always be the collection of the 
 free alkaloid by means of a portion of the immiscible solvent. 
 This solution is drawn off into a beaker, the solvent evap- 
 
 gquibb'a Pattern 
 
 Fig. 54. 
 
260 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 orated over a water-bath, using gentle heat, and the dry 
 residue in the beaker dissolved in a measured quantity of 
 standard acid V.S. The excess of acid V.S. is then titrated 
 with a standard alkali V.S. in the presence of a suitable 
 indicator. 
 
 In connection with this, attention is called to the fact that 
 some of the most important methods of isolating alkaloids 
 in drug assays and in toxicological investigations depend upon 
 the solubility of the free alkaloids in ether, chloroform, benzene, 
 etc., and the relative insolubility of alkaloidal salts in the 
 same solvents. There are, however, a number of exceptions 
 to this rule, and special attention is called to the fact that in 
 many cases alkaloids pass from decidedly acid aqueous solu- 
 tions (in which they certainly occur as alkaloidal salts) into 
 chloroform and ether, in the shaking-out methods. While 
 the amount of alkaloid so dissolved by these solvents is never 
 very great, still the quantity is an appreciable one. This 
 behavior occurs in the case of caffeine, colchicine, and narco- 
 tine; and under certain conditions also in the case of strychnine, 
 atropine, veratrine and other bases. This transfer of alkaloid 
 to chloroform occurs more particularly in the case of alka- 
 loids of weak basic character, and when the solution is neutral 
 or only feebly acid, or when the alkaloid is in combination 
 with a comparatively weak acid, as citric, tartaric, etc. Fur- 
 thermore, some alkaloidal salts, notably those of weak bases, 
 are transferred as such to chloroform, especially the salts of 
 hydrochloric acid, hydrobromic acid, and nitric acid. In the 
 case, however, of the sulphates, phosphates, tartrates, and 
 citrates of strongly basic alkaloids, no transfer occurs, or at 
 most, only minute quantities of the alkaloids pass over. In 
 order to prevent a transfer of alkaloid or alkaloidal salt out 
 of an aqueous solution to an immiscible solvent, the use of 
 sulphuric acid is to be recommended, and in all toxicological 
 
VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 261 
 
 investigations due regard should be paid to the above named 
 conditions. For more details, see the paper by Edward Schaer, 
 Proc. A. Ph. A., 1906, 425. 
 
 GENERAL METHODS OF ASSAYING DRUGS 
 
 No rule can be formulated as to the method of extraction, 
 or the solvent to be employed, which can be applied to all 
 drugs. Each must be dealt with in accordance with the 
 properties of the contained alkaloids and their state of com- 
 bination. Several methods, however, are in use which may 
 be applied to a large number of different drugs, and with 
 slight special modifications to many more. 
 
 Kebler's Modification of the Keller Method.* Treat 10 
 gms. of the dry powdered drug in a 250-cc. flask with 25 gms. 
 of chloroform and 75 gms. of ether; stopper the fl^sk securely, 
 agitate well for a few minutes and add 10 gms. of ammonia- 
 water and shake frequently during one hour. Then on adding 
 5 gms. more of the ammonia-water and shaking, the suspended 
 powder agglutinates into a lump, leaving the solution clear 
 after a few minutes' standing. Then proceed by A or B. 
 
 A. When the mixture has completely separated, 50 gms. 
 (representing 5 gms. of the drug) are poured off into a beaker 
 and heated on a water-bath until the solvent is evaporated. 
 Ten cc. of ether are then added and again evaporated. The 
 varnish-like residue is then dissolved in 15 cc. of warm alcohol 
 and water added to slight permanent turbidity, then the 
 indicator is added, followed by an excess of standard acid 
 solution, and the mixture retitrated with standard alkali. 
 
 B. When the mixture has completely separated, pour off 
 50 gms. into a separatory funnel, and add 20 cc. of acidulated 
 water, agitate, and when the liquids have separated draw 
 
 * J. A. C. S., XVII, 828. 
 
262 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 off the aqueous solution into a second separatory funnel. 
 Repeat this operation with two more portions of 15 cc. of 
 acidulated water. Now render the contents of the separatory 
 funnel alkaline by adding ammonia-water. This liberates the 
 alkaloids, which are then separated by treatment with a mix- 
 ture of chloroform, three parts (by volume), and ether, one 
 part, using three successive portions, first 20 cc, then twice 
 IS cc. 
 
 The chloroform-ether solution is heated on a water-bath 
 until the solvent is evaporated, and then the varnish-like 
 residue treated twice with 8 cc. of ether and again evaporated. 
 
 The residue is then dissolved in 15 cc. of alcohol, water 
 added to slight permanent turbidity, and then the indicator. 
 Titrate in usual way with decinormal acid and centinormal 
 alkali. 
 
 A serioJs objection to this method lies in the taking of a 
 so-called aliquot part: First, because of the well-known solu- 
 bility of ether in water, and conversely of water in ether, as 
 a result of which the volume of the ethereal stratum is materially 
 changed. Furthermore, commercial ether contains variable 
 quantities of alcohol, hence the change in volume will not 
 always be the same. 
 
 Another source of error in the aliquot part is found in 
 the volatile nature of the solvents used. In warm weather 
 it is impossible to avoid loss by volatilization, hence the aliquot 
 part taken is too large. 
 
 W. A. Puckner * has described a modification of the 
 Keller method which avoids the use of the aliquot part. He 
 uses only one-half of the ethereal solvent for the maceration, 
 and after the usual maceration transfers the drug to a small 
 percolator in which, after the ethereal solution has been well 
 
 * Ph. Rev. XVI, 180, and XX, 457. 
 
VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 263 
 
 drained off, the marc is percolated with the same menstruum 
 to complete exhaustion. The quantity of ethereal solvent 
 required is not materially greater than in the Keller method, 
 while the quantity of alkaloid obtained for weighing or titrating 
 is larger because it represents the whole of the sample taken 
 for the assay. In the case of drugs containing a very small 
 proportion of alkaloid this is an important advantage. 
 
 The objection to this plan is that the transfer of the mass 
 from the flask in which the maceration has been conducted 
 to a suitable percolator, which should not be more than 3 cm. 
 in diameter, requires very dextrous manipulation, or it will 
 be attended with loss of alkaloid. 
 
 Assay of Galenical Preparations. 
 
 J. U. Lloyd's Method. One gm. of a solid extract which 
 has been dissolved in 5 to 8 cc. of an alcoholic menstruum or 
 a corresponding volume of the tincture evaporated to this 
 bulk, or 5 cc. of the fluid extract in a flat-bottomed porcelain 
 mortar, are mixed with 2 cc. of a solution of perchloride of 
 iron. To this is added sodium bicarbonate with constant 
 trituration until a stiff magma * results. Extract this magma 
 by repeated trituration with chloroform, using first 20 cc. 
 and then three portions of 10 cc. each, decanting them severally 
 by means of a guiding-rod, being careful that no suspended 
 portions of the magma are drawn off. In order to make sure 
 that all of the alkaloid has been extracted, add 5 cc. more of 
 chloroform, draw it off, evaporate on a watch-glass, dissolve 
 
 * The ferric hydroxid which is produced in the above process serves to 
 attract most of the tannates, gums, vegetable acids, and coloring matters, 
 while the excess of sodium bicarbonate liberates the alkaloids, which are then 
 dissolved by the chloroform. If the fluid extract is strongly alcoholic the 
 chloroform will not separate easily, in which case the addition of a few cc. of 
 water containing a very little glucose will cause a sharp separation. 
 
26-1 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 residue in dilute sulphuric acid, and test for alkaloids by 
 Wagner's or Mayer's reagent. ^ 
 
 The mixed chloroformic extracts are collected and may be 
 estimated volumetrically as follows: 
 
 Method A. To be used if the chloroformic extract is not 
 colored. 
 
 The chloroformic so'ution is evaporated to dryness in a 
 
 flask placed on a water-bath. To this residue is added an 
 
 N 
 accurately measured excess of — sulphuric V.S. and the solution 
 
 diluted with a little water, the indicator added and the excess 
 
 N 
 of standard acid solution estimated by titrating with — potas- 
 
 N 
 sium hydroxid V.S. The number of cc. of the — alkali V.S. 
 
 5° 
 
 N 
 used, divided by 2 and subtracted from the volume of — acid 
 
 25 
 
 V.S. originally added, will give the number of cc. of the latter 
 
 required for the alkaloid. This number, multiplied by the 
 
 proper factor, will give the total alkaloid present in the fluid 
 
 extract. 
 
 Example. The chloroformic residue obtained from 5 cc. 
 
 of a fluid extract of hyoscyamus was dissolved in 12 cc. of 
 
 N . N 
 
 — acid V.S., the solution titrated with — alkali V.S. 20.6 
 
 25 50 
 
 cc. of the latter were used. 
 
 N . . N 
 
 20.6 cc. of — V.S. is the equivalent of 10.^ cc. of — V.S. 
 50 ^ ^25 
 
 N 
 10.3 cc. subtracted from 12 cc, the amount of — - acid originally 
 
 N 
 added, leaves 1.7 cc, the quantity of — acid V.S., which was 
 
 25 
 
 N 
 required for the alkaloid. This, multiplied by the — factor 
 
VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 265 
 
 for total alkaloids of hyoscyamus, 0.01157 gm., gives the 
 quantity of alkaloids present in the 5 cc, which quantity 
 multiplied by 20 gives the per cent: 
 
 1.7X0.01157 = 0.019669 gm.X 20 = 0.39338 per cent. 
 
 Method B. This method may be employed if the chloro- 
 formic extract is highly colored, the indicator fluorescein 
 being used. 
 
 The residue from the evaporation of chloroform is dis- 
 solved in 10 cc. of acid-free alcohol; then water is added to 
 
 N 
 slight turbidity, followed by a measured excess of — acid V.S.; 
 
 N 
 then the titration is completed with — alkali V.S. The first 
 
 appearance of fluorescence marks the completion of the reaction. 
 This is best observed by holding the flask over a dark surface 
 and viewing by reflected light. 
 
 Method C. This is to be used in the case of highly colored 
 extracts. It consists in removing the alkaloid in a pure state 
 by shaking out in a separating funnel with immiscible sol- 
 vents. The chloroformic extract is shaken out with several 
 portions of acidulated water; this removes the alkaloid, leaving 
 resins, fats, coloring matters, etc., in the chloroform. The 
 acid alkaloidal solution is then treated with ether in a second 
 separator, ammonia added to alkaline reaction, and the alka- 
 loid thus liberated by ammonia dissolves in the ether, from 
 which it is obtained by evaporation and estimated acidimetrically. 
 
 Another Method. Add ammonia and shake out directly 
 with an immiscible solvent. Wash the alkaloid out of this 
 with acidulated water and again shake out with a suitable 
 immiscible solvent. 
 
 J. Katz's Method (Arch. d. Pkarm., 1898, i; Am. Dr., 
 1898, 281). This method has the advantage of all other methods 
 
266 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 in that it enables one to rapidly and accurately estimate the 
 alkaloids in a preparation without the necessity of applying 
 heat for any purpose whatever during the process. The assay 
 may be made in from one to three hours. Twenty-five cc. of 
 the tincture of an alcoholic strength of about 45 per cent are 
 placed in a separatory funnel, i cc. of a 33 per cent solution 
 of soda added, and the mixture shaken for 5 minutes with 
 50 cc. of ether, and set aside until the liquids have com- 
 pletely separated into two layers. The lower dark-colored 
 aqueous layer is drawn off into a beaker. The ethereal layer 
 which, besides the alkaloid, has taken up most of the alcohol 
 and some coloring matter, is shaken up with 3 cc. of water, 
 which, after separating, is drawn off and added to the first 
 aqueous liquid. The ethereal liquid is then poured into a 
 suitable flask, while the combined aqueous liquid is further 
 treated with two portions of ether (25 cc. each) the ether to 
 contain about 10 per cent of alcohol. These ethereal extractions 
 are washed each with 1.5 cc. of water; the first extraction 
 after washing may be added to the original ethereal solution, 
 while the second is reserved and later on employed for washing 
 the flask. The ethereal solution, which contains still some 
 traces of aqueous fluid containing alkali, is dehydrated by 
 treatment with 2 or 3 gms. of exsiccated calcium sulphate, 
 and finally filtered into a glass-stoppered flask containing 50 
 cc. of water. 
 
 N 
 Titration by means of acid V.S. is employed, using 
 
 alcholic solution of iodeosin (1-250) as indicator. 
 
 The Method as above described is obviously applicable 
 only to such alkaloids as are readily soluble in ether. If an 
 estimation of alkaloids insoluble or only slightly soluble in 
 ether, but soluble in chloroform, is to be made, the method- 
 is modified as follows: 
 
VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 267 
 
 Twenty-five cc. of the tincture of 45 per cent alcoholic 
 strength are vigorously shaken for 5 minutes with 30 cc. of a 
 mixture of i part of chloroform and 2 parts of ether. The 
 solution so obtained is washed with 3 cc of a 30 per cent 
 solution of sodium chlorid. This operation is repeated twice, 
 using each time 15 cc. of the ether-chloroform mixture and 
 washing with 1.5 cc. of sodium chlorid solution. 
 
 If the Separation of the aqueous and ethereal liquids is not 
 distinct an additional 2 or 3 gms. of sodium chlorid may be 
 used. This prevents the emulsification, which, if pure water 
 were employed, would occur. 
 
 If the Tincture to be assayed contains more than 45 per 
 cent of alcohol it is necessary to add water to reduce it to 40 
 or 50 per cent. 
 
 Tinctures containing Chlorophyll or fat or fatty acids must 
 first be deprived of these constituents, as these substances, 
 possessing acid properties inferior to that of iodeosin, will act 
 the part of an alkali toward it and thus be recorded as alkaloid. 
 
 To remove the Chlorophyll and fatty acids, acidulate a 
 mixture of equal parts of the tincture and water with a few 
 drops of sulphuric acid, shake with talcum during several 
 hours, and, after subsidence of the latter, filter. Of this 
 filtrate 25 gms. (not cc, on account of the admixture of alcohol 
 and water causing change of volume) are taken and the 
 alkaloid estimated in the manner already described, after 
 first removing, if necessary, the last traces of fat by a single 
 shaking of the acid solution with petroleum ether. 
 
 For the Assay of extracts i to 1.5 gms. are dissolved in 
 from 40 to 50 cc. of 45 per cent alcohol to make a solution 
 containing less than 3 per cent extractive. For the assay of 
 fluid extracts 10 cc. are taken. 
 
 ' For details of special assays, see the author's Manual of 
 Volumetric Analysis. 
 
268 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The Influence of the Presence of Certain Volatile Solvents 
 
 upon the accuracy of alkaloidal titrations must be borne in 
 mind. Alcohol in may instances influences the color changes 
 of the indicators to the extent of rendering them indistinct 
 or entirely unrehable; while ether or chloroform materially 
 diminishes the sensitiveness of many of our most valued 
 indicators, among them phenolphthalein, rosahc acid, congo- 
 red, and luteol. On the other hand, fluorescein and gallein 
 may be mentioned as being more sensitive in the presence of 
 alcohol or even of ether. It is advisable, therefore, in most 
 cases, to perform the titration without the presence of the 
 above-named solvents. 
 
 Indicators Vary in their Degree of Sensitivenness to the 
 same Alkaloid, hence the choice of an indicator in a particular 
 case is a matter of importance. The following table, based 
 upon that of Klippenberger, will serve as an aid in the selection. 
 The smallest quantity which will procure a distinct tint should 
 be taken. 
 
 Atropine. Lacmo'.d, Fluorescein, lodeosin, Haematoxylin. 
 
 Brucine. Cochineal, lodeosin, Haematoxylin. 
 
 Cocaine. Lacmoid, Fluorescein, Cochineal, Hsematoxylin. 
 
 Coniine. Cochineal, Lacmoid, lodeosin, Haematoxylin. 
 
 Codeine. lodeosin, Lacmoid, Haematoxylin. 
 
 Emetine. lodeosin, Cochineal. 
 
 Morphine. Hcematoxylin, Cochineal, Lacmoid. 
 
 Quinine. HcBinatoxylin, Azolitmin, Fluorescein. 
 
 Strychnine. Azolitmin, lodeosin, Haematoxylin. 
 
 Sparteine. Hcematoxylin, Azolitmin. 
 
CHAPTER XIII 
 
 ESTIMATION OF PHENOL 
 
 Preparation of Decinormal Bromin V.S. (Koppeschaar's 
 Solution), Br= 79.92; 7.992 gm. in i liter. 
 
 KBr =119.02 NaBr =102.92 
 KBr03 = 167.02 NaBr03 = 150.92 
 
 This solution does not contain free bromin, but it contains 
 two salts, a bromid and a bromate, which, when treated with 
 hydrochloric acid, liberate a definite quantity of bromin. 
 
 It is made as follows: Dissolve 3 gms. of sodium bromate 
 and 50 gms. of sodium bromid (or 3.2 gms. of potassium 
 bromate and 50 gms. of potassium bromid) in sufficient water 
 to make 900 cc. 
 
 Transfer 20 cc. of this solution by means of a pipette into 
 a bottle having a capacity of about 250 cc, provided with a 
 glass stopper; add 75 cc. of water, then 5 cc. of pure hydro- 
 chloric acid, and immediately insert the stopper. 
 
 Shake the bottle a few times, then remove the stopper 
 just sufficiently to qu'ckly introduce 5 cc. of potassium iodid 
 solution (1-5), taking care that no bromin vapor escape, and 
 immediately stopper the bottle. 
 
 Agitate the bottle thoroughly, remove the stopper and 
 rinse it and the neck of the bottle with a little water so that 
 the washings flow into the bottle, then add from a burette 
 decinortnal sodium thiosulphate until the color of the free 
 iodin is nearly all discharged, then add a few drops of starch 
 
 269 
 
270 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 N 
 solution, and continue the titration with — thiosulphate V.S. 
 
 until the blue color disappears. 
 
 N 
 Note the number of cc. of the — sodium thiosulphate 
 
 thus used, and dilute the bromin solution so that equal volumes 
 
 N 
 of it and the — sodium thiosulphate will exactly correspond 
 
 to each other under the above-mentioned conditions. 
 
 Example. Assuming that the 20 cc. of bromin solution 
 
 N 
 required 25.2 cc. of the — thiosulphate to completely absorb 
 
 the iodin, the bromin solution must be diluted in the pro- 
 portion of 20 to 25.2; that is, each 20 cc. must be diluted to 
 make 25.2 cc. 
 
 Thus if 850 cc. are left, they must be diluted to make 
 107 1 cc, and the solution is decinormal. 
 
 A new trial should always be made after diluting, and 
 the bromin solution should correspond, volume for volume, 
 with the decinormal sodium thiosulphate. 
 
 The first step in the preparation of this solution is to 
 dissolve the salts; then hydrochloric acid is added, which 
 liberates a definite quantity of bromin, as the equation illus- 
 trates: 
 
 SNaBr -fNaBrOs +6HC1 = 6NaCl -hsBra -I-3H2O. 
 
 The stopper should be inserted into the bottle as soon 
 as the hydrochloric acid has been added, in order that no 
 bromin vapor escape, and the bottle rotated so as to mix the 
 acid thoroughly with the liquid. 
 
 The next step is to determine the quantity of bromin 
 which a definite volume of solution will liberate. The bromin 
 solution should be of such strength that 1000 cc. of it will 
 
ESTIMATION OF PHENOL 271 
 
 contain 7.976 gms. of available bromin. Bromin, like chlorin, 
 liberates iodin from potassium iodid, and is estimated in 
 the same manner. 
 
 One atomic weight of ibdin is liberated by one atomic 
 weight of bromin: 
 
 Br2 + 2KI = 2KBr + l2. 
 
 Thus by determining the quantity of iodin liberated the 
 quantity of bromin is found. 
 
 N 
 The iodin is determined by the — sodium thiosulphate, 
 
 one liter of which represents 12.692 gms. of iodin, which is 
 equivalent to 7.992 gms. of bromin, as shown by the following 
 equation : 
 
 (Bra) = I2 + 2(Na2S203 + sH20) 
 20 )159.84 20 )253.84 20 )496.44 j^ 
 
 7.992 gms. 12.692 gms. 24.822 gms. or 1000 cc. — V.S. 
 
 10 
 
 = 2NaI + Na2S406 + 10H2O. 
 
 The Assay of Phenol. Dissolve i gm. of the sample in 
 sufficient water to make 500 cc. of solution at the standard 
 temperature. Twenty cc. of this solution containing 0.04 gm. 
 of the sample are transferred to a glass-stoppered bottle, having 
 a capacity of about 200 cc. 
 
 To this, 30 cc. of decinormal bromin, followed by 5 cc. 
 of hydrochloric acid, are added, and the bottle immediately 
 stoppered, and shaken repeatedly during half an hour. 
 
 Then the stopper is removed just sufficiently to introduce 
 5 cc. of a 20 per cent aqueous solution of potassium iodid, 
 being careful that no bromin escape. 
 
 The bottle is then thoroughly shaken and the neck rinsed 
 with a little water he washingSi being allowed to flow into 
 the bottle. 
 
272 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The solution is now ready for titration, and the decinormal 
 sodium thiosulphate is delivered in from a burette, until the 
 iodin is almost completely absorbed; then add a few drops 
 of starch T.S., and continue the titration until the blue color 
 is just discharged. 
 
 About T cc. of chloroform may be used instead of starch. 
 The precipitated tribromphenol interferes somewhat with the 
 end-reaction when starch is used, and frequently with old 
 phenol solutions * the precipitate possesses a bluish color 
 which is not removed by an excess of sodium thiosulphate 
 and which makes the end-reaction difficult. This difficulty 
 is overcome by the use of a small quantity of chloroform which 
 dissolves the tribromphenol and admits of a very sharp end- 
 reaction. 
 
 When chloroform is used, the end-reaction is very clearly 
 
 defined, and is known by a colorless aqueous solution and 
 
 the chloroform being free from any tinge of pink, due to 
 
 traces of iodin 
 
 N 
 Note the number of cc. of — thiosulphate used; deduct 
 
 N 
 this number from 30 cc. (the quantity of — bromin originally 
 
 N 
 added), and the quantity of — bromin which went into com- 
 bination with the phenol is obtained. 
 
 N 
 Each cc. of — bromin represents 0.001556 gm. of pure 
 
 phenol. 
 
 N 
 Example. Assuming that 6 cc. of — sodium thiosulphate 
 
 were required to discharge the color of the starch iodid, this 
 
 F. X. Moerk, A. J. Ph., 1904, 475. 
 
ESTIMATION OF PHENOL 273 
 
 deducted from 30 cc. leaves 24 cc, the quantity which com- 
 bined with the phenol. 
 
 0.001566X24 = 0.037584 gm. 
 0.037584X100 
 
 0.04 
 
 93.96 per cent of pure phenol. 
 
 The above method originated with Koppeschaar, and is 
 the only volumetric method by which accurate results may 
 be obtained. 
 
 It is based upon the fact that bromin reacts with phenol, 
 producing an insoluble precipitate of tribromphenol. 
 
 The titration is not made directly; but the phenol solution 
 is treated with an excess of standard bromin solution in the 
 presence of some hydrochloric acid. The hydrochloric acid 
 liberates the bromin, and the freed bromin reacts with the 
 phenol, as shown by the equations : 
 
 (a) 5NaBr + NaBrOs + 6HC1 = 6NaCI + 3H2O + 3Br2 
 
 (b) CeHsOH + 3Br2 = CeHaBrsOH + 3HBr. 
 
 6 )94 6 )479.92 
 
 10 )15-66 10 )79-92 j^ 
 
 1.566 gms. 7-992 gms. or 1000 cc. — bromin V.S. 
 
 N 
 
 Thus each cc. of the — bromin represents 0.001566 gm. 
 
 of pure phenol. 
 
 The bromin solution which was added in excess, and 
 
 the liberated bromin of which is not fixed by phenol, is then 
 
 N 
 fond by residual titration with — sodium thiosulphate after 
 
 the addition of some potassium iodid. 
 
 The decinormal bromin solution and the decinormal sodium 
 thiosulphate solution being equivalent, each cc. of the latter 
 
274 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 consumed represents one cc. of the former. Then by sub- 
 tracting the number of cc. of the sodium thiosulphate solution 
 and from the number of cc. of bromin solution originally added, 
 the quantity of the latter which was actually consumed by 
 the phenol present is found. This number, when multiplied 
 by the factor for phenol, then gives the quantity of pure phenol 
 present. 
 
 The hydrochloric acid used in the above estimation must 
 contain no free chlorin. The potassium iodid must be free 
 from iodate. The starch T.S. should not be added imtil 
 most of the free iodin has been taken up, and the color of 
 the solution has diminished to light yellow. 
 
 The carbolic acid should be diluted with water before 
 titration, and should never be stronger than o.i gm. in 25 cc. 
 
CHAPTER XIV 
 
 SOME TECHNICAL EXAMINATION METHODS FOR FATS, 
 OILS AND WAXES 
 
 The Acid Value or Proportion of Free Fatty Acids. This 
 indicates the number of milligrams of KOH required to neutral- 
 ize the free fatty acids in i gm. of oil, fat or wax. This standard 
 
 N N 
 alkali used is in alcoholic solution and may be — , — , or weaker, 
 
 depending upon the nature of the fat. Phenolphthalein is the 
 indicator. The fat is dissolved, according to Geissler, in ether, 
 but alcohol or purified methylated spirit, chloroform or a mix- 
 ture of alcohol and ether may be used. The solvent, whichever 
 is used, must be free from acidity, and should be neutralized 
 
 N 
 with — alkali if necessary. 
 
 The Process. Ten gms. of the oil are accurately weighed 
 into a flask and about 50 cc. of solvent added. A few drops 
 of phenolphthalein are then added and the titration with 
 
 N 
 alcoholic — potassium hydroxid solution begun, shaking con- 
 stantly until the first appearance of a pink color. Care must 
 be taken not to add too great an excess of the alkali, other- 
 wise saponification will occur. A small excess may, however, 
 be added and retitrated with standard acid; a more distinct 
 end-point is then obtained. In the case of waxes or solid 
 fats, .the solvent is added, heat applied until it boils, and the 
 
 titration at once started, 
 
 275 
 
276 THE ESS1:NTIALS OF VOLU^rETRIC ANALYSIS 
 
 N 
 One cc. of — KOH = 0.02805 g'^- °^ KOH or 28.05 mgms. 
 
 .The number of cc. used, multiplied by 28.05 and then 
 divided by the weight of oil taken, gives the milligrams of 
 KOH neutralized by the free fatty acids of the oil, i.e., the 
 acid value. 
 
 The Saponification Value (Kottstorfer Number).* This 
 indicates the number of milligrams of KOH required for the 
 complete saponification of one gram of fat or oil. Reagents 
 required are: 
 
 Alcoholic Potassium Hydroxid Solution, made by dissolving 
 40 gms. of potassium hydroxid in i liter of 95 per cent alcohol. 
 
 Half-Normal Hydrochloric Acid Solution. Each cc. of 
 which =28.05 mgms. of KOH. 
 
 Indicator. Phenolphthalein i gm. in 100 cc. of 95 per 
 cent alcohol. 
 
 The Process. Into an Erlenmeyer flask capable of holding 
 200 cc, accurately weigh about 1.5 gms. of the fat (previously 
 purified and filtered). Run into this from a burette 25 cc. 
 of the alcoholic potassium hydroxid solution. Then place a 
 small funnel in the mouth of the flask or cover it loosely with 
 a watch crystal, and set it in a steam-bath for half an hour 
 or until the fat is entirely saponified. The operation is facili- 
 tated by occasional agitation. The flask is then removed, 
 its contents cooled, and titrated with the half-normal hydro- 
 chloric acid, using phenolphthalein as indicator. The alcoholic 
 potassium hydroxid solution is standardized by conducting a 
 blank experiment similar in every detail to the above with the 
 exception that the fat is omitted. 
 
 The Kottstorfer number is then ascertained by subtracting 
 the number of cc. of the standard hydrochloric acid used in 
 the analysis from the number necessary to neutralize 25 cc. 
 
 * J. Kottstorfer, 1879, Zeitschr. anal. Chem., XVIII, 199. 
 
TECHNICAL KXA.\[INATION METHODS FOR FATS 277 
 
 of the alcoholic potassium hydroxid solution in the blank 
 experiment, multiplying the difference by 28.05 and dividing 
 by the weight of fat taken. 
 
 Example. 1.5 gms. of the fat was treated with 25 cc. of 
 alcoholic potassium hydroxid solution, under the above described 
 
 N 
 conditions, and in titrating the excess, 21.5 cc. of — hydro- 
 chloric acid was required. In the blank experiment 25 cc. 
 of the alcoholic potassium hydroxid required 32 cc. of the 
 
 N 
 
 — acid. Then 32— 21.5 = 10.5. 
 
 10.5X28.05 
 
 ^196. 
 
 196 is the milligrams of KOH neutralized by i gra. of the 
 oil, or the Kottstorfer saponification number. 
 
 TABLE SHOWING REQUIREMENT AS TO SAPONIFICATION 
 
 NUMBER 
 
 Lard cjil 195 to 197 
 
 Almond oil (expressed) . . 191 to 200 
 
 Cottonseed oil 191 to 196 
 
 Linseed oil 187 to 195 
 
 Cod-liver oil i75 to 185 
 
 Olive oil 191 to 195 
 
 Castor oil 179 to 180 
 
 Theobroma oil 188 to 195 
 
 Croton oil 212 to 218 
 
 Yellow wax 90 to 96 
 
 The Reichert Number or Volatile Fatty Acid Value. This 
 indicates the number of cc. of decinormal KOH required to 
 neutralize the volatile fatty acids distilled from 2.5 gms. of 
 fat. This method is conducted as follows: 2.5 gms. of the 
 clear filtered fat are taken in an Erlenmeyer flask of about 
 150-cc. capacity with i gm. of potassium hydroxid and 20 cc. 
 of 80 per cent alcohol, and the whole digested on a water-bath 
 (rotating the flask frequently) until saponification is complete 
 and the alcohol all removed. 50 cc. of water are then added, 
 
278 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 and 20 cc. of dilute sulphuric acid (1:10), and the mixture 
 distilled. The distillate is collected in a 50-cc. flask into 
 which is set a funnel with a wetted filter to collect any insoluble 
 fat acid. The first 10 or 20 cc. of distillate are returned to 
 the flask, and the 50 cc. distilled. This is treated with a few 
 drops of phenolphthalein and then titrated with decinormal 
 potassium hydroxid. The number of cc. consumed is the 
 Reichert Number. 
 
 In the Meissl modification (the Reichert-Meissl method) 
 5 gms. of the fat are taken and a more complete distillation 
 of the volatile acids effected. The method is particularly 
 useful in determining the genuineness and purity of butter. 
 
 The Reichert-Meissl Number. This is undoubtedly the 
 best process for detecting the admixture of foreign fats with 
 butter. This process depends upon the fact that butter con- 
 tains certain constituents which, when appropriately treated, 
 yield volatile acids in much larger quantity than is obtained 
 from any of the practicable substitutes for butter. These 
 acids are principally butyric and caproic. The process con- 
 sists in saponifying the fat with an alkali, then separating 
 the fatty acids by neutralization, and distilling off the volatile 
 acids for titration with standard alkali. 
 
 The operations involved in this process do not admit of 
 any arbitrary variation, and reliable and comparable results 
 can only be obtained by strictly adhering to the prescribed 
 details. 
 
 The following process is adopted by the Association of 
 Ofi&cial Agricultural Chemists. The solutions required are: 
 
 Sodium Hydroxid Solution. One hundred gms. of sodium 
 hydroxid are dissolved in 100 cc. of distilled water. The 
 alkali should be as free as possible from the carbonates, and 
 be preserved out of contact with the air. 
 
 Potassium Hydroxid Solution. One hundred gms. of the 
 
TECHNICAL EXAMINATION METHODS FOR FATS 279 
 
 purest potassium hydroxid are dissolved in 58 cc. of hot dis- 
 tilled water, cooled in a stoppered vessel, and the clear liquid 
 decanted and preserved out of contact with the air. 
 
 Sulphuric Acid. Mix 200 cc. of the strongest acid with 
 1000 cc. of water. 
 
 Alcohol of about 95 per cent, redistilled from caustic soda. 
 
 Standard Barium Hydroxid Solution, accurately standard- 
 ized, approximately decinormal. 
 
 Indicator. Dissolve i gm. of phenolphthalein in 100 cc. 
 of 95 per cent alcohol. 
 
 The process: 
 
 Weighing the Butter. The butter to be examined should 
 be melted and kept in a dry, warm place at about 60° C. for 
 two or three hours, until the water and curd have entirely 
 settled out. The clear supernatant fat is poured off and 
 filtered through a dry filter-paper in a jacketed fxmnel con- 
 taining boiling water. Should the filtered fat in a fused state 
 not be perfectly clear, it must be filtered a second time. This 
 is to remove all foreign matter and any trace of moisture. 
 The saponification flasks are prepared by thoroughly washing 
 with water, alcohol and ether, wiping perfectly dry on the 
 outside, and heating for one hour at the temperature of boiling 
 water. The flasks should then be placed in a tray by the 
 side of the balance and covered with a silk handkerchief until 
 they are perfectly cool. They must not be wiped with a silk 
 handkerchief within fifteen or twenty minutes of the time 
 they are weighed. The weight of the flasks having been 
 accurately determined, they are charged with the melted fat 
 in the following way: 
 
 A pipette with a long stem, marked to deliver 5.75 cc, 
 is warmed to a temperature of about 50° C. The fat, having 
 been poured back and forth once or twice into a dry beaker 
 in order to thoroughly mix it, is then taken up in the pipette 
 
280 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 and the nozzle of the pipette carried nearly to the bottom of the 
 flask, having been previously wiped to remove any adhering 
 fat, and 5.75 cc. of fat are allowed to flow into the flask. 
 After the flasks have been charged in this way they should 
 be recovered with the silk handkerchief, allowed to stand 
 fifteen or twenty minutes, and again weighed. 
 
 The Saponification. Three methods may be employed: 
 I. Under Pressure with Alcohol. Ten cc. of 95 per cent 
 
 alcohol are added to the fat in 
 the flask, and then 2 cc. of the 
 caustic soda solution. A soft 
 cork stopper is now inserted in 
 the flask and tied down with 
 a piece of twine. The saponi- 
 fication is then completed by 
 placing the flask upon the 
 water or steam bath (see Fig. 
 55). The flask during the 
 saponification, which should 
 last one hour, should be 
 gently rotated from time to 
 time, being careful not to pro- 
 ject the soap for any distance 
 up its sides. At the end of 
 an hour the flask, after having 
 been cooled nearly to room 
 temperature, is opened. 
 
 Fig. 55. 
 
 2. Under Pressure without the Use of Alcohol. Place 2 
 cc. of the potassium_ hydroxid in the flask containing the fat, 
 which must be roimd-bottomed and made of well-annealed 
 glass to resist the pressure; cork, and heat as in the previous 
 method. Rotate the flask very gently during the saponifica- 
 tion, taking great care that none of the fat rises on the sides 
 
TECHNICAL EXAMINATION METHODS FOR FATS 281 
 
 of the flask out of reach of the alkali. Potash makes a softer 
 soap than soda and thus allows a complete saponification 
 without the use of alcohol. This method avoids the danger 
 of formation of esters and the trouble of removing the alcohol 
 after saponification. 
 
 3. With a Reflux Condenser and the Use of Alcohol. Place 
 10 cc. of the 95 per cent alcohol In the flask containing the 
 fat, add 2 cc. of the sodium hydroxid solution with a reflux 
 condenser (a glass tube not less than i m^eter in length is 
 allowable) and heat on the steam bath until the saponification 
 is complete. 
 
 Removal of the Alcohol. The stoppers having been laid 
 loosely in the mouth of the flask, the alcohol is removed by 
 dipping the flask into a steam bath. The steam should cover 
 the whole of the flask except the neck. After the alcohol is 
 nearly removed, frothing may be noticed in the soap, and 
 to avoid any loss from this cause or any creeping of the soap 
 up the sides of the flask, it should be removed from the bath 
 and shaken to and fro until the frothing disappears. The 
 last traces of alcohol vapor may be removed from the flask 
 by waving it briskly, mouth down, to and fro, or better, by 
 a current of carbon dioxid free air. 
 
 Dissolving the Soap. After the removal of the alcohol 
 the soap should be dissolved by adding 135 cc. of recently 
 boiled distilled water (or 132 cc. if potassium hydroxid was 
 used in the saponification), warming on the steam bath, with 
 occasional shaking until solution of the soap is complete. 
 
 Setting Free the Fatty Acids and Distilling. Cool to from 
 60° to 70° C, throw in a few pieces of pumice stone, add 5 
 cc. of the dilute sulphuric acid (or 8 cc. if potassium hydroxid 
 was used in the saponification), stopper as in the method of 
 saponification, and heat on the water-bath until the fatty 
 acids form a clear, transparent layer on top of the water. 
 
282 THE ESSENTIALS OF VOLUIMETRIC ANALYSIS 
 
 This may take several hours. Cool to room temperature, 
 add a few pieces of pumice stone, and connect with a glass 
 condenser, as in Fig. 56. 
 
 Heat slowly with a naked flame until ebullition begins, 
 and distil, regulating the flame in such a way as to collect no 
 cc. of distillate in as nearly thirty minutes as possible. 
 
 Mix this distillate, filter through a dry filter, and titrate 
 100 cc. with the standard barium hydroxid solution, using 
 
 Fig. 56. 
 
 0.5 cc. of phenolphthalein as indicator. The red color should 
 remain unchanged for two or three minutes. 
 
 Increase the number of cubic centimeters of tenth-normal 
 alkali used by one-tenth, divide by the weight of fat taken, 
 and multiply by five to obtain the Reichert-Meissl number. 
 Correct the result by the figure obtained in a blank exper- 
 
 r 
 
 iment. 
 
 WTien treated as above described, 5 gms. of genuine butter 
 
 N 
 never yields less acidity than is represented by 24 cc. of — 
 
 alkali. It is true, however, that the butter made from the 
 milk of a single cow, especially towards the end of her period 
 of lactation, has been known to fall sHghtly below this figure. 
 
TECHNICAL I'XAMINATION METHODS FOR FATS 283 
 
 but the average butter, as produced from the mixed milk of 
 a herd, usually requires 27 cc. or more. Oleomargarin requires 
 about I cc. of beef-fat, and lard about the same. Cacao 
 butter requires about 7 cc. 
 
 The percentage of butter-fat in a mixture of fats, 5 gms. 
 being taken: (w— 0.6) X 3.65 = percentage of true butter-fat. 
 
 lodin Absorption Number of Fats and Oils (Hiibl's Num- 
 ber).* This is the percentage of iodin absorbed by a fat 
 or an oil under certain conditions. In other words it is the 
 number of parts of iodin absorbed by 100 parts of an oil. 
 
 Reagents required: 
 
 (a) Hiibl's Iodin Solution. Dissolve 25 gms. of pure iodin 
 in 500 cc. of alcohol, and mix this solution with 500 cc. of 
 alcohol containing 30 gms. of pure mercuric chlorid, and set 
 aside for twenty-four hours. The mercuric chlorid solution 
 should be filtered if necessary before it is mixed with the 
 alcoholic iodin solution. This solution loses its strength 
 rapidly, and should therefore be tested before using. 
 
 (b) Decinormal sodium thiosulphate. 
 
 (c) Potassium iodid solution, 20 gms. in 100 cc. 
 
 (d) Starch paste indicator. 
 
 The Process. To about 0.3 gm.| of the fat or oil, accu- 
 rately weighed and dissolved in 10 cc. of chloroform, contained 
 in a glass-stoppered bottle of 250 cc. capacity, add 25 cc. of 
 the iodin solution. Stopper the bottle securely and place in 
 a cool, dark place for four hours.} At the expiration of this 
 
 * Dingler's PoIyt._ Jour., 1884, 281; Am. Ch. Jour., VI, 285. 
 
 t In the case of drying oils which have a very high absorbent power, as 
 linseed oil, use from 0.15 to 0.20 gm.; for oil of theobroma and similar fats 
 use 0.80 gm. 
 
 I The time allowed does not give the complete iodin absorption power 
 of an oil or fat and cannot be compared with determinations in which six to 
 twelve hours have been used. It gives, however, very satisfactory comparative 
 results, but the time factor must be very closely observed. 
 
284 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 time, the mixture must still possess a brown color; if it does 
 not, a further measured quantity of the iodin solution must 
 be added and the mixture again set aside. Finally 20 cc. 
 of the potassium iodid solution are added together with 50 
 cc. of water, and the mixture titrated with the decinormal 
 soditun thiosulphate xmtil the color is almost discharged, 
 when a few drops of starch indicator are added and the 
 titration continued until the solution is colorless. 
 
 The Standardization of the Iodin Solution is effected by 
 subjecting it to the same treatment as in the assay, but with 
 the oil omitted, and at the end of four hours titrating with 
 the decinormal thiosulphate. 
 
 The difference between the number of cc. of thiosulphate 
 solution used in the blank test and the number used in the 
 actual assay, is multiplied by 12.59, and this divided by 3, 
 gives the iodin value of the oil under analysis. 
 
 When the quantity of the oil or fat taken is not 0.3 gm., 
 then the product is not divided by 3, but by the figure corre- 
 sponding to the quantity taken; thus if 0.15 gm. are taken 
 the product is divided by 1.5. 
 
 Another way of making the calculation is as follows: 
 
 The difference between the cc. of thiosulphate used in 
 the blank test and the cc. used in the analysis, is multiplied 
 by 0.01259, then by 100, and the product divided by the 
 weight in grams of the oil taken for analysis. 
 
 Example: 
 
 Number of cc. of thiosulphate used in blank test 45.4 
 
 " " analysis 26.1 
 
 a II 
 
 li II 
 
 representing iodin absorbed 19.3 
 
 . 19.3X0.01259X100 
 
 Iodin number is =80.9. 
 
 0.3 ^ 
 
TECHNICAL EXAMINATION METHODS FOR FATS 285 
 
 This method, as is well known, is based upon the fact 
 that the unsaturated glycerids in the oils form addition products 
 with the iodin. The mercuric chlorid and iodin contained 
 in the alcoholic solution interact with a formation of mercuric 
 chloriodid and iodin chlorid; the latter is supposed to be 
 the active agent. 
 
 HgCl2+l2=HgClI+ICl. 
 
 The mercuric chlorid also acts the part of a carrier of halogen 
 similar to that played by mercury in the Kjeldahl process 
 when dissolving the substance in sulphuric acid. 
 
 Gill and Adams, J. A. C. S., xxii, 12, call attention to 
 the fact that not only addition products, but also substitution 
 products, are formed in this process, which vary in amount 
 with the time of action and the strength of the solution. This 
 is a feature which interferes with the accurate determination 
 of the iodin number, giving as it does a higher figure. In 
 order to prevent this formation of substitution products, and 
 thus overcome the discrepancy in results, these authors suggest 
 the use of mercuric iodid instead of mercuric chlorid, making 
 the solution with methyl alcohol (free from acetone and 
 anhydrous). They claim that by the use of a solution so 
 made a truer iodin absorption number is obtained. A great 
 disadvantage of the Hiibl method is that the solution quickly 
 loses its strength; another is, the length of time required 
 for the absorption. As above described, four hours are required, 
 but this is not sufficient time to allow of complete absorption 
 of the iodin. It is, however, a good practice to have a fixed 
 time limit; the process then gives very satisfactory com- 
 parative results. 
 
2Sf) THK ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 TABLE SHOWING lODIN ABSORPTION NUMBER FOR SOME 
 COMMON OILS 
 
 Hubl's No. 
 4 Hours. 
 
 Almond oil (expressed) yS . 2 
 
 Butter 35.4 
 
 Castor oil 87.3 
 
 Cocoanut Oil 8.4 
 
 Cod-liver oil 144-3 
 
 Cottonseed oil io4-3 
 
 Lard oil 69.5 
 
 Linseed oil ... / '79-6 
 
 Mustard oil 113. i 
 
 Olive oil 86 . 1 
 
 Oleomargarin 52.1 
 
 Peanut oil 96.3 
 
 Rape oil 102.4 
 
 Seame oil 106.4 
 
 Theobroma oil 35-4 
 
CHAPTER XV 
 
 ESTIMATION OF SUGARS 
 
 Preparation of Fehling's Solution, (a) The Copper Solu- 
 tion. 34.67 gms. of carefully selected small crystals of pure 
 cupric sulphate are dissolved in suf&cient water to make, 
 at or near 15° C. (59° F.), exactly 500 cc. Keep in small 
 well-stoppered bottles. 
 
 (6) The Alkaline tartrate Solution. 173 gms. of potassium 
 and sodium tartrate (Rochelle salt) and 75 gms. of potassium 
 hydroxid are dissolved in sufi&cient water to make, at or near 
 15° C. (59° F.), exactly 500 cc. Keep in small rubber-stoppered 
 bottles. 
 
 For use, equal quantities of the two sulutions should be 
 mixed at the time required. 
 
 One molecular weight of water-free glucose will reduce 
 five molecular weights of cupric oxid, i.e., 180.12 gms. of 
 glucose will reduce 1248.7 gms. of crystallized copper sul- 
 phate (CUSO4-I-5H2O). If pure chemicals are used in the 
 preparation of this solution there will be no need of standard- 
 izing it, but if the solution is old and its titer doubtful, the 
 following method of standardization may be employed. 
 
 The Standardization. Dissolve 0.95 gm. of pure cane 
 sugar in 50 cc. of water, add 2 cc. of hydrochloric acid, and 
 heat to 70° C. for ten minutes. Then neutralize with sodium 
 carbonate and dilute to i liter. 
 
 Fifty cc. of this solution should exactly reduce 10 cc. of 
 Fehling's solution, when treated as described below. 
 
 287 
 
288 THE ESSENTIALS OF \'OLUMETRIC ANALYSIS 
 
 Ten cc. of the mixed Fehling's solution is equivalent to 
 
 Glucose 0-050 
 
 Maltose 0.0806 
 
 Inverted cane-sugar 0.0475 
 
 Inverted starch 0.045 
 
 Lactose 0.0678 
 
 Estimation of Reducing Sugars. {Glucose, Maltose, In- 
 verted Cane-sugar, Lactose and Inverted Starch. 0.5 gm. or 
 less of the sugar is dissolved in 100 cc. of distilled water, 
 and put into a burette. Ten cc. of the mixed Fehling's solu- 
 tion, accurately measured, is put into a loo-cc. Erlenmeyer 
 flask, or in a white porcelain dish, diluted with 40 cc. of 
 distilled water, and rapidly heated to boiling. The solution 
 should remain perfectly clear and retain its blue color. If 
 it does not it should be rejected. 
 
 The hot, clear blue solution is immediately titrated with 
 the sugar solution, which should be added in small portions at 
 a time, boiling after each addition until the copper is com- 
 pletely precipitated and the blue color of the solution is entirely 
 destroyed. The solution is then boiled about two minutes 
 longer, and the amount of the sugar solution used is read off. 
 
 The Calculation. Ten cc. of Fehling's solution are always 
 taken; and whatever the quantity of glucose or other sugar 
 solution is required to effect its complete reduction, that 
 quantity contains the equivalent of 10 cc. of Fehling's solu- 
 tion, as shown in the table above. 
 
 Thus if 12 cc. of the sugar solution were required to reduce 
 10 cc. of Fehling's solution, the 12 cc. contain c.05 gm. of 
 glucose or 0.0806 gm. of maltose, etc. One hundred cc. of the 
 solution therefore contain x gm. of glucose. 
 
 o.oi^Xioo ^ , 
 
 — = 0.416 gm. glucose. 
 
 12 
 
ESTIMATION OF SUGARS 289 
 
 In order to obtain reliable results it is important that the 
 process be carried out exactly as laid down in the above 
 directions, and that the quantity of sugar present in solution 
 be no greater than one per cent. The degree of heat and 
 the time occupied in the process, as well as the concentration 
 of the Fehling's solution, have a very important bearing upon 
 the accuracy of the results. 
 
 It is advisable to complete the titration in as short a time 
 as possible. A preliminary test should always be made, in 
 which the approximate quantity of the solution required is 
 found; then a second and more accurate titration can be 
 done in which the sugar solution may be added more boldly, 
 and the time of boiling and exposure of the copper solution 
 to the air much lessened. The complete reduction of the 
 copper (using tindiluted Fehling's), after the addition of the 
 requisite quantity of sugar, does not take place instantly. 
 The time required varies somewhat with the different sugars. 
 For instance, with glucose, invert sugar and levulose, the 
 reduction is not complete until after heating two minutes; 
 with maltose four minutes, and with lactose six minutes are 
 required. 
 
 If the sugar to he examined by either glucose, maltose, or 
 lactose, it may be titrated directly; but if it be sugar-cane, it 
 must first be inverted. This is done by dissolving the sugar 
 (0.5 gm.) in about 100 cc. of water, adding 3 or 4 drops of 
 strong hydrochloric acid, and boiling briskly for ten or fifteen 
 minutes. This is then allowed to cool, neutralized with potas- 
 sium hydroxid, and made up to 100 cc. with water. 
 
 The sugar in urine may be estimated by this process. The 
 urine is placed in the burette and run into the boiling Fehling's 
 solution in the usual manner. If it contain a large quantity 
 of sugar, it must be diluted two or three times. 
 
 In estimating with Fehling's solution it is well to attach a 
 
290 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 rubber tube eight to twelve inches in length to the lower end 
 of the burette, so that the boiling need not be done directly 
 under the burette, and thus cause incorrect readings through 
 the expansion of the liquid therein. 
 
 Determination of the End-point, It is always somewhat 
 difficult to determine the exact point at which the blue color 
 disappears, owing to the presence of the precipitated suboxid 
 of copper. This difficulty may be overcome by the addition 
 of some substance which will prevent the precipitation of 
 the cuprous oxid, such as ammonium hydroxid or potassium 
 ferrocyanid. When the latter is used the disappearance of 
 the blue color can then be readily seen, as the solution remains 
 clear to the end, turning from blue to green, and finally brown, 
 which indicates the end of the reaction. 
 
 Professor Bartley reports this method as accurate, reliable, 
 and rapid, provided the solution be not boiled during the 
 reduction. He recommends adding to the Fehling's solution 
 in the porcelain basin lo cc. of a lo per cent freshly prepared 
 solution of potassium ferrocyanid and 30 cc. of water. The 
 ferrocyanid does not precipitate the copper in alkaline solution. 
 
 L. Beulaygue (Compt. rend., 138, 51) suggests the following 
 method for determining the end-reaction when titrating sugaT 
 solution in the usual manner with Fehling's reagent, using 
 solution of sodium monosulphid as the indicator. When the 
 end of the reaction is near, a little of the hot solution is applied 
 by means of a glass rod, to two superimposed white filter 
 papers. The upper one acts as a filter, retaining the par- 
 ticles of cuprous oxid. The lower paper is withdrawn, and the 
 moist spot touched with a drop of the sodium monosulphid 
 reagent, when an immediate black stain of cupric sulphid 
 is formed if the reaction is not complete. By successive 
 spotting out and testing in this manner, a very accurate reading 
 of the end-reaction may be obtained. It is important, when 
 
ESTIMATION OF SUGARS 291 
 
 Standardizing the Fehling's solution, that the same indicator 
 should be employed. Potassium ferrocyanid in solution, 
 acidified by either hydrochloric or acetic acid, may be employed 
 in a similar manner. The end-reaction in then indicated by 
 the disappearance of the red color from the last spot. 
 
 E. F. Harrison,* who has employed Fehling's solution 
 somewhat extensively in quantitative sugar estimations, has 
 found the indicators usually recommended to determine the 
 end-point of reaction to be unsatisfactory. It was suggested 
 by him that the action of cupric salts in liberating iodin from 
 iodid might be utilized for this purpose with advantage, and 
 his experiments determine the superiority of this over the 
 other indicators heretofore proposed. The indicator is pre- 
 pared by boiling 0.05 gm. of starch with a few cc. of water, 
 adding 10 gm. of potassium iodid and diluting to 100 cc. 
 This indicator should be prepared as required. In use 0.5 
 to i.o cc. of this solution is acidified with about 5 or 10 
 drops of acetic acid, and one drop or more of the liquid .in 
 process of titration added. As long as unreduced copper is 
 present, a color is produced, varying from red to blue, and 
 of greater or less intensity, according to the nearness of the 
 end-point. The production of no color marks the end of 
 the reaction. The indicator is available with one drop of a 
 solution containing one part of cupric sulphate in twenty 
 thousand parts. 
 
 Estimation of Maltose in Malt Extract. A half per cent 
 sohition of the sample is prepared, and titrated into 10 cc. 
 of Fehling's solution in the manner described above. 
 
 Estimation of Starch After Inversion. This method con- 
 sists in converting the starch into glucose and then estimating 
 the glucose with Fehling's solution. 
 
 Two gms. of the starch are mixed with water, and boiled 
 
 * Trans. Brit. Ph. Conf., 1903, 568-9. 
 
292 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 for 15 minutes; the solution is then cooled, and diluted to about 
 200 cc. Eighteen cc. of strong hydrochlroic acid are then 
 added, and the solution heated under a return condenser for 
 two or three hours. The solution is then cooled, neutralized 
 with potassium hydroxid and diluted to 250 cc. This is put 
 intO'a burette and titrated into 10 cc. of Fehling's solution, as 
 described above, under reducing sugars. 
 
 In estimating the starch in baking powder, 2 to 5 gms. 
 of the powder are introduced into an Erlenmeyer flask, 150 
 to 200 cc. of a 4 per cent solution of hydrochloric acid are 
 added and the solution gently boiled for four hours, after 
 which the flask and contents g,re cooled, neutralized by adding 
 sodium hydroxid, and made up to a definite volume. It is 
 then ready for testing with Fehling's solution. 
 
 Ten cc. of Fehling's solution is the equivalent of 0.0475 
 gm. of invert starch. 
 
 Estimation of Starch after Inversion by Means in Diastase. 
 The treatment of starch with malt infusion or pure diastase 
 at a temperature not above 70° C, readily converts the starch 
 into maltose, but the solution also contains, besides maltose, 
 various dextrins in proportions varying with the temperature 
 at which the diastase acts. The digestion may vary from 
 fifteen minutes to fifteen hours. Complete conversion of the 
 starch may be determined by testing occasionally with iodin. 
 A blank experiment should be made, especially if the digestion 
 is carried on beyond half an hour. A like quantity of the 
 same diastase solution should be digested at the same tem- 
 perature and for the same time, and the amount of sugar 
 found deducted from the total quantity found in the analysis. 
 Faulenbach* makes use of the following solution of diastase: 
 Crush 3.5 kilos of fresh green malt, treat with a mixture of 
 two liters of water and four liters of glycerin and let stand 
 
 * Zeitschr. f. physiol. Chem., VII, 510; and Chem. Centralh., 1883, 632. 
 
ESTIMATION OF SUGARS 293 
 
 for one week, stirring occasionally; then express and filter. 
 
 This solution is very stable. Five drops of it will dissolve 
 
 1 gm. of starch; 15 drops of it contain a quantity of carbo- 
 hydrate =0.001 gm. of glucose. 
 
 A quantity of the substance to be tested (containing about 
 
 2 gms. of starch) is boiled to gelatinize the starch. Fifteen 
 drops of the diastase solution are then added, and the mix- 
 ture digested at 63° C. It is then filtered to separate the 
 undissolved cellulose, etc., and heated with 20 cc. of hydro- 
 chloric acid in a water-bath for three hours. The acidity 
 is then just destroyed by means of caustic soda, the glucose 
 determined, o.ooi gm. deducted, and the starch calculated 
 from the glucose. 
 
 O'Sullivan * employs pure diastase, prepared as follows : 
 Pour sufficient water over 2 or 3 kilos of finely crushed 
 pale malt to just cover it. Let stand for three or four hours, 
 then express and filter the solution. Add alcohol (sp.gr. 0.83) 
 imtil the liquid above the flocculent precipitate becomes opal- 
 escent. Collect the precipitate, wash it with alcohol (sp.gr. 
 0.86 to 0.88) then with absolute alcohol and press it between 
 linen. Finally dry it in a vacuum over sulphuric acid. 
 
 Estimation of the Diastasic Value of Malt Extract. The 
 diastasic value of malt extracts may also be determined by 
 estimating the amount of maltose produced by a given amount 
 of the extract in a given time, when brought in contact with 
 an excess of gelatinized starch solution. It is always necessary 
 to estimate the copper-reducing power of the extracts with 
 Fehling's solution upon a separate sample, and to deduct this 
 from the total reducing power found after treatment with 
 the starch. 
 
 The process is briefly as follows: A definite quantity (say 
 30 cc.) of gelatinized starch, made from the best Bermuda 
 
 * Jour. Chem. Soc, XLV (1884), i. 
 
294 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 arrowroot, of 3 per cent strength, is placed in a flask and 
 heated to 40° C. A weighed amount of the malt extract (say 
 20 cc. of a I per cent solution) is then added, and the temper- 
 ature kept at 40° C. for exactly half an hour. At the end of 
 this time some sodium hydroxid solution is added in order 
 to check the further action of the diastase upon the starch 
 (heating to 100° C. accomplishes the same result). . The solution 
 is then diluted to a definite volume with water and the maltose 
 produced, estimated by Fehling's solution in the usual manner. 
 The quantity of reducing sugar originally present in the sample 
 must be previously determined, and this amount deducted 
 from the total amoimt found after treatment with starch, 
 and the remainder calculated as maltose. Ten cc. of Fehling's 
 solution is the equivalent of 0.0806 gm. of maltose. 
 
CHAPTER XVI 
 
 ESTIMATION OF FORMALDEHYDE* 
 
 The Ammonia Method (Legler). This method is based 
 upon the reaction between free ammonia and formaldehyde 
 in which hexamethylene-tetramin is formed. It is for ordinary 
 purposes sufl&ciently accurate. 
 
 It is this method which is recommended by Lederle f 
 for use in the laboratory of the New York City Health Depart- 
 ment, and by Prescott in the laboratory of the Universtiy of 
 Michigan. 
 
 The assay is conducted as follows: 2 cc. of the solution 
 
 are placed in a glass-stoppered bottle, the stopper of which 
 
 N 
 is thickly coated with petrolatum, and 50 cc. of — ammonia 
 
 solution added; let stand twelve hours, shaking occasionally. 
 
 ■ , . . ., N 
 Then determme the excess of ammonia by titratmg with — 
 
 sulphuric acid solution, using rosolic acid or litmus as 
 indicator. The excess of ammonia subtracted from the 
 quantity added gives the quantity which combined with the 
 formaldehyde, and thus the amoimt of the latter is ascer- 
 tained. 
 
 * Berichte d. Chem. Ges., XVI., 1335, 1883. 
 t Am. Drug., 1897, 246. 
 
 295 
 
296 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 The reaction is represented as follows: 
 
 6CH2O + 4NH3 = (CH2)6N4 + 6H2O. 
 4)180 4 )68.04 Hexamethylenetetramin 
 
 " ^5 2)17^ j^ 
 
 22.5 gms. 8.5 gms. or 1000 cc. — V.S. 
 
 0.0225 gm. 0.0085 gm. or i cc. " " 
 
 N 
 Assuming that 22 cc. of — sulphuric acid were employed 
 
 N 
 in the titration, 22 cc. of the — ammonia solution must have 
 
 2 
 
 been in excess, hence 28 cc. of the latter went into combina- 
 tion with the formaldehyde. Thus the 2 cc. of formaldehyde 
 solution contained 28X0.0225=0.63 gm. 
 
 A. G. Craig* says that the chief difficulty in using the 
 Legler method is the volatility of the ammonia. The difficulty 
 is not so much the loss of strength in the standard solution, 
 but the loss during the determination. He proposes the follow- 
 ing scheme by which this error is removed. 
 
 Prepare a normal solution of sulphuric acid. Make up an 
 approximately normal solution of ammonia, the exact strength 
 being immaterial. Procure several three-ounce prescription 
 bottles with smooth sides and close-fitting soft rubber stoppers. 
 Prepare a methyl orange solution. Procure a boiler in which 
 the bottles may be immersed to the neck without upsetting 
 (a large beaker will do). Take as much of the sample as 
 will contain 0.5 gm. of formaldehyde. Measure with the 
 pipette, 25 cc. of the ammonia solution into each of the bottles, 
 and to half of them add a sample of formaldehyde; stopper 
 tightly. If the necks of the bottles are small, the stoppers 
 need not be tied down. Place the bottles in the boiler, add 
 cold water to the necks, and heat to boiling. Boil for one 
 
 * J. A. C. S., XXIII, 642 (igoi). 
 
ESTIMATION OF FORMALDEHYDE 297 
 
 hour, and cool by running in cold water slowly, being careful 
 not to allow the cold water to touch the hot bottles. Titrate 
 with sulphuric acid and methyl orange to the first indication 
 of a color change. Take the difference between the readings 
 for the blanks and those for the samples, as the ammonia 
 consumed in normal cc. Of this difference, i cc. =0.0601 
 gm. of formaldehyde. 
 
 The Legler method is also liable to error because the 
 compoimd formed is a weak base, and as such combines 
 with acid, while at the same time it is liable to decompose 
 into ammonia and formaldehyde and thus give an indefinite 
 end-point when the residual ammonia is titrated with acid. 
 Error is also liable to be introduced, through the presence of 
 carbonic acid in the ammonia water, which, with the indicator 
 rosolic acid, gives no sharp end-reaction. 
 
 The Ammonixxm Chlorid Method. In this method a solu- 
 tion of ammonium chlorid is used, from which ammonium 
 is evolved by treatment with sodium hydroxid. The excess of 
 alkali is then determined by titration with standard solution 
 of sulphuric acid. This method, as devised by H. Schifif,* 
 and Modified by C. A. Male,f is as follows: 
 
 Introduce 2 gms. neutral ammonium chlorid, dissolved in 
 20 cc. of water, into a flask or bottle of about 200-cc. capacity, 
 having a well-fitting stopper. Dilute 10 cc. of the formaldehyde 
 solution to 100 cc. with water, and neutralize with sodium 
 hydroxid solution, as the formaldehyde solution generally con- 
 tains varying quantities of formic acid. Add 20 cc. of this 
 neutralized solution to the ammonium chlorid solution, then 
 
 N 
 25 cc. of — NaOH, and immediately stopper the flask, and 
 
 * Chem. Ztg., XXVII, 14 (1903)- 
 
 t Pharm. Jour., June, 1905, 844. See also Carl E. Smith, A. J. Ph., 
 LXX, 86, (1898). 
 
29S THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 leave for one hour. Afterwards determine the excess of alkali 
 
 N 
 with — H2SO4, using rosolic acid or litmus solution as indi- 
 cator, both of which give up a sharp end-reaction. The 
 reaction and calculation is based upon the following equation : 
 
 2NH4CI-F3CH20-h2KOH = N2(CH2)3 + 2KCH5H20. 
 
 N 
 I cc. — KOH is equivalent to 0.045 S^- of formic aldehyde. 
 
 This modified method is quite as simple, and gives results 
 practically identical with that of Romijn. 
 
 Oxidation by Hydrogen Dioxid (Blank and Finkenbeiner) * 
 This method depends upon the use of hydrogen dioxid for 
 oxidizing formaldehyde into formic acid, in alkaline solution. 
 The formic acid so produced neutralizes a portion of the alkali, 
 and the excess of the latter is then determined by titration 
 with standard acid. The method gives good results and can 
 be very rapidly carried out. 
 
 Three cc. of the formaldehyde solution are placed into a 100- 
 cc. Erlenmeyer flask, the latter closed with a well-fitting stopper, 
 and the weight of the solution carefully taken. 50 cc. of 
 normal NaOH V.S. are then added, and followed immediately, 
 but slowly, through a small funnel, by 50 cc. of a 3 per cent 
 hydrogen dioxid solution, previously neutralized with normal 
 NaOH V.S., using a drop of litmus solution. The solution is 
 allowed to stand about 15 minutes, or until the reaction has 
 ceased. The funnel and the sides of the vessel are then 
 rinsed with distilled water, a few drops of litmus solution 
 added and the unconsumed alkali titrated with normal H2SO4 
 V.S. The cc. of the latter deducted from the 50 cc. of nor- 
 mal NaOH V.S. originally taken gives the quantity of the 
 
 * Berichte d. Chem. Ges., XXXI, 2979 (1898); and A. J. Ph., 1899, 486. 
 
ESTIMATION OF FORMALDEHYDE 299 
 
 alkali V.S. which was consumed in the reaction. The dif- 
 ference multiplied by 0.03 gm. gives the weight of CH2O 
 in the sample. To find per cent multiply the result by 100 
 and divide by the weight of sample taken for analysis. 
 
 In the original process, double, normal sodium hydroxid is 
 used.* 
 
 The reaction is as follows: 
 
 CH20 + H202+NaOH = NaCOOH + 2H20 
 
 H2O2 oxidizes formaldehyde to formic acid, HCOOH, which 
 reacts with and neutralizes NaOH. 
 
 The lodometric Method (Romijn).t This method, which 
 is considered the most rapid, most accurate, and most readily 
 applied, depends upon the fact that iodin in the presence of 
 an alkali acts as an indirect oxidizing agent, giving, when 
 formaldehyde is present, the iodid of the base and formic acid: 
 
 CH20 + l2 + 2NaOH = 2NaI-t-CHOOH+H20. 
 
 As modified by L. Renter % it is as follows: 
 Twenty cc. of 35 to 40 per cent formaldehyde are intro- 
 duced into a graduated flask and distilled water added to bring 
 up the volume to 500 cc. Of this thoroughly mixed fluid, 5 
 cc. are introduced into a bottle capable of being perfectly 
 closed. 30 cc. of normal sodium or potassium hydroxid are 
 
 N 
 added, and then — iodin solution allowed to flow in from 
 
 a burette with constant agitation, until the fluid remains 
 a bright yellow color (36-60 cc). The shaking is then vig- 
 orously continued for one minute, when 40 cc. of normal 
 
 * See C. Allen Lyford, J. A. C. S., XXIX, 1227 (1907) "The Action of 
 Barium Peroxid and Hydrogen Peroxid upon Formaldehyde." 
 
 tZeitschr. anal. Chem., XXXVI, 18-24 (1897); ibid., XXXIX, 60-O3 
 (1900). 
 
 + Ph. Rev., 1903, 207. 
 
300 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 sulphuric acid are added, and then the excess of iodin titrated 
 
 N 
 with decinormal thiosulphate solution. Each cc. of — iodin 
 
 solution consumed corresponds to 0.0015 g™- formaldehyde. 
 A blank titration should always be made. 
 
 A solution containing as much as 5 per cent may be 
 accurately estimated by this method, provided not more than 
 2 gms. be taken. The method is, however, not accurate in 
 the presence of other aldehydes. F. O. Taylor,* commenting 
 upon this method, says that it is quite satisfactory, but that 
 the quantities of formaldehyde and reagents used are unnec- 
 essarily large and cumbersome, and the method is hence modi- 
 fied by him as follows: 
 
 From a weighing bottle, consisting of a small Erlenmeyer 
 
 flask, fitted with a perforated rubber stopper through which 
 
 passes a dropper, and containing about 25 or 30 cc. of the 
 
 formaldehyde solution, weigh out accurately about 10 gms. 
 
 of the solution into a stoppered 500-cc. flask and fill this to 
 
 the mark with distilled water. For titration remove 5 cc. of 
 
 this solution, corresponding to o.oi gm. of the weighed quan 
 
 tity of formaldehyde, and put into a 200-cc. Erlenmeyer flask. 
 
 Into another flask put 5 cc. of water for a blank titration. 
 
 To both now add 20 cc. of normal NaOH and then 20 cc. 
 
 N . 
 of an opproximately — iodin solution, whose exact strength 
 
 need not be loiown, and let stand for five or ten minutes for 
 
 the entire completion of the reaction. Now add 25 cc. of 
 
 N 
 normal H2SO4 and titrate the excess of iodin with — Na2S203. 
 
 10 
 
 The difference between the cc. of thiosulphate used on the 
 
 N 
 assay and the cc. of the blank is the number of cc. of — iodin 
 
 ID 
 
 * Bull. Ph., Aug., 1903, 323. 
 
ESTIMATION OF I'ORMALDlCIIVnE 301 
 
 N 
 consumed by the formaldehyde. Each cc. of — iodin so 
 
 used equals 0.0015 S^- oi CH2O. 
 
 The Potassium Cyanid Method.* This method is especially 
 applicable to solutions containing small quantities of formal- 
 dehyde. It depends upon the fact that potassium cyanid and 
 formaldehyde combine to form an addition product, in which 
 one molecule of potassium cyanid combines with one molecule 
 of formaldehyde, as shown by the following equation : 
 
 H 
 
 CHoO + KCN = N=C— C— O— K. 
 
 H 
 
 In the estimation, the formaldehyde is mixed with a known 
 
 quantity of potassium cyanid (in excess), the excess of the 
 
 latter being determined by the use of standard silver nitr'Ete 
 
 solution, and thus the quantity of potassium cyanid, which 
 
 combined with formaldehyde is found, and from this the 
 
 quantity of formaldehyde is calculated. 
 
 N 
 The process is carried out as follows: f Ten cc. of — silver 
 
 10 
 
 nitrate are treated with six drops of 50 per cent nitric acid 
 
 in a so-cc. flask. Ten cc. of a solution of potassium cyanid 
 
 (containing i gm. of KCN in 500 cc. of water) are then added 
 
 and well shaken. An aliquot part of the filtrate, say 25 cc, 
 
 N 
 is then titrated by Volhard's method with — ammonium 
 ■' 10 
 
 sulphocyanate for excess of silver. 
 
 Another 10 cc. of silver nitrate solution is then acidified 
 
 * Romijn, Zeitschr. anal. Chem., XXXVI, 18-24 (1897). 
 t Bernard H. Smith, J. A. C. S., XXV, 1032 (1903). 
 
302 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 with nitric acid and treated with lo cc. of the potassium 
 
 cyanid solution to which has been added a measured quantity 
 
 of dilute formaldehyde solution. The whole is made up to 
 
 50 cc. and then filtered. 25 cc. of the filtrate are titrated 
 
 N 
 with — ammonium sulphocyanate for excess of silver as 
 
 before. 
 
 The difference between these two results multiplied by 
 
 2 gives the amount of potassium cyanid which was used by 
 
 N 
 the formaldehyde, in terms of — sulphocyanate. 
 
 N 
 Each cc. of — V.S. represents 0.003 g™- oi formaldehyde. 
 
 The best results by this method are obtained, if the solu- 
 tion of formadlehyde is diluted to below i per cent. With 
 I per cent solutions it is necessary to use 15 cc. of the silver 
 solution. In estimating very dilute solutions, it is advisable 
 to use a 200-cc. flask and to take 100 cc. of the filtrate for 
 the titration. It is possible by this method to determine 
 with accuracy one part in 100,000. 
 
CHAPTER XVII 
 
 ESTIMATION OF ALCOHOL IN TINCTURES AND 
 BEVERAGES 
 
 The quantity of alcohol contained in dilute spirit, which 
 leaves no residue upon evaporation, may be ascertained by 
 taking the sp.gr. and referring to the alcohol table. When 
 taking the specific gravity, the temperature of the liquid 
 should be 15!° C. (60° F.). 
 
 In Wines, Beer, Tinctures and other alcoholic liquids 
 containing vegetable matter, the sp.gr. of the sample is taken 
 at i5f° C. (60° F.) and noted. A certain quantity (say 
 100 cc.) is measured off and evaporated to one half, or until 
 all odor of alcohol has passed off, the evaporation being con- 
 ducted without ebullition, in order that particles of the material 
 may not be carried off by the steam. The liquid left is 
 then diluted with distilled water, cooled to 60° F. and made 
 up to the original volume (100 cc), and the sp.gr. taken. 
 Lastly, we calculate: the sp.gr. before evaporating is divided 
 by the sp.gr. after evaporating, and the quotient will be the 
 sp.gr. of the water and alcohol only of the liquor. Then 
 by referring to the alcohol table the percentage of alcohol 
 contained in the liquor is obtained. 
 
 Example. The liquor before evaporating had a sp.gr. 
 of 0.9951; after evaporation and dilution to 100 cc. the sp.gr. 
 was found to be 1.0081. 
 
 -^^55 =Qggy^ tlie sp.gr. of the contained spirit. 
 1. 008 1 
 
 303 
 
304 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 TABLE FOR ASCERTAINING THE PERCENTAGES RESPECTIVELY 
 OF ALCOHOL BY WEIGHT, BY VOLUME, AND AS PROOF 
 SPIRIT FROM THE SPECIFIC GRAVITY. 
 
 Condensed from the excellent Alcohol Tables of Mr. Hehncr in the 
 "Analyst," vol. v. pp. 43-63. 
 
 Specific 
 Gravity 
 
 Absolute 
 
 Alcohol 
 
 by w'ght. 
 
 Per cent. 
 
 Absolute 
 
 Alcohol 
 
 by vol'me. 
 
 Per cent. 
 
 Proof S 
 Spirit. G 
 Per cent. 
 
 pecific 
 ravity 
 15-5'- 
 
 Absolute 
 
 Alcohol 
 
 by w'ght- 
 
 Per cent. 
 
 Absolute 
 
 Alcohol 
 
 bv vol'me. 
 
 Per cent. 
 
 Proof 
 
 Spirit- 
 
 Per cent. 
 
 I . 0000 
 
 0.00 
 
 0-00 
 
 0.00 
 
 9489 
 
 35.0s 
 
 41.90 
 
 73-43 
 
 
 9999 
 
 0.05 
 
 0.07 
 
 0. 12 
 
 9479 
 
 35-55 
 
 42.4s 
 
 74-39 
 
 
 9989 
 
 0.58 
 
 0.73 
 
 1. 28 
 
 9469 
 
 36.06 
 
 43.01 
 
 75-37 
 
 
 9979 
 
 1 . 12 
 
 1.42 
 
 2.4» 
 
 9459 
 
 36-61 
 
 43.63 
 
 76.45 
 
 
 9969 
 
 1-75 
 
 2 . 20 
 
 3-8s 
 
 9449 
 
 37-17 
 
 44.24 
 
 77-53 
 
 
 99S9 
 
 2-33 
 
 2-93 
 
 S-13 
 
 9439 
 
 37-72 
 
 44.86 
 
 78.61 
 
 
 9949 
 
 2.89 
 
 3-62 
 
 6-34 
 
 9429 
 
 38-28 
 
 4S.47 
 
 79-68 
 
 
 9939 
 
 3-47 
 
 4-34 
 
 7-61 
 
 9419 
 
 38-83 
 
 46-08 
 
 80.75 
 
 
 9929 
 
 4 .06 
 
 5-08 
 
 8-90 
 
 9409 
 
 39-35 
 
 46.64 
 
 81.74 
 
 
 9919 
 
 4.69 
 
 5.86 
 
 10.26 
 
 9399 
 
 39-85 
 
 47.18 
 
 82.69 
 
 
 9909 
 
 S-3I 
 
 6.63 
 
 11-62 
 
 9389 
 
 40.35 
 
 47.72 
 
 83.64 
 
 
 9899 
 
 5-94 
 
 7-40 
 
 12.97 
 
 93 79 
 
 40-85 
 
 48-26 
 
 84.58 
 
 
 9889 
 
 6.64 
 
 8.27 
 
 14-50 
 
 9369 
 
 41-35 
 
 48.80 
 
 85-53 
 
 
 9879 
 
 7.33 
 
 9-13 
 
 15-99 
 
 9359 
 
 41.85 
 
 49-34 
 
 86.47 
 
 
 9869 
 
 8.00 
 
 9-95 
 
 17-43 
 
 9349 
 
 42-33 
 
 49-86 
 
 87.37 
 
 
 9859 
 
 8.71 
 
 10.82 
 
 18-96 
 
 9339 
 
 42.81 
 
 50-37 
 
 88.26 
 
 
 9849 
 
 9-43 
 
 II . 70 
 
 20.50 
 
 9329 
 
 43.29 
 
 50-87 
 
 89.15 
 
 
 9839 
 
 10.15 
 
 12.58 
 
 22.06 
 
 9319 
 
 43.76 
 
 SI. 38 
 
 90.03 
 
 
 9829 
 
 10.92 
 
 13-52 
 
 23.70 
 
 9309 
 
 44.23 
 
 SI. 87 
 
 90.89 
 
 
 9S19 
 
 II .69 
 
 14.46 
 
 25-34 
 
 9299 
 
 44.68 
 
 52.34 
 
 91.73 
 
 
 9809 
 
 12.46 
 
 lS-40 
 
 26.99 
 
 9289 
 
 45.14 
 
 52.82 
 
 92.56 
 
 
 9799 
 
 13 23 
 
 16.33 
 
 28.62 
 
 9279 
 
 45.59 
 
 53.29 
 
 93.39 
 
 
 9789 
 
 14.00 
 
 17.26 
 
 30.26 
 
 9269 
 
 46.05 
 
 53-77 
 
 94.22 
 
 
 9779 
 
 14.91 
 
 18.36 
 
 32-19 
 
 9259 
 
 46.50 
 
 54.24 
 
 95.05 
 
 
 9769 
 
 IS-7S 
 
 19-39 
 
 33-96 
 
 9249 
 
 46.96 
 
 54.71 
 
 95.88 
 
 
 9759 
 
 16.54 
 
 20.33 
 
 35.63 
 
 9239 
 
 47.41 
 
 55.18 
 
 96.70 
 
 
 9749 
 
 17-33 
 
 21.29 
 
 37-30 
 
 9229 
 
 47.86 
 
 55.65 
 
 97.52 
 
 
 9739 
 
 18.15 
 
 22.27 
 
 39-03 
 
 9219 
 
 48.32 
 
 56. 11 
 
 98.34 
 
 
 9729 
 
 18.92 
 
 23.10 
 
 40-64 
 
 9209 
 
 48.77 
 
 56.58 
 
 99.16 
 
 
 9719 
 9709 
 
 I9-7S 
 20.58 
 
 24.18 
 25-17 
 
 42-38 
 
 9199 
 
 49.20 
 
 57.02 
 
 99 . 93 
 
 
 44-12 
 
 
 
 
 
 
 9699 
 9689 
 
 21.38 
 22 .15 
 
 26-13 
 27-04 
 
 45-79 
 47-39 
 
 9198 
 
 49.24 
 
 57. 06 
 
 loo.ooPs 
 
 
 9679 
 9669 
 
 22 . 92 
 
 23 69 
 
 27-95 
 28-86 
 
 48.98 
 
 
 
 
 
 
 50.57 
 
 9189 
 
 49.68 
 
 57.49 
 
 100. .76 
 
 
 9659 
 
 24.46 
 
 29.76 
 
 52.16 
 
 9179 
 
 50.13 
 
 57.97 
 
 101.59 
 
 
 9649 
 
 25.21 
 
 30.65 
 
 53-71 
 
 9169 
 
 50.57 
 
 58.41 
 
 102.35 
 
 
 9639 
 
 25-93 
 
 31-48 
 
 5S-I8 
 
 9159 
 
 51 .00 
 
 58-85 
 
 103.12 
 
 
 9629 
 
 26.60 
 
 32-27 
 
 56-55 
 
 9149 
 
 51.42 
 
 59.26 
 
 103.85 
 
 
 9619 
 
 27 . 29 
 
 33 -06 
 
 57-94 
 
 9139 
 
 51.83 
 
 59-68 
 
 104-58 
 
 
 9609 
 
 28.00 
 
 33-89 
 
 59-40 
 
 9129 
 
 52.27 
 
 60. 12 
 
 105.35 
 
 
 9599 
 
 28.62 
 
 34.61 
 
 60.66 
 
 9119 
 
 52.73 
 
 60.56 
 
 106.15 
 
 
 9589 
 
 29.27 
 
 35-35 
 
 61.95 
 
 9109 
 
 53.17 
 
 61 .02 
 
 106.93 
 
 
 9579 
 
 29-93 
 
 36-12 
 
 63-30 
 
 9099 
 
 53.61 
 
 61.45 
 
 107.69 
 
 
 9569 
 
 30.50 
 
 36-76 
 
 64-43 
 
 9089 
 
 54.05 
 
 61.88 
 
 108.45 
 
 
 9559 
 
 31 .06 
 
 37-41 
 
 65-55 
 
 9079 
 
 54.52 
 
 62.36 
 
 109.28 
 
 
 9549 
 
 31-69 
 
 38-11 
 
 66-80 
 
 9069 
 
 55.00 
 
 62.84 
 
 110.12 
 
 
 9539 
 
 32-31 
 
 38-82 
 
 68-04 
 
 9059 
 
 55.45 
 
 63-28 
 
 110.92 
 
 
 9529 
 
 32.94 
 
 39-54 
 
 69.29 
 
 9049 
 
 55.91 
 
 63-73 
 
 III .71 
 
 
 95J9 
 
 33-53 
 
 40-20 
 
 70.46 
 
 9039 
 
 56.36 
 
 64.18 
 
 112.49 
 
 
 9509 
 
 34- 10 
 
 40.84 
 
 71-58 
 
 9029 
 
 56.82 
 
 64.63 
 
 113 .26 
 
 ■9499 
 
 34-57 
 
 41 -37 
 
 72.50 
 
 9019 
 
 57.25 
 
 65-05 
 
 113.99 
 
ALCOHOL IN TINCTURES AND BEVERAGES 
 
 305 
 
 Specific 
 Gravity 
 
 Absolute 
 Alcohol 
 by w'ght. 
 Per cent. 
 
 Absolute 
 
 Alcohol 
 
 by vol'me. 
 
 Per cent. 
 
 Proof 
 
 Spirit. 
 
 Per cent. 
 
 Specific 
 
 Gravity 
 
 15.5°. 
 
 Absolute 
 
 Alcohol 
 
 by w'ght. 
 
 Per cent. 
 
 Absolute 
 Alcohol 
 yy vol'me. 
 Per cent. 
 
 Proof 
 
 Spirit. 
 
 Per cent. 
 
 .9009 
 
 57.67 
 
 65.45 
 
 114.69 
 
 .8429 
 
 82.19 
 
 87.27 
 
 152.95 
 
 .8999 
 
 58.09 
 
 65.85 
 
 115.41 
 
 .8419 
 
 82.58 
 
 87.58 
 87.88 
 
 153.48 
 
 .8989 
 
 58.55 
 
 66.29 
 
 116. 18 
 
 .8409 
 
 82.96 
 
 154.01 
 
 .8979 
 
 59.00 
 
 66.74 
 
 116.96 
 
 .8399 
 
 83.35 
 
 88.19 
 
 154.54 ■ 
 
 .8969 
 
 59-43 
 
 67.15 
 
 117.68 
 
 • 8389 
 
 83.73 
 
 88.49 
 
 155-07 
 
 ■ 8959 
 
 59.87 
 
 67.57 
 
 118. 41 
 
 .8379 
 
 84.12 
 
 88.79 
 
 155-61 
 
 .8949 
 
 60.29 
 
 67.97 
 
 119. 12 
 
 .8369 
 
 84.52 
 
 89.11 
 
 156. r6 
 
 ■ 8939 
 
 60.71 
 
 68.36 
 
 119.80 
 
 .8359 
 
 84.92 
 
 89.42 
 
 156.71 
 
 .8929 
 
 61. 13 
 
 68.76 
 
 120.49 
 
 .8349 
 
 85.31 
 
 89.72 
 
 157.24 
 
 .8919 
 
 61.54 
 
 69.15 
 
 121 . 18 
 
 .8339 
 
 85.69 
 
 90.02 
 
 157.76 
 
 .8909 
 
 61.96 
 
 69.54 
 
 1 2 1 . 86 
 
 .8329 
 
 86.06 
 
 90.32 
 
 158.28 
 
 .8899 
 
 62.41 
 
 69.96 
 
 122.61 
 
 .8319 
 
 86.46 
 
 90.61 
 
 158.79 
 
 .8889 
 
 62.86 
 
 70.40 
 
 123.36 
 
 .8309 
 
 86.85 
 
 90.90 
 
 i'!9.3i 
 
 .8879 
 
 63 30 
 
 70.81 
 
 124.09 
 
 .8299 
 
 87.23 
 
 91 .20 
 
 159.82 
 
 .8869 
 
 63.74 
 
 71.22 
 
 124.80 
 
 .8289 
 
 ll-^" 
 
 91.49 
 
 160.33 
 
 .8859 
 
 64.17 
 
 71 .62 
 
 125.51 
 
 .8279 
 
 88.00 
 
 91-78 
 
 160.84 
 
 .8849 
 
 64.61 
 
 72.02 
 
 126.22 
 
 .8269 
 
 88.40 
 
 92.08 
 
 161.37 
 
 ■ 8839 
 
 65.04 
 
 72.42 
 
 126.92 
 
 .8259 
 
 88.80 
 
 92.39 
 
 161 . 91 
 
 .8829 
 
 65.46 
 
 72.80 
 
 127-59 
 
 .8249 
 
 89.19 
 
 92.68 
 
 162.43 
 
 .8819 
 
 65.88 
 
 73-19 
 
 128-25 
 
 .8239 
 
 89.58 
 
 92.97 
 
 162.93 
 
 .8809 
 
 66.30 
 
 73-57 
 
 128-94 
 
 .8229 
 
 89.96 
 
 93.26 
 
 163.43 
 
 • 8799 
 
 66.74 
 
 73-97 
 
 129.64 
 
 .8219 
 
 90.32 
 
 93.52 
 
 163.88 
 
 .8789 
 
 67-17 
 
 74-37 
 
 130-33 
 
 .8209 
 
 90.68 
 
 93.77 
 
 164.33 
 
 .8779 
 
 67.58 
 
 74.74 
 
 130-98 
 
 .8199 
 
 91.04 
 
 94.03 
 
 164.78 
 
 ■ 8769 
 
 68.00 
 
 75-12 
 
 131-64 
 
 .8189 
 
 91.39 
 
 94.28 
 
 165.23 
 
 -8759 
 
 68.42 
 
 75-49 
 
 132-30 
 
 .8179 
 
 91.75 
 
 94.53 
 
 165.67 
 
 .8749 
 
 68.83 
 
 75.87 
 
 132.95 
 
 .8169 
 
 92.11 
 
 94.79 
 
 166.12 
 
 • 8739 
 
 69.2s 
 
 76.24 
 
 133.60 
 
 .8159 
 
 92.48 
 
 95.06 
 
 166.58 
 
 .8729 
 
 69.67 
 
 76.61 
 
 134.2s 
 
 .8149 
 
 92.85 
 
 95.32 
 
 167.04 
 
 • 8719 
 
 70-08 
 
 76.98 
 
 134.90 
 
 ■ 8139 
 
 93.22 
 
 95.58 
 
 167.50 
 
 .8709 
 
 70.48 
 
 77.32 
 
 135.51 
 
 .8129 
 
 93 ■ 59 
 
 98.54 
 
 167.96 
 
 .8699 
 
 70.88 
 
 77.67 
 
 136.13 
 
 .8119 
 
 93.96 
 
 96.11 
 
 168.24 
 
 .8689 
 
 71 .29 
 
 78.04 
 
 136.76 
 
 .8109 
 
 94.31 
 
 96.34 
 
 168.84 
 
 ■ 8679 
 
 71.71 
 
 78.40 
 
 137.40 
 
 .8099 
 
 94 66 
 
 96.57 
 
 169.24 
 
 .8669 
 
 72. 13 
 
 78.77 
 
 138.05 
 
 .8089 
 
 95.00 
 
 96.80 
 
 169.65 
 
 .'8659 
 
 72.57 
 
 79.16 
 
 138.72 
 
 .8079 
 
 95.36 
 
 97.05 
 
 170.07 
 
 .8649 
 
 73.00 
 
 79-54 
 
 139.39 
 
 .8069 
 
 95.71 
 
 97.29 
 
 170.50 
 
 • 8639 
 
 73.42 
 
 79.90 
 
 140.02 
 
 .8059 
 
 96.07 
 
 97.53 
 
 170.99 
 
 .8629 
 
 73.83 
 
 80.26 
 
 140.65 
 
 .8049 
 
 96.40 
 
 97. 75 
 
 171.30 
 
 .8619 
 
 74.27 
 
 80.64 
 
 141.33 
 
 .8039 
 
 96.73 
 
 97.96 
 
 171.68 
 
 .8609 
 
 74.73 
 
 8i .04 
 
 142.03 
 
 .8029 
 
 97.07 
 
 98.18 
 
 172.05 
 
 -8599 
 
 75. 18 
 
 81.44 
 
 142.73 
 
 .8019 
 
 97.40 
 
 98.39 
 
 172.43 
 
 .8589 
 
 75.64 
 
 81.84 
 
 143.42 
 
 .8009 
 
 97.73 
 
 98.61 
 
 172.80 
 
 .8579 
 
 76.08 
 
 82.23 
 82.58 
 
 144.10 
 
 .7999 
 
 98.06 
 
 98.82 
 
 173.17 
 
 ■ 8569 
 
 76.50 
 
 144.72 
 
 ■ 7989 
 
 98.37 
 
 99.00 
 
 173-50 
 
 -8559 
 
 76.92 
 
 82.93 
 
 145.34 
 
 .7979 
 
 98.69 
 
 99.18 
 
 173-84 
 
 -8549 
 
 77-33 
 
 83.28 
 
 145.96 
 
 .7969 
 
 99.00 
 
 99.37 
 
 174-17 
 
 .8535 
 
 77.75 
 
 83.64 
 83.98 
 
 146.57 
 
 .7959 
 
 99.32 
 
 99.57 
 
 174.52 
 
 .8529 
 
 78.16 
 
 147.17 
 
 .7949 
 
 99.65 
 
 99.77 
 
 174.87 
 
 .8519 
 
 78.56 
 
 84.31 
 
 147.75 
 
 .7939 
 
 99.97 
 
 99.98 
 
 175.22 
 
 .8509 
 
 78.96 
 79-36 
 
 84-64 
 84-97 
 
 148.32 
 148.90 
 
 
 
 
 
 -8499 
 
 
 
 
 
 .8489 
 
 79.76 
 
 85.29 
 
 149-44 
 
 ■^ 
 
 
 
 
 .8479 
 
 80.17 
 
 85.63 
 
 150.06 
 
 
 Absolute 
 
 Alcohol. 
 
 
 .8469 
 .8459 
 .S449 
 
 80.58 
 81.00 
 81 .40 
 
 §5-9' 
 86.32 
 86.64 
 
 150.67 
 151.27 
 151.83 
 
 
 
 
 
 
 
 
 
 .8439 
 
 81.80 
 
 86.96 
 
 152.40 
 
 .7938 
 
 100.00 
 
 100.00 
 
 175-25 
 
306 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Then by referring to the table we find that this sp.gr. cor- 
 responds to 7.33 per cent, by weight, of absolute alcohol. 
 
 Another Way is to boil the liquid in a retort, condense 
 the vapor, and when all the alcohol has passed over add 
 suflScient water to the distillate to make up the original vol- 
 ume, at the temperature of 15!° C. (60° F.). Then, by 
 taking the sp.gr. of this diluted distillate, the quantity of 
 absolute alcohol is found by reference to the table. This 
 latter method requires the taking of the sp.gr. but once and 
 gives more accurate results. 
 
PART III 
 
 A FEW GASOMETRIC ANALYSES 
 
 CHAPTER XVIII 
 
 THE NITROMETER 
 
 For general gas analysis, and for the rapid estimation of 
 such substances as ethyl nitrite, hydrogen dioxid, urea, bleaching 
 powder, manganese dioxid, etc., an instrument called the 
 nitrometer is used. 
 
 The apparatus in its simplest form is shown in Fig. 57. 
 It consists of a measuring tube, a, of 50- or 
 loo-cc. capacity, and graduated in tenths 
 of a cc. This is connected by means of a 
 stout rubber tube with an open equilibrium 
 tube, I, also called " control-tube," " pres- 
 sure-tube," or " level-tube." Both tubes are 
 preferably provided with a globular expan- 
 sion near the lower end, and are held by 
 suitable clamps upon a stand, in such a 
 manner that either tube may be readily and 
 quickly clamped at a higher or lower level. 
 The measuring tube is fitted at the top 
 with a stop-cock, c, and a graduated glass 
 tube or cup, d. Some nitrometers are pro- 
 vided with a three-way stop-cock, so 
 arranged that according to the way it is 
 turned, it will discharge the contents of the 
 cup either into the measuring tube below, or out into the 
 
 Fig. 57. 
 
 307 
 
308 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 waste opening which is usually placed at e, or it will discharge 
 the contents of the measuring tube into the waste opening. 
 
 With this apparatus gases can be rapidly and accurately 
 measured at definite temperature and pressure. 
 
 In measuring the gas the instrument is filled with some 
 liquid in which the gas is insoluble — generally mercury. In 
 many cases a saturated solution or salt may be used. 
 
 Suppose we fill the instrument with mercury in such quan- 
 tity that when the stop-cock is opened and the control-tube 
 raised, the mercury will rise as far as the top, and about two 
 inches in the control-tube. 
 
 The top is now closed, the control-tube lowered, and a 
 little carbonic acid gas admitted 'through e. The top is 
 then again closed, and the instrument allowed to stand until 
 its contents have acquired the temperature of the room. A 
 centigrade thermometer suspended to the stand will then give 
 the temperature of the gas. 
 
 The control-tube is now raised or lowered so as to make 
 the level of the liquid in both tubes the same. This makes 
 the pressure in the tube the same as the atmospheric pressure 
 outside, and, by referring to a barometer standing near, this 
 pressure is ascertained. 
 
 We now have a definite volume of the gas at a known 
 temperature and pressure. 
 
 It now only remains to read off the volume of the gas, 
 and correct it to the normal temperature and pressure by 
 Charles' and Boyle's laws, respectively. 
 
 The normal temperature and pressure is o° C. and 760 
 mm. pressure. 
 
 The weight of the gas in grams may then be calculated 
 from its volume by multiplying the number of cc. at the nor- 
 mal temperature and pressure, by the weight of i cc. of the 
 gas in grams. 
 
THE NITROMETER 309 
 
 This weight may be found as follows: 
 
 looo cc. of hydrogen at normal temperature and pressure 
 weigh o.o8g6 gm. One cc. of H then weighs 0.0000896 gm. 
 
 One cc. of oxygen weighs sixteen times as much, and i 
 cc. of nitrogen weighs fourteen times as much. Therefore, 
 by multiplying the weight of i cc. of H by the atomic weight 
 of an elementary gas, or half the molecular weight of a com- 
 pound gas, the weight of i cc. of that gas is obtained. 
 
 According to the Law of Charles, the volume of a gas 
 under constant pressure varies directly with the absolute tem- 
 perature. 
 
 All gases expand or contract by -jt^ of their volume for 
 each centigrade degree of temperature increased or decreased. 
 
 We may regard a gas at 0° C. as having passed through 
 273° C. In other words, 273° below zero must be regarded 
 as the absolute zero, and 0° C. as 273° absolute temperature. 
 
 Thus the absolute temperature centigrade is the observed 
 temperature + 273°. 
 
 Example. A given volume of oxygen gas at 15° C. measures 
 20 cc. What will it measure at 0° C. ? 
 
 o°-|-273°X20 273° X 20 
 
 Boyle's Law. The volume of a confined gas is inversely 
 proportional to the pressure brought to bear upon it. That 
 is, the less the pressure, the greater the volume, and vice versa. 
 
 Rule. Multiply the observed volume by the observed pres- 
 sure, and divide by the normal pressure. 
 
 Example. A given volume of gas at 750 mm. pressure 
 measures 20 cc. What will it measure at 760 mm. (the nor- 
 mal pressure) ? 
 
 7 CO X 20 cc. . 
 
 — =10.73 cc. Ans. 
 
 760 
 
310 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 Now let us take an example in which both laws are involved. 
 
 A given volume of oxygen at 15° C. subjected to a pressure 
 of 750 mm. measures 20 cc. What will it measure at the 
 normal temperature and pressure? — i.e., 0° C. and 760 mm. 
 
 In the first example we find that 20 cc. of oxygen at 15° 
 C. will measure at 0° C. 18.95 cc Then 
 
 7150X18.05 cc. 
 
 — ^ = 18.70 cc. Ans. 
 
 760 ' 
 
 Now to find the weight of this volume of o.xygen we pro- 
 ceed as follows: 
 
 I cc. of H weighs 0.0000896 gm.; 
 
 I cc. of O weighs 16X0.0000896=^0.0014336 gm.; 
 
 18.70 cc. of 0=18.70X0.0014336 gm., or 0.02680832 gm. 
 
 The method of using the nitrometer for the gasometric 
 assay of various substances is illustrated in the assay of spirit 
 of nitrous ether. 
 
CHAPTER XIX 
 
 ASSAY OF NITRITES 
 
 Spirit of Nitrous Ether. This is an alcoholic solution of 
 ethyl nitrite (C2H5N02 = 75.05), yielding, when freshly pre- 
 pared and tested in the nitrometer, not less than 1 1 times its 
 own volume of nitrogen dioxid (NO = 30.01). 
 
 AVhen nitrites are mixed with an excess of KI and acid- 
 ulated with H2SO4, iodin is liberated, and all the nitrogen 
 of the nitrite is evolved in the form of NO, as shown in the 
 equation 
 
 2C2H5N02-F2KI+2H2S04=2C2H50H-H2KHS04 + l2 + 2NO. 
 150.1 60.02 
 
 The process is conducted as follows: 
 
 Open the stop-coct of the measuring tube, raise the control- • 
 tube, and pour into the latter a saturated solution of NaCl 
 until the measuring tube, including the bore of the stop-cock, 
 is completely filled. Then close the stop-cock and fix the 
 control-tube at a lower level. Now introduce into the funnel 
 at the top of the measuring tube a weighed quantity (about 
 4 gms.) * of spirit of nitrous ether; open the stop-cock and 
 allow the spirit to run into the nitrometer, being careful that 
 no air enters at the same time. Ten cc. of potassium iodid 
 
 * It is convenient to take 5 cc. accurately measured, and calculate its 
 weight by multiplying by the specific gravity. 
 
 3U 
 
312 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 T.S. are now added in the same manner, and followed by 
 lo cc. of normal sulphuric acid V.S. Effervescence takes 
 place immediately, and after thirty to sixty minutes, when 
 the volume of gas has become constant, the control-tube is 
 lowered so as to make the level of the liquid in both tubes 
 the same, and the volume of the gas in the graduated tube 
 read off. 
 
 This volume, multiplied by 0.003099 gm., gives the weight 
 of ethyl nitrite in the spirit taken for analysis. The product 
 multiplied by 100, and then divided by the weight of the spirit 
 taken, gives the per cent of pure ethyl nitrite present. 
 
 The temperature correction is one-third of one per cent 
 of the total percentage found, for each degree, additive, if 
 the temperature is below 25" C. The barometric correction 
 is -^ of one per cent for each millimeter, additive if above, 
 subtractive if below 760. 
 
 The volume of NO generated at the ordinary indoor tem- 
 perature (assumed to be at or near 25° C, 77° F.) should 
 not be less than 55 cc. if 5 cc. of the spirit are taken, corre- 
 sponding to about 4 per cent of pure ethyl nitrite. 
 
 Sodium chlorid solution is used in the above assay, because, 
 owing to its density, the spirit will float on top, and the gas 
 evolved will not dissolve in it. At the same time the expense 
 of using mercury is saved. It is important that no air be 
 allowed to get into the measuring tube, because this would 
 convert the NO into a higher oxid of nitrogen, which would 
 dissolve in the salt solution, and thus vitiate the result. 
 
 If it is required to correct the volume of gas evolved at 
 higher temperatures, to its corresponding volume at 0° C, 
 the calculations involved are as explained below: 
 
 Example. Five cc. of spirit of nitrous ether (sp.gr. 0.823) 
 are treated in a nitrometer, and the NO evolved measures 
 55 cc. 
 
ASSAY OF NITRITES 313 
 
 The temperature at which the operation is conducted is 
 25° C, and the atmospheric pressure normal. 
 
 What per cent of ethyl nitrite is present in the sample ? 
 
 By consulting the equation given above, it will be seen that 
 one molecular weight of NO = 30.01 is evolved from one 
 molecular weight of ethyl nitrite, 75.05. 
 
 Now reduce the volume of gas liberated at 25° C. to its 
 corresponding volume at 0° C. Thus 
 
 273° + 25":5S: :273° + o°:x. ^ = 50.4 cc. 
 
 Thus the gas evolved from 5 cc. of the spirit of nitrous ether, 
 measured at 0° C, is 50.4 cc. 
 
 The next step in the calculation is to find how much ethyl 
 nitrite each cc. of the evolved NO represents. One liter of 
 hydrogen at 0° C. and normal pressure weighs 0.0896 gm. 
 
 By multiplying this weight by half the molecular weight 
 of NO, the weight of 1000 cc. of the latter gas is obtained; 
 this will be found to be 1.344. Now if 1.344 gms. of NO 
 measures 1000 cc, 30.01 gms. will measure 22,328.86 cc. 
 
 1.344:1000: : 30.01 :x. x=22,328.86. 
 
 Then if 22,328.86 cc. of NO are evolved by, and conse- 
 quently represent, 75.05 gms. of ethyl nitrite, as the equation 
 shows, I cc. of NO will represent 0.0033618 gm. of pure 
 ethyl nitrite. 
 
 Now, since in the above example 50.4 cc. of gas were 
 evolved at 0° C, the 5 cc. of spirit of nitrous ether examined 
 must contain 
 
 50.4X0.003618 gm. =0.1694 gm. 
 
 of pure ethyl nitrite. 
 
314 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 In order to determine the percentage strength, the weight 
 of the spirit taken must be known. This may be found by 
 multiplying the measure by the specific gravity, 5 cc.X0.823 
 = 4.115 gms. Then 
 
 4.115 gms.:o.i6943 gm.: :ioo:x. x^4 per cent. 
 
 Amyl Nitrite is a liquid containing about 80 per cent of 
 amyl nitrite (principally iso-amyl nitrite), C5HiiN02= 117.09, 
 together with variable quantities of undetermined compounds. 
 
 The assay is as follows: Take about 3 cc. of the amyl 
 
 nitrite, which has been previously shaken with 0.5 gm. of 
 
 potassium bicarbonate and carefully decanted. Put into a 
 
 tared loo-cc. measuring flask, and weigh it accurately; add 
 
 alcohol to bring the volume to exactly 100 cc. Ten cc. of this 
 
 alcoholic solution are introduced into the nitrometer as directed 
 
 for spirit of nitrous ether; 10 cc. of potassium iodid solution 
 
 N 
 (20 per cent) and 10 cc. of — H2SO4 V.S. are then added, 
 
 and the volume of NO generated, measured at the ordinary 
 indoor temperature (25° C. or 77° F.), should be about 40 cc. 
 Each cc. at this temperature represents 0.0048 gm. of pure 
 amyl nitrite. 
 
 Sodium Nitrite (NaN02 = 69.01). This, like the other 
 nitrites mentioned, when treated with potassium iodid and 
 sulphuric acid, is decomposed, and NO is given off. The 
 reaction is here illustrated: 
 
 2NaN02 + 2KI + 2H2S04 
 
 = K2S04 + Na2S04 + 2H20+2NO+l2. 
 
 A molecule of NaN02 (69.01) evolves, when properly 
 treated, one molecule of NO (30.01). 
 
ASSAY OF NITRITES 315 
 
 The assay process is as follows: Weigh out 0.15 gm. of 
 
 NaNOa, dissolve it in about 5 cc. of water, and introduce 
 
 the solution into a nitrometer. This is followed by a solution 
 
 N 
 of I gm. of KI in 6 cc. of water and 15 cc. of — H2SO4. 
 
 The gas which is liberated should measure not less than 50 
 cc. at 15° C. (59° F.) or 51.7 cc. at 25° C. (77° F.), corresponding 
 to not less than 97.6 per cent of the pure salt. Each cc. at 
 25° C. represents 0.002837 E^- ^^'^ ^^ 0° C. 0.0030916 gm. 
 of pure NaNO... 
 
 Nitric Acid in Nitrates. This may also be effected by the 
 use of the nitrometer. 
 
 When a nitrate is shaken up with an excess of sulphuric 
 acid and mercury, the nitrate is decomposed and NO is evolved, 
 as seen- in the following equation: 
 
 2KNO3 +4H2SO4 + 3Hg = 3HgS04 + K2SO4 + 2NO + 4H2O. 
 
 2)202.22 ■ 2 )60.02 
 
 loi.ii 3°-°i 
 
 Thus each molecule of the nitrate radical NO3 gives off a 
 molecule of NO. 
 
 Not more than 0.2 gm. of nitrate should be taken for 
 analysis, since, if this quantity is exceeded, the volume of 
 gas evolved will! be greater than the instrument can conve- 
 niently hold. In this estimation the nitrometer is filled with 
 mercury instead of brine; the nitrate is dissolved in 5 cc. 
 of water, introduced into the nitrometer, and followed by 
 excess of strong sulphuric acid. The instrument is well 
 shaken for some time, and when action has ceased and the 
 contents have cooled down to the temperature of the room, 
 the level is adjusted and the volume of NO read off and cal- 
 culated in the usual way. 
 
316 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 ESTIMATION OF NITROGEN DIOXID 
 
 NO = 30.01; I liter at 0° C. and 760 mm. = 1.344 gms. 
 at 25° C. and 760 mm. = 1.2394 gms. 
 
 ONE CUBIC CENTIMETER OF NITROGEN DIOXID IS THE 
 EQUIVALENT OF: 
 
 Nitrogen dioxid, NO =30.01 
 
 Amyl nitrite, C5H,iN02= 117.09 
 Ethyl nitrite, C2H5NOj = 75.os . 
 Sodium nitrite, NaNOj^bg.oi. . 
 
 At 0° C. and 
 760 mm. 
 Grara. 
 
 0.001344 
 
 0-005245 
 
 0.0033618 
 
 0.0030916 
 
 At 2S° C. anJ 
 
 760 mm. 
 
 Gram. 
 
 0.0012394 
 0.0048 
 u. 003099 
 0.002837 
 
CHAPTER XX 
 HYDROGEN DIOXID 
 
 As stated in a previous chapter, hydrogen dioxid when 
 acted upon by an acidulated solution of potassium perman- 
 ganate, is decomposed and oxygen is evolved. One-half of 
 this oxygen comes from the dioxid and the other half from 
 the permanganate. 
 
 Therefore if i cc. of the dioxid be treated in this way 
 and 20 cc. of oxygen are evolved, the strength of the solu- 
 tion is ten volumes. 
 
 The nitrometer may be used for this estimation. 
 
 This instrument is charged with a concentrated solution 
 of sodium sulphate (which in this case is better than brine), 
 and I cc. of the dioxid introduced from the funnel, followed 
 by excess of solution of permanganate acidulated with sul- 
 phuric acid. 
 
 This latter solution should be of such strength that when 
 the reaction is completed, the solution should still have a 
 purple color. 
 
 The reaction is thus illustrated: 
 
 5H202+3H2S04-f2KMn04=K2S04+2MnS04+8H20-|-s02. 
 
 By the use of Squibb's Urea Apparatus (Fig. 61) the 
 estimation may be easily and rapidly made. 
 
 Into the generating bottle is put about 30 cc. of a strong, 
 acidulated solution of potassium permanganate, and a small 
 test-tube containing i cc. of H2O2 is carefully introduced. 
 The two liquids must not be allowed to come in contact. 
 
 317 
 
318 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The larger flask is filled with water, or better, a solution 
 of sodium sulphate; the connection is then made by means 
 of the rubber tube, and the generating-bottle tipped over and 
 agitated so that the liquids will mix and the reaction take 
 place. 
 
 The liberated oxygen then passes into the larger bottle, 
 displacing an equal volume of water, which is collected and 
 measured. Half of this volume represents the volume strength 
 of the H2O2. 
 
 An Improvised Nitrometer may be used. The author has 
 found the following instrument convenient: 
 
 To the bottom of an ordinary 50-cc. burette is attached 
 a suitable length of rubber tubing, to the other end of which 
 is attached another burette or ungraduated tube, which serves 
 as a control-tube. 
 
 Into the top of the burette is fitted a rubber stopper, 
 through which passes a short glass tube, which is connected 
 by means of a rubber tube to a generating-bottle similar to 
 that used with Squibb' s Urea Apparatus. Into the control- 
 tube is poured the solution of sodium sulphate, sufficient to 
 fill the burette to the zero mark and have the surface of the 
 liquid in both tubes on a level. 
 
 About 30 cc. of strong permanganate solution acidulated 
 with sulphuric acid are now placed in the generating-bottle, 
 and then the small test-tube or homeopathic vial, containing 
 exactly i cc. of hydrogen dioxid, is placed in. The generating- 
 bottle is then stoppered and agitated, the evolved gas passes 
 over, and forces down the liquid in the burette; the control- 
 tube is then lowered so as to bring the surfaces of the liquid 
 in both tubes on a level. 
 
 The reading is then taken. 
 
 Each cc. of gas represents one-half volume of oxygen 
 evolved from the dioxid if i cc. of the latter is used. Each cc. 
 
HYDROGEN DIOXID 319 
 
 of oxygen evolved from i cc. of the dioxid represents also 0.0017 
 gm. of absolute H2O2, or 0.0008 gm. of available oxygen. 
 
 Thus if from i cc. of the solution of hydrogen dioxid 20 
 cc. of gas are evolved, it is a so-called ten-volume solution 
 and contains 0.0017X20=0.034 gm. of absolute H2O2, or 
 0.0008X20=0.016 gm. of available oxygen. 
 
 According to Naylor and Dyer (Trans. Brit. Ph. Conf., 
 1901, 339) the gasometric permanganate method is unreliable, 
 because under the conditions of the test sulphuric acid added 
 to the brine solution in the nitrometer naturally liberates a 
 Uttle hydrochloric acid, and this in the presence of perman- 
 ganate becomes to some extent decomposed into chlorin. It 
 is the uncertainty as to the extent to which the chlorin is 
 absorbed by the water, which renders the accuracy of the 
 method doubtful. The results of this method are uniformly 
 too high, whether the gas be collected over mercury, over 
 saturated magnesium sulphate, or over brine, and in the 
 latter case considerably higher. But when the dichromate 
 V.S. is used (without acid), closely concordant results are 
 obtained, whether the gas be collected over mercury, or the 
 other solutions. The evolution of oxygen by the latter 
 method is slower than when permanganate is used, but the 
 oxygen obtained represents the volume available in the sample. 
 
 In the Hypochlorite Method, the nitrometer is filled with 
 a saturated solution of sodium chlorid. Two cc. of the hydro- 
 gen dioxid are admitted into the measuring tube, the funnel 
 tube filled with a little water, and this also let in, then 20 cc. of 
 the chlorinated lime solution introduced. From this point the 
 procedure is the same as in the gasometric permanganate method. 
 
 Ca(C10)2 + 2H202 = 04-^CaCl2 + 2H20. 
 
 The presence of preservatives, except inorganic ones, gives 
 low results. 
 
320 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 The Hypobromite Method. W. M. Dehn (J. A. C. S., 
 XXIX (9), 13 1 5) describes an accurate and rapid determination 
 of iiydrogen dioxid by means of sodium hypobromite, using 
 a ureometer. The reaction involved is 
 
 H202 + NaBrO = NaBr + H20 + 02. 
 
 The apparatus is shown in Fig. 58. The following descrip- 
 tion of the method is by Dehn from the 
 Journal of the American Chemical Society. 
 " The stop-cock E is opened and the stop- 
 cocks D and F are closed; then the solution of 
 sodium hypobromite * is poured in at the top 
 of C until it fills the tubes A and C to some 
 point above the stop-cock E. 
 
 " The stop-cock E is then closed and the 
 stop-cock F is opened, so that the hypobro- 
 mite in C may rim down to the constricted 
 portion; the solution in A is then sustained 
 by atmospheric pressure. The stop-cock D 
 (arranged so as to deliver only in the two 
 directions of a right-angle triangle) is turned 
 from the position shown in the figure and is 
 so controlled that B may first be washed with 
 a little of the hydrogen peroxid and then be 
 filled with the same to a readable height on 
 the scale. Upon turning D so as to admit a 
 regulated volume of the hydrogen peroxid solution an immediate 
 evaluation of oxygen results. After admitting most of the hypo- 
 bromite held above E, and letting stand for a minute or two 
 so as to drain properly, the columns of hypobromite in A 
 
 Fig. 58. 
 
 * This solution is prepared as directed under Estimation of Urea, except 
 that it is finally diluted with an equal volume of water. 
 
HYDROGEN DIOXID 
 
 321 
 
 and C are brought to the same level, the volume of oxygen 
 is read off and its weight and that of the corresponding 
 hydrogen dioxid are calculated by the usual formulas." 
 
 The author of this method also gives the following table, 
 by means of which the cc. of oxygen, under various conditions 
 of temperature and pressure, may be calculated into milli- 
 grams of hydrogen dioxid, and claims that by the use of this 
 instrument, this hypobromite method and the table for calcu- 
 lating the assay of hydrogen dioxid, is not only rapid and 
 accurate, but the necessity of preparing and correcting standard 
 solutions is avoided and the presence of the usual preserva- 
 tives used in the dioxid solution may be ignored. 
 
 WEIGHT 
 
 IN MILLIGRAMS 
 
 OF H2O, 
 
 CORRESPONDING 
 
 TO ONE 
 
 
 CUBIC 
 
 CENTIMETER OF 
 
 MOIST 
 
 OXYGEN 
 
 
 t, mm. 
 
 728 
 
 732 
 
 736 
 
 740 
 
 744 
 
 748 
 
 4° 
 
 I . 2664 
 
 1.2734 
 
 1.2802 
 
 1.2872 
 
 1.2942 
 
 I.3011 
 
 8 
 
 I . 2463 
 
 I-2531 
 
 I . 2600 
 
 I . 2669 
 
 1.2736 
 
 1 . 2805 
 
 12 
 
 I. 2251 
 
 1-2317 
 
 I.23S7 
 
 1.2454 
 
 1.2522 
 
 1.2589 
 
 i6 
 
 1.2044 
 
 1.2111 
 
 I. 2178 
 
 1.2245 
 
 1.2311 
 
 1-2378 
 
 20 
 
 1.1S17 
 
 I. 1884 
 
 I . 1948 
 
 I. 2015 
 
 I . 2080 
 
 I. 2145 
 
 24 
 
 1.T583 
 
 I . 1649 
 
 I.1719 
 
 I. 1777 
 
 I . 1843 
 
 I. 1907 
 
 28 
 
 I - 1345 
 
 1.1411 
 
 I. 1476 
 
 I -1538 
 
 I . 1603 
 
 I. 1665 
 
 32 
 
 I . 1085 
 
 1.1149 
 
 1. 1 2 13 
 
 I. 1275 
 
 1-1338 
 
 I.I40I 
 
 36 
 
 1.0843 
 
 1.0905 
 
 1.0967 
 
 I. 1050 
 
 I . 1093 
 
 I- "55 
 
 40 
 
 1.0605 
 
 1.0666 
 
 1.0725 
 
 1.0786 
 
 1.0849 
 
 1.0909 
 
 ■ //mm. 
 
 752 
 
 756 
 
 760 
 
 764 
 
 768 
 
 4° 
 
 1. 3081 
 
 1-3151 
 
 1.3222 
 
 1.3290 
 
 r-3359 
 
 8 
 
 1.2876 
 
 1-2944 
 
 1. 3014 
 
 I. 308 I 
 
 
 315° 
 
 12 
 
 1.2657 
 
 1.2726 
 
 1.2823 
 
 I . 2860 
 
 
 2928 
 
 16 
 
 I . 2444 
 
 I. 2512 
 
 1.2578 
 
 I . 2946 
 
 
 2713 
 
 20 
 
 I. 2213 
 
 1.2279 
 
 1.2345 
 
 I. 24 10 
 
 
 2475 
 
 24 
 
 I. 1972 
 
 I . 2036 
 
 I. 2 100 
 
 I. 2169 
 
 
 2230 
 
 28 
 
 1. 173 1 
 
 I. 1796 
 
 I. 1857 
 
 I. 1922 
 
 
 1986 
 
 32 
 
 I. I 465 
 
 1-1528 
 
 I. 1589 
 
 1. 1562 
 
 
 171S 
 
 36 
 
 1.1214 
 
 I. 1279 
 
 I-I34I 
 
 1. 1402 
 
 
 1465 
 
 40 
 
 1.0971 
 
 I -1033 
 
 I . 1094 
 
 I-II55 
 
 
 1216 
 
CHAPTER XXI 
 
 ESTIMATION OF SOLUBLE CARBONATES BY THE USE OF 
 THE NITROMETER 
 
 The nitrometer may be used for estimating ammonium 
 carbonate in aromatic spirit of ammonia. 
 
 The nitrometer in this case must be charged with mercuiy, 
 as the liberated CO2 is soluble in aqueous liquids. 
 
 A given volunle of the spirit is introduced into the nitrometer 
 followed by an excess of dilute HCl, and the evolved gas then 
 read off; and from its quantity the proportion of ammonium 
 carbonate may be calculated by applying the equation 
 
 (NH4) 2CO3 + 2HCI = 2 NH4CI + H2O + CO2. 
 
 *g6 *44 
 
 The volume of gas liberated must first be reduced to its 
 corresponding volume at 0° C. 
 
 Each cc. of CO2 at 0° C. weighs 0.001966 gm. Now if 
 44 gms. of CO2 represent 96 gms. of normal ammonium car- 
 bonate, how much ammonium carbonate does 0.001966 gm. 
 of CO2 represent? 
 
 44 : 96 : : o.ooi 966 :x. x--= 0.004289 gm. 
 
 Thus each cc. of CO2 at normal pressure and 0° C. 
 represents 0.004289 gm. of (NH4)2C03, approximately. 
 
 * The atomic weights are approximate. 
 
 322 
 
CHAPTER XXII 
 
 ESTIMATION OF UREA IN URINE 
 
 The determination of urea is based upon the fact that 
 when urea is decomposed by an alkaline hypochlorite or 
 hypobromite, carbon dioxid and nitrogen are given off, as 
 the equation shows: 
 
 CO(NH2)2+3NaBrO = 3NaBr + C02 + N2 + 2H20. 
 
 The liberated N may be measured, and from its quantity 
 the amount of urea calculated. The other products of the 
 decomposition go into solution. 
 
 The hypobromite solution is prepared as follows: loo 
 gms. Na,OH are dissolved in 250 cc. of water, and when this 
 solution has become cold 25 cc. of bromin are added, and 
 the solution kept cold. This solution contains sodium hypo- 
 bromite, bromate and hydroxid. It readily undergoes decom- 
 position, and should therefore always be freshly prepared 
 when wanted for use. To 15 cc. of the NaOH solution add 
 I cc. of bromin. 
 
 The solution of sodium hypochlorite is generally preferred 
 to the hypobromite, because it is more stable, just as effica- 
 cious, and the disagreeable handling of bromin is obviated. 
 
 Various forms of apparatus have been devised for the 
 quantitative estimation of urea. 
 
 The Doremus Ureometer (Fig. 59) is the simplest of these. 
 The long arm of the ureometer is filled with the hypobromite 
 solution, and then i cc. of the urine is introduced by the aid 
 of the pipette. The pipette is introduced through the bulb 
 
 323 
 
324 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 as far as it will go in the bend, and the nipple is then gently 
 but steadily compressed, being careful that no air is admitted. 
 
 The volume of the liberated gas is read off after the froth 
 has subsided. 
 
 The ureometer indicates, according to its graduation, either 
 milligrams of urea in i cc. or grains of urea per fluid ounce 
 of urine. 
 
 It also indicates by the signs +, N, and — whether the 
 
 ^^ 
 
 Fig. 59. 
 
 Fig. 60. 
 
 urea is present in an increased, normal or decreased quantity. 
 Either Knop's or Squibb's solution, or Liquor Sodee Chloratse 
 U. S. P. may be used in this instrument. Knop's solution is 
 that described above. Squibb's solution contains potassium 
 bromid as well as bromin. It is prepared by taking an equal 
 weight of bromin and of potassium bromid, and adding eight 
 times as many cc. of water as there were grams of bromin 
 taken. For use mix equal volumes of this solution with the 
 sodium hydroxid solution above described. 
 
ESTIMATION OF UREA IN URINE 
 
 32.J 
 
 The Hinds-Doremus Ureometer. This apparatus, which 
 is shown in Fig. 60, is capable of giving more exact results 
 than the original apparatus, because the i cc. of urine can 
 be delivered more accurately. It consists of a bulb with an 
 upright tube a, graduated like the original, so that each of 
 the smallest divisions represents o.ooi gm. of urea in the urine 
 used. The lower portion of this tube is, in connection with 
 a smaller tube c, graduated with a capacity of 2 cc; between 
 these tubes a glass stop-cock is situated. Closing the stop- 
 cock h, Knop's or Squibb's fluid (diluted one half) or liquor 
 sodae chloratas U. S. P. is introduced into tube a so as to 
 
 Fig. 61. 
 
 completely fill it. The apparatus is then placed in an upright 
 position and the smaller tube c is filled to the zero mark with 
 urine. The stop-cock is then turned slowly so as to admit 
 gradually i cc. of the urine to tube a. After fifteen minutes 
 the reading is taken. If the reading be 0.015 ^.nd the amount 
 of urine taken was i cc, then multiplying by 100 gives 1.5 
 per cent of urea. 
 
 Squibb's Urea Apparatus (Fig. 61) is a very simple appa- 
 ratus, and can be easily improvised in a drug store. It con- 
 sists of two wide-mouthed bottles, the larger of which, C, 
 capable of holding about 250 cc, is fitted with a rubber stopper, 
 through which is passed a curved delivery-tube and a short 
 straight tube, the latter connected by a piece of rubber tubing 
 
326 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 to the short glass tube in the rubber stopper of the smaller 
 or generating-bottle B. In the generating-bottle is a small 
 test-tube A. 
 
 Into the test-tube A is placed 5 cc. of urine, and into the 
 smaller bottle B is put 20 cc. of the h)^obromite solution, 
 or strong liquor sodae chloratae. The test-tube is then placed 
 in the generating-bottle B, being careful that the urine and 
 the reagent do not come in contact. The larger bottle C is 
 now filled with water and the two bottles connected by the 
 rubber tube, the larger bottle being placed on its side upon 
 a block, and when all connections are tight, the generating- 
 bottle is shaken so that the urine will mix with the reagent. 
 
 Decomposition takes place, and the generated gas passes 
 into the bottle C, displacing water, which is caught in a grad- 
 uated cylinder or other measuring vessel. The volume of 
 water displaced is equivalent to the volume of gas evolved. 
 
 Each cc. of nitrogen gas evolved at 0° C. and normal 
 pressure represents 0.00268 gm. of urea. Then by multi- 
 plying the number of cc. evolved by this number the quantity 
 of urea in the 5 cc. of urine taken is ascertained. 
 
 The volume of gas obtained when the operation is con- 
 ducted at ordinary temperatures should always be reduced 
 to its corresponding volume at 0° C. and 760 mm. 
 
 The factor 0.00268 is found in the following manner: 
 
 1000 cc. of H at 0° C. weigh 0.0896 gm.; 
 1000 cc. of N at 0° C. weigh 1.255 gms. 
 
 By the equation it is seen that 60.02 gms. of urea evolve 
 when decomposed 28.02 gms. of N. 
 
 CO(NH2)2-h3NaBrO = 3NaBr-hC02-hN2-l-2Pl20. 
 60.02 28.02 
 
ESTIMATION OF UREA IN URINE 327 
 
 Now we v/ill find the volume occupied by 28.02 gms. of 
 N at 0° C. 
 
 1.255 g"^s. of N=iooo cc. 
 
 gms. cc. gms. cc. 
 
 1.255:1000: : 28.02:^. x = 22326. 
 
 Thus 60.02 gms. of urea evolve 22326 cc. of N. One cc. 
 of N thus represents 0.00268 gm. of urea. 
 
APPENDIX 
 
 DESCRIPTION OF INDICATORS 
 
 The following list includes the more reliable indicators in 
 common use, arranged alphabetically. 
 
 AliT-arin- alkalies = ;?ed 
 
 This dye was first found in the roots of madder {Rubia 
 tinctorium) but is now also obtained synthetically. A half 
 per cent solution in alcohol is employed as an indicator. 
 
 AzoHtmin: ^^Z^ 
 
 This is the color principle to which litmus owes its value 
 as an indicator. Its extraction is explained under litmus. 
 It is a high-priced article and is in consequence seldom used, 
 the purified litmus tincture being preferred. 
 
 Brazil Wood Solution: addf'=y"««™''"''''^ 
 
 Boil 50 gms. of finely-cut Brazil-wood (the heart-wood 
 of Peltophorum dubium (Sprengel) Briton, nat. ord. LeguminoscB) 
 with 250 cc. of water during half an hour, replacing from 
 time to time. Allow the mixture to cool; strain; wash the 
 contents of the strainer with water until 100 cc. of strained 
 liquid are obtained; add 25 cc. of alcohol and filter. 
 
 This indicator is especially serviceable for the titration of 
 
 328 
 
DESCRIPTION OF INDICATORS 329 
 
 alkaloids, but it is useless in the presence of sulphurous acid, 
 sulphites, or sulphids, as these substances decolorize the solu- 
 tion. 
 
 Cochineal: alkalies =FJ«/«i 
 
 acids = V ellowtsh-rcd 
 
 Cochineal is the dried female insect, Pseudococciis cacti 
 Linne. 
 
 The test solution is made by macerating 3 gms. of unbroken 
 cochineal for four days in 250 cc. of a mixture of i part of 
 alcohol and 3 parts of water, .by volume. It should be neu- 
 tralized with ammonia water before using. 
 
 It is a very valuable indicator, especially for carbonates 
 of the alkalies and alkaline earths, because it is not affected 
 by CO2. It is also useful for titrating alkaloids, alkalies, 
 alkali earths, ammonia, and inorganic acids, but is useless 
 for most organic acids. 
 
 Congo Red: t^^Zlte 
 
 Congo red is a sodium tetrazo-diphenyl-naphtionate. It 
 occurs in commerce in the form of lumps of a reddish-brown 
 color. It is readily soluble in both water and alcohol. Its 
 solution is exceedingly sensitive to free acids even in the presence 
 of acid salts, and is likewise very sensitive to free carbonic 
 and acetic acids. It may be employed for estimating free 
 mineral acids, in the presence of most organic acids. 
 
 The test solution contains i per cent of the dye and 10 
 per cent of alcohol. 
 
 CiaWfin • a^^a-yies = Bright red 
 uauein. ^^jjg = Pale brown 
 
 Anthracene violet or pyrogallol-phthalein was proposed by 
 M. Dechan for use as an indicator. 
 
 It is prepared by heating a mixture of one part of phthalic 
 
330 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 anhydrid and two parts of pyrogallol, and finally recrys- 
 tallizing in a similar way to phenolphthalein. 
 
 It is described as a dark reddish crystalline solid, possessing 
 a greenish luster. It is nearly insoluble in water, but readily 
 soluble in alcohol. In commerce it is frequently found as a 
 paste, mixed with water. 
 
 It forms a red coloration with alkalies, which is changed 
 to yellowish-brown on addition to an acid in excess. 
 
 It is said to be more delicate toward alkalies than phenol- 
 phthalein, and may be used in its stead for titrating many 
 of the alkaloids. It may be used in the presence of ammonia 
 or its salts. It indicates sharply with the organic acids. A 
 solution in. rectified spirit i-iooo is generally employed. 
 
 Hematoxylin: t^r^UlL to oran.e 
 
 A peculiar principle obtained from logwood, having the 
 composition Ci6Hi406, and crystallizing with one or three 
 molecules of water. It is an effiorescent yellowish-rose colored 
 substance, but when pure is said to be colorless, reddening 
 on exposure to light. 
 
 It is soluble in hot water or alcohol. Its alcoholic solution 
 is largely used as an indicator in the titration of alkaloids, 
 for which it is considered the indicator par excellence. 
 
 The solution is prepared by dissolving 0.3 gm. of the well- 
 crystallized material in 100 cc. of alcohol. In titrating use 
 about three drops of this solution. 
 
 Its color reaction with carbonates and bicarbonates is 
 interesting. When added to a solution of alkali-bicarbonate 
 the reaction requires many seconds, and results in a gradually 
 deepening carm.ine-red which is permanent, while in the case 
 of soluble carbonates the reaction is instantaneous, a purple- 
 red which changes rapidly through cherry, eosin-red to orange. 
 
DESCRIPTION OF INDICATORS 331 
 
 The reaction with ammonium carbonate is similar to that 
 with bicarbonate. 
 
 lodeosin: ^'^t^lfX^'^ 
 
 Tetra-iodo-fluorescein Erythrosin B. This indicator is use- 
 ful for minute quantities of alkali, as for instance, such as 
 may be dissolved out from glass on contact with water. It 
 is used in connection with highly dilute standard solutions 
 only. The iodeosin solution is made by dissolving 0.002 gm. 
 of the indicator in 1000 cc. of pure ether. Titration with this 
 indicator is carried out by introducing 50 to 100 cc. of the 
 liquid to be titrated into a stoppered bottle and adding 10 to 
 20 cc. of the ethereal indicator solution and setting aside 
 after shaking. The ethereal layer as well as the fluid will 
 be colorless if the latter is neutral, but if traces of alkali are 
 present the rose-red tint passes into the aqueous liquid leaving 
 the ether colorless. If the fluid is acid, the ethereal layer is 
 yellow. If preferred, 4 or 5 drops of a 1-10,000 alcoholic 
 solution of the indicator may be added to the liquid, and 
 ether then added. lodeosin is particularly useful in titration 
 of alkaloids, especially those of weak basicity, as emetine, 
 atropine, morphine, etc. 
 
 lacmoid- alkalies = B/«e 
 Lacmoia. ^^^^ _^^^ 
 
 Lacmoid is somewhat allied to litmus, but differs from 
 it in many respects. It is a product of resorcin, and may 
 be prepared by heating gradually to 110° C. a mixture of 
 100 parts of resorcin, 5 parts of sodium nitrite, and 5 parts 
 of water. After the violent reaction moderates it is heated 
 to 120° C. until ammonia ceases to be evolved. The residue 
 is then dissolved in warm water and the lacmoid precipitated 
 therefrom by HCl; the free acid is then removed by washing 
 and the residue dried. 
 
332 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 This constitutes the commercial lacmoid, which is not in 
 a sufficiently pure state to be used as an indicator; its puri- 
 fication is effected according to Forster by treating the powder 
 with boiling water and acidulating the resulting blue solution 
 with hydrochloric acid. After a few hours the precipitate is 
 collected and washed with a little cold water and carefully 
 dried, or it is dissolved in alcohol and the solution evaporated. 
 Even after careful purification, lacmoid solution may still 
 exhibit a violet tinge, which is a disturbing factor in accurate 
 work. To remedy this defect, Forster suggests the addition 
 of 5 gms. of beta-naphthol green to 3 gms. of purified lacmoid 
 dissolved in 700 cc. of water and 300 cc. of alcohol. A 
 sample of lacmoid which is only sparingly soluble in water 
 should be rejected. The purer the article the more readily 
 does it dissolve in water. A good T.S. may be made by 
 dissolving one part in 12 parts of 20 per cent alcohol, filtering 
 and evaporating in vacuo over sulphuric acid. Of the residue 
 obtained in this way 0.2 gm. is dissolved in 100 cc. of alcohol. 
 
 Lacmoid paper is prepared by dipping slips of calendered 
 unsized paper into the blue or red solution and drying them. 
 
 Lacmoid is slightly affected by carbonic-acid gas. It may 
 be used cold for the alkaline and earthy hydroxids, arsenites, 
 and borates and the mineral acids. The carbonates and 
 bicarbonates of the alkalies and alkali earths are titrated- 
 hot with this indicator. 
 
 Many of the metallic salts, such as the sulphates and 
 chlorids of iron, copper and zinc, which are more or less 
 acid to litmus, are neutral to lacmoid; therefore free acids 
 in such solutions may be estimated by its aid. 
 
 Lacmoid paper reacts alkaline with the chromates of potas- 
 sium or sodium, but neutral with the dichromates, so that 
 a mixture of the two or of chromic acid and dichromate may 
 be titrated by its aid. 
 
DESCRIPTION OF INDICATORS 333 
 
 I :^m..K f T n^^..,^\ alkalies = Btee 
 
 Litmus {Lacmus): ^^j^^ ^^^^ 
 
 A pigment obtained by the fermentation of certain lichens, 
 principally from Roccella tincloria and R. fuciformis, but also 
 from other species of lichen. 
 
 It occurs in commerce in small, friable, light cakes or 
 cubes, of a violet color. 
 
 The coloring principles of litmus are azolitmin, erythro- 
 litmin, and erythrolein. The first, which is the most important, 
 is soluble in water, but insoluble in alcohol. The other two 
 are readily soluble in alcohol, but only sparingly soluble in 
 water. 
 
 The process for making litmus test solution consists in 
 exhausting coarsely powdered litmus with boiling alcohol. 
 
 The residue is then digested with about an equal weight 
 of cold water so as to dissolve the excess of alkali present. 
 
 The blue solution thus obtained, after being acidulated, 
 may be used to make red litmnx-paper. Finally, the residue 
 is extracted with about five times its weight of boiling water 
 and the solution filtered. 
 
 The filtrate is preserved as test solution in wide-mouthed 
 bottles, stoppered with loose plugs of cotton to exclude dust, 
 but to admit air. 
 
 When kept in closed vessels litmus solution gradually 
 loses color, but this returns upon exposure to air and conse- 
 quent absorption of oxygen. 
 
 The fermentation to which the loss of color is due niay 
 be prevented by saturating the solution with NaCl or by the 
 addition of thymol or phenol. 
 
 The British Pharmacopoeia recommends the boiling of litmus 
 in powder with three succcessive portions of rectified spirit, 
 and then to digest the residue in distilled water and filter, 
 the object of these steps in the process being to get rid of 
 
334 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 the greater portion of erythrolitmin and erythrolein, which 
 
 are soluble in alcohol. Then by treating the residue with 
 
 water a larger proportion of azolitmin is dissolved, and the 
 
 solution is contaminated with very little of the other two 
 
 principles. 
 
 Litmus test solution should be of such strength that 3 
 
 drops added to 50 cc. of water will impart to the latter a 
 
 N 
 distinct color. If one drop of — acid or alkali solution be 
 
 ^ 10 
 
 added to this the color should change to red or blue. 
 
 Litmus may be used in a very large number of titrations. 
 It is of value in the titration of most mineral acids and of a 
 few organic acids, e.g., benzoic and oxalic. It is also useful 
 in the titration of alkaline hydroxids when the latter are free 
 from carbonates. 
 
 But for carbonates, bicarbonates, etc., a reliable end-reaction 
 can only be obtained by boiling the solution during the titra- 
 tion, in order to dispel the liberated CO2. 
 
 Free CO2 has an acid reaction with litmus, and interferes 
 very much with the finding, of the end-reaction. 
 
 Litmus may be used for ammonia and for borax. It is 
 of no use for phosphoric or arsenic acid, nor for sulphurous 
 acid, phosphates, or arsenates, because the change of tint 
 is too gradual. 
 
 It is unsatisfactory in titrating many organic acids, e.g., 
 tartaric and citric, but may be used for oxalic or benzoic, 
 as before stated. 
 
 Sometimes it is required to perform a titration with litmus 
 at night. Gas- or lamp-light is not adapted for showing the 
 reaction satisfactorily, but by using a monochromatic light, 
 such as the sodium flame, a very sharp line of demarcation 
 may be found. 
 
 The operation should be conchacted in a dark room, using 
 
DESCRIPTION OF INDICATORS 33.J 
 
 a piece of platinum foil sprinkled with salt or a piece of pumice- 
 stone saturated with a solution of salt, heated in a Bunsen 
 flame. 
 
 The red color then appears perfectly colorless, while the 
 blue appears like a mixture of ink and water. 
 
 I iif Ar>l • alkalies = FcBow 
 mieoi. ^^jjg ^Colorless 
 
 Chemically it is an oxy-chlor-diphenyl-quinoxalin. It was 
 suggested as an indicator by Autenrieth. 
 
 The solution for the purpose of an indicator is prepared 
 by dissolving i part in loo parts of alcohol. Of this, four 
 drops are sufficient for 50 cc. of fluid to be titrated. 
 
 In sensitiveness, luteol exceeds both litmus and phenol- 
 phthalein. It is more sensitive toward ammonia than Nessler's 
 solution. Ten cc. of a solution containing one drop of am- 
 monia water per liter, is colored yellow immediately upon add- 
 ing luteol, whereas with Nessler's solution it takes quite some 
 time before a reaction is obtained. 
 
 Methyl Orange: ^'^f^ ^I^'f" " 
 
 Poirrier's Orange III, Tropaeolin D, Helianthin, Mandarin- 
 orange, Para-sulpho-benzeneazo-dimethylanilin. 
 
 This is prepared by the action of diazo-sulphanilic acid 
 upon dimethylanilin; the acid so formed is converted into 
 a sodium or ammonium salt, purified by reprecipitation with 
 HCl, and again converted into a sodium or ammonium salt. 
 If prepared carefully and from the purest materials, it is a 
 bright orange-red powder, perfectly soluble in water and 
 slightly in alcohol; but it is often found in commerce as a 
 dull orange-brown powder, often not completely soluble in 
 water. Many conflicting statements have been made by opera- 
 
336 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 tors as to the value of methyl orange as an indicator, which 
 have tended to bring this indicator into disrepute. 
 
 Sutton has examined many specimens, but has not found 
 any in which the impurities sensibly affected its delicate action. 
 He claims that the common error is the use of too much 
 indicator, and that some eyes are more sensitive to a change 
 of tint than others. 
 
 Methyl orange is no doubt a very good indicator, but 
 practice with it must be had in order to obtain good results. 
 The author has found one sample which had a beautiful 
 orange color, but which was absolutely useless as an indicator. 
 
 A. H. Allen describes the characters and tests of a good 
 article as follows: 
 
 1. Aqueous solution, not precipitated by alkalies. (Orange 
 I becomes red-brown: orange II brownfsh-red.) 
 
 2. Hot concentrated aqueous solution yields with HCl 
 microscopic acicular crystals of the free sulphonic acid, soon 
 changing to small lustrous plates or prisms having a violet 
 reflection. (Orange I gives yellow-brown color or flocculent 
 precipitate; orange II brown-yellow precipitate.) 
 
 3. Dissolves in concentrated H2SO4 with a reddish or 
 yellowish-brown color, which on dilution becomes fine red. 
 
 4. BaCly yields a precipitate. 
 
 5. CaCl2 yields no precipitate. (Orange I gives a red 
 precipitate.) 
 
 6. Pb(C2H302)2 yields an orange-yellow precipitate. 
 
 7. MgS04 in dilute solutions precipitates the coloring 
 matter in microscopic crystals. 
 
 Methyl-orange T.S. is made by dissolving i gm. of methyl 
 orange in 1000 cc. of water. Add to it carefully diluted sul- 
 phuric acid in drops until the liquid turns red and just ceases 
 to be transparent. Then filter. 
 
 The great value of this indicator consists in the fact that 
 
DESCRIPTION OF INDICATORS 337 
 
 it is not affected by carbonic-acid gas, sulphurated hydrogen, 
 or silicic, oleic, stearic and many other acids. 
 
 It answers well for ammonia, but it is useless for most 
 of the organic acids. Phosphoric and arsenic acids are rendered 
 neutral to methyl orange when only one third of the acid has 
 combined with the base, the end-reaction being well defined. 
 (Phenolphthalein indicates neutrality when two-thirds of acid 
 are combined.) 
 
 This indicator should not be employed when titrating with 
 standard solutions which are weaker than decinormal, nor 
 should it be used in any hot titrations, nor in excessive quan- 
 tities. Two drops are sufficient for 50 cc. of the fluid to be 
 titrated, or just enough to give it a faint tint. 
 
 ,, ,. J carbonates = i?ed 
 Phenacetolin : { hydroxids = Yellow 
 
 acids = Yellow 
 
 This indicator is prepared by heating together for several 
 hours equal molecular weights of phenol, glacial acetic acid, 
 and sulphuric acid in a vessel provided with a reflux con- 
 denser. The product is then thoroughly washed with water 
 to remove excess of acid and dried for use. It is only very 
 slightly soluble in water, but dissolves readily in alcohol, 
 forming a greenish-brown solution. 
 
 The solution yields with alkali hydroxids a scarcely per- 
 ceptible pale yellow, but with normal carbonates of the alkalies, 
 sulphids, and with ammonia it gives a decided pink color; 
 with bicarbonates a more intense pink, while with acids a 
 golden yellow. 
 
 This indicator is useful for estimating the amount of alkali 
 or alkali earth hydroxids in the presence of carbonate, unless 
 the hydroxid is present in too small a quantity. Ammonia 
 must not be present. The titration is carried out by adding 
 the acid until a faint red color appears; this indicates that 
 
338 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 
 the alkali hydroxid or the lime has been neutralized. The 
 further addition of the acid intensifies the red until the car- 
 bonate present in the mixture is neutralized, when a golden- 
 yellow color appears. The proportion of alkali hydroxid must 
 be far in excess of the carbonate in order to obtain reliable 
 results; furthermore, considerable practice is required in the 
 use of this indicator in order to accustom the eye to the color 
 changes. 
 
 A convenient strength of solution is i : loo in alcohol. 
 
 Phenolphthalein (C.oH:404): ^'^t^Icl/... 
 
 Preparation. Five parts of phthalic anhydrid (C8H4O3), 
 10 parts of phenol (CeHeOH), and 4 parts of H2SO4 are 
 heated -together at 120° to 130° C. for several hours. The 
 product is then boiled with water, and the residue, which 
 consists of impure phenolphthalein, is dissolved in dilute 
 soda solution and filtered. By neutralizing this solution the 
 phenolphthalein is precipitated and may be purified by crys- 
 tallization from alcohol; or the alcoholic solution may be 
 boiled with animal charcoal, filtered, and the phenolphthalein 
 reprecipitated by boiling water. 
 
 Uses. Phenolphthalein is a very valuable indicator; is 
 extremely sensitive, and exhibits a well-marked and prompt 
 change from colorless to pink, and vice versa. A few drops 
 of the solution of the indicator show no color in neutral or 
 acid liquids, but the faintest excess of alkali produces a sudden 
 change to red. 
 
 It may be employed in the titration of mineral and organic 
 acids and most alkalies, but it is not suited for the titration 
 of ammonia or its salts. It is very sensitive to CO2, and 
 therefore in estimating carbonates the liquid must be boiled. 
 
DESCRIPTION OF INDICATORS 339 
 
 as with litmus. It is inapplicable for borax, ejJcept in the 
 presence of glycerin, because the color gradually fades away 
 as the acid is added. One great advantage which phenol- 
 phthalein possesses is that its indications may be clearly 
 read in many colored liquids; another is that it may be used 
 in alcoholic liquids or in mixtures of alcohol and ether, and 
 therefore many organic acids which are insoluble in water 
 may be accurately titrated by its help. 
 
 Phenolphthalein T.S. is a one per cent solution in alcohol. 
 
 1, ,. f carbonates =B/Mfi 
 
 Poirrier Blue (C4B): ''"^'''' j hydroxids ^Red 
 acids =Blue 
 
 This indicator, which is closely allied to Gentian Blue in 
 properties, is obtained by the action of sulphuric acid on 
 triphenylrosanilin. It is a blue powder with a coppery luster. 
 It dissolves in water and in alcohol, yielding blue solutions. 
 KOH and NaOH change the color to red, but ammonia 
 decolorizes at. It is employed as an indicator in aqueous 
 solution 1:500. This indicator is exceedingly sensitive to 
 acids. Borax and boric acid give a blue color; in the titra- 
 tion of boric acid the red color does not appear until the acid 
 is completely neutralized. This indicator is recommended 
 for the titration of hydrocyanic acid, toward which it is especially 
 sensitive, the alkaline cyanids are alkaline in reaction to most 
 indicators, but C4B does not show an alkaline reaction until 
 the HCN is completely neutralized, and a minute excess of 
 the alkali hydroxid has been added. C4B is of the character 
 of a weak acid and its salts are very unstable; they are decom- 
 posed by water alone when in very great dilution, therefore 
 the indicator must be used in sufficient quantity. The addition 
 of a few drops of alcohol facilitates the color change, which 
 is indeed a very sharp one. 
 
340 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 
 Resazurin: alkalies =Bto« 
 
 acids =Rea 
 
 This is a new indicator for alkalimetry, proposed by Crismer. 
 It is prepared as follows: Dissolve 4 gms. of resorcin in 300 
 cc. of anhydrous ether and add 40 to 45 drops of nitric acid 
 (sp.gr. 1.25) saturated with nitrous anhydrid. Allow the 
 mixture to stand in a cold place for two days, whereupon 
 a deposit of blackish crystals, having a reddish-brown reflec- 
 tion, will be formed in the bottom of the vessel. The super- 
 natant clear red liquid is decanted and the crystals washed 
 with ether until the washings show a blue color with ammonia 
 water. 
 
 Resazurin (C12H7NO4) is slightly soluble in water, more 
 so in alcohol and freely soluble in acetic ether. It produces 
 a blue solution with water, alkalies, and alkali carbonates, 
 which are turned red upon the addition of a slight excess of 
 acid. To use this indicator in alkalimetry, Crismer recom- 
 mends the following solution: Resazurin 0.2 gm. dissolved 
 
 m 40 cc. of — ammonia solution, and made up to 1000 cc. 
 
 with distilled water. 
 
 This is deep blue in color and keeps well. Two or three 
 drops are sufficient to color 200 cc. of liquid. 
 
 This indicator is not suited for the titration of nitric acid 
 or monobasic organic acids, and it is not very sensitve to 
 carbonic acid. It is, however, extremely sensitive to alkalies. 
 If the solution is acidulated to a rose-red color and heated 
 in a white glass flask, the solution will turn blue through 
 the alkaline reaction of the dissolved glass before the boiling- 
 point is reached. 
 
 This indicator is especially useful for borax. 
 
DESCRIPTION OF INDICATORS 341 
 
 Rosolic Acid (C.0H14O3): ^"'ZZo. 
 
 ■This compound is also called Aurin and Coralline, and 
 is prepared as follows: 
 
 A mixture of phenol and sulphuric acid is placed upon 
 a water-bath, and oxalic gradually added, waiting each time 
 till the evolution of gas ceases, and using less oxalic acid than 
 is required to attack all the phenol. 
 
 In this process the oxalic acid is decomposed into CO, 
 CO2, and H2O. The CO immediately reacts with the phenol 
 and forms rosolic acid, as the following equation shows : 
 
 3C6H5OH + 2CO = C20H14O3 + 2H2O. 
 
 Commercial rosolic acid is a mixture of several derivatives 
 among them the above, methylaurin C20H16O3 and others. 
 Commercial poeonin (also known as Aurin R.) [chiefly C19H14O3] 
 may be used in place of rosolic acid. 
 
 Rosolic acid is soluble in diluted alcohol. Its color is 
 pale yellow, not changed by acid, but turns violet-red with 
 alkalies. 
 
 It is an excellent indicator for the mineral acids and strong 
 bases, weak ammoniacal solutions, oxalic acid and other or- 
 ganic acids, except acetic. 
 
 The test solution is made by dissolving i gm. of the com- 
 mercial rosolic acid in 10 cc. of diluted alcohol and then adding 
 enough water to make 100 cc. 
 
 T-«««-.^i:« fr\ni\ ■ alkalies =FeKoa' 
 Tropaeolm (OO). ^^-^^ ^Yellowish-red 
 
 This is used in the form of a solution containing 0.5 gm. 
 to 1000 cc. of alcohol. 
 
 Turmeric Tincture. Digest any convenient quantity of 
 
342 THE ESSENTIALS OT VOLUMETRIC ANALYSIS 
 
 ground curcuma-root (from Curcuma longa Linnd, nat. ord. 
 ScilaminecB) repeatedly with small quantities of water, and 
 throw this liquid away. Then digest the dried residue for 
 several days with six times its weight of alcohol and filter. 
 
 Turmeric Paper. Impregnate white, unsized paper with 
 the tinctiu-e and dry it. 
 
 The color principle of turmeric is curcumin. It is seldom 
 used in volumetric analysis, except in the form of turmeric 
 paper. For high-colored solutions curcumin gives no reaction 
 with acids, but becomes brown with alkalies. There is another 
 color principle in turmeric besides curcumin, which is, however, 
 useless in that it is indifferent to alkalies; it is soluble in water, 
 and extracted by digestion with water, after which the cur- 
 cumin is dissolved out with alcohol. 
 
 Turmeric paper is especially useful, because of its peculiar 
 reaction with boric acid, with which it develops a brown color 
 after drying, and which color, when touched with caustic 
 soda solution is changed to dark green. 
 
INDEX 
 
 PAGE 
 
 Acetate, lead 167 
 
 — potassium 80 
 
 — sodium 81 
 
 Acetic acid 100 
 
 Acid, acetic 100 
 
 — arsenous ig6 
 
 decinormal V.S 239 
 
 — • — solution of - 197 
 
 V.S., use of in reduction 238 
 
 — defined 56 
 
 — value of fats and oils 275 
 
 — boric 100 
 
 — chromic 1 73, 230 
 
 — citric loi 
 
 — hydriodic 100 
 
 — hydriodic by sulphocyanate method iig 
 
 — hydriodic, syrup of 1 20 
 
 — hydrobromic 100, 117 
 
 -by sulphocyanate method 117 
 
 by Volhard's method 117 
 
 using chromate as indicator 117 
 
 — hydrochloric ^6 
 
 normal 61 
 
 action of, on permanganate 142 
 
 standardization of '. . 64 
 
 standardization by sodium carbonate 63 
 
 — hydrocyanic 1 23 
 
 using chromate indicator ' 1 24 
 
 potassium iodid indicator 1 26 
 
 — hypophosphorous 99, 164 
 
 — lactic 103 
 
 — nitric 99. ii9> 3^5 
 
 3-13 
 
344 INDEX 
 
 ... . PAGE 
 
 Acid, nitrous 162 
 
 — oxalic jQ, 
 
 and oxalates !„ 
 
 decinormal .' gj 
 
 — phosphoric n^ 
 
 — rosolic ly 
 
 — sulphuric ng 
 
 normal 64 
 
 — sulphurous 100, 200, 202 
 
 — tartaric loi 
 
 Acidimetry gg 
 
 — and alkalimetry 57 
 
 Acids, estimation of, by neutralization gg 
 
 in salts 104 
 
 — haloid 116 
 
 — inorganic 5,4 
 
 — organic 100 
 
 — quantity to be taken for assay g(, 
 
 — weighing of, for assay g4 
 
 volatile 05 
 
 Alcohol, in tinctures and beverages 303 
 
 Alcoholometric table 304 
 
 Alizarin 32g 
 
 Alkali bicarbonates and carbonates mixed yc 
 
 — carbonates 68 
 
 — hydro xid and carbonate mixed ^3 
 
 — hydroxid, standard solution, preservation of no 
 
 — hydroxids, estimation g^ 
 
 — iodids 231 
 
 — metals, in their salts ' g3 
 
 — standard solutions, preparation of go 
 
 Alkali earth hydroxids g4^ gc 
 
 and carbonates mixed gy 
 
 salts 84, 86, 104 
 
 — earths, organic salts of yy 
 
 Alkalis combined with non-volatile acids g^ 
 
 ■ with volatile acids 83 
 
 ^ in presence of sulphites ^j 
 
 — organic salts of 76 
 
 Alkalimetry jg 
 
 — and acidimetry 57 
 
 Alkaloids, estimation of 249 
 
INDEX 345 
 
 PAGE 
 
 Alkaloids, separation of 258 
 
 Alum 104 
 
 Ammonium bromid 112 
 
 — carbonate 71 
 
 — chlorid as an impurity in the bromid 113 
 
 Ammonia water 67 
 
 stronger 68 
 
 Amyl nitrite 316 
 
 Anions 18 
 
 Anthracene violet 329 
 
 Antimonic compounds i gg 
 
 Antiraonous oxid 198 
 
 Antimony compounds ig8 
 
 — and potassium tartrate 200 
 
 Apparatus, cleaning of 39 
 
 — used in volumetric analysis 28 
 
 — use of 3g 
 
 Arsenic trioxid 196 
 
 Arsenite of potassium solution 197 
 
 Arsenous acid 196 
 
 solution 197 
 
 decinormal V.S 239 
 
 V.S., use of in reduction 238 
 
 — anhydrid 196 
 
 — compounds 196 
 
 direct percentage assay of 197 
 
 — iodid ig7 
 
 — oxid 196 
 
 standardization of iodine with 190 
 
 Atomic weights xii 
 
 multiples of xi 
 
 Azolitmin 3^8 
 
 Barium chlorid 104 
 
 — dioxid 15s. 158 
 
 — peroxid 158 
 
 — soluble salts of 86 
 
 Benzoate, sodium 81 
 
 Berzelius' system of oxids - 54 
 
 Bicarbonate of sodium 71 
 
 — of potassium 7° 
 
 Bismuth salts 104 
 
346 INDEX 
 
 PAGE 
 
 Bisulphite, sodium 20,- 
 
 Bitartrate, potassium yq 
 
 Bleaciiing powder 242 
 
 Boric acid loo 
 
 Boyle's Law ogg 
 
 Brazil wood ,28 
 
 test solution 17 
 
 Bromates 235 
 
 Bromate, potassium 236 
 
 Bromid, ammonium 112 
 
 Bromids m 
 
 Bromin free 218 
 
 — V.S ■ 269 
 
 — water 223 
 
 Burette, automatic 30 
 
 — connected with reservoir 31 
 
 — clamps 33 
 
 — glass stop-cock 29 
 
 — holder 33 
 
 — Mohr's 28 
 
 Burettes, special forms of 44 
 
 Burette supports 33 
 
 Butter, examination of 279 
 
 Calcium carbonate 85, 153 
 
 — chlorid 87 
 
 — hypophosphite 166 
 
 — salts 153, 166 
 
 ■ — soluble salts of 1 86 
 
 Calculating results 48 
 
 Calibration of instruments. 45 
 
 Calx chlorinata 220 
 
 Cane sugar, inverted 288 
 
 Carbonate, ammonium 71 
 
 — and hydroxid of alkali mixed 73 
 
 — calcium 85 
 
 -^ of lithium 71 
 
 — potassium 69 
 
 — of sodium (anhydrous) 71 
 
 (crystalllized) 70 
 
 — sodium, normal V.S 86 
 
 Carbonates and bicarbonates of alkalies, ^xed 1 175 
 
INDEX 347 
 
 PAGE 
 
 Carbonates and hydroxids of alkali earths, mixed 87' 
 
 — of alkalies 68 
 
 — soluble, assay of b\' tlie use of the nitrometer 322 
 
 Cathions 18 
 
 CiB 339 
 
 Centinormal solutions 10 
 
 Charles' Law 309 
 
 Chlorate potassium 236 
 
 Chlorates 234 
 
 Chlorid, barium 104 
 
 — calcium 87 
 
 — ferric 237 
 
 — of lime 220 
 
 — sodium 112 
 
 V.S 108 
 
 Chlorids iii 
 
 Chlorin, in bleaching powder 242 
 
 — in chlorin water 241 
 
 — free 218 
 
 — water 218 
 
 Chlorinated lime 220 
 
 ■ — soda solution 222 
 
 Chloiometry 238 
 
 Chromates 167, 210, 224, 230 
 
 Chromic acid i73. 230, 234 
 
 — anhydrid 230 
 
 — oxids 174 
 
 Chromium trioxid 230 
 
 Chrdmophoric theory 20 
 
 Citrate lithium 81 
 
 — potassium 80 
 
 Citric acid i°i 
 
 Cobalt salts i°4 
 
 Cochineal i7.- 3=9 
 
 Coefficients for calculating analyses 51 
 
 Cold way, titration 69 
 
 Congo red 329 
 
 Copper, Fleitmann's method 177 
 
 Cream of tartar 79 
 
 Cyanid potassium 127 
 
 Cyanogen '^i 
 
 Cylinder, graduated 38 
 
348 INDEX 
 
 PAGE 
 
 Decinormal solutions lo 
 
 Diastasic value of malt extract 293 
 
 Bichromate, potassium 230 
 
 analysis by means of 178 
 
 preparation of V.S 179 
 
 Digestion methods 233 
 
 Dioxid, hydrogen 223 
 
 — manganese 228, 234, 243 
 
 Direct percentage estimations 48 
 
 — ■ table of quantities for 102 
 
 Dissociation theory 18 
 
 Distillation methods 224 
 
 Double normal solutions 10 
 
 Eau de Javelle 222 
 
 Elements, list of xii 
 
 Empirical permanganate solutions, use of 144 
 
 — solutions II 
 
 End-reaction 16 
 
 Eosin 17 
 
 Erdmann's float 43 
 
 Erythrosin B 331 
 
 Factors 51 
 
 Fats, waxes and oils 275 
 
 Fehling's solution 287 
 
 end-point 290 
 
 Ferric alum solution 119 
 
 — chlorid i6o,*237 
 
 — salts 210, 234, 237 
 
 estimation of by means of permanganate 159 
 
 reduction of 159 
 
 Ferrous ammonium sulphate 140 
 
 ■ — carbonate saccharated 149, 183 
 
 — iodid, syrup of ' 120 
 
 — salts 181 
 
 — sulphate 148, 185 
 
 Ferrum reductum 150 
 
 Flask, liter 37 
 
 — measuring 37 
 
 Flasks, titration 3$ 
 
 Fluorescein 18 
 
INDEX 349 
 
 PAOE 
 
 Formaldehyde. 295 
 
 Free fatty acids 275 
 
 Galenical preparations 263 
 
 Gallein 329 
 
 Gasometric analyses 307 
 
 Gay-Lussac's method for haloid salts T 115 
 
 General methods of assaying drugs 261 
 
 General principles 4 
 
 Glucose 288 
 
 Gordin's modified alkalimetric assay 255 
 
 Gravimetric method, the i 
 
 Grethan's pipette 95 
 
 Grouvelle's bleaching fluid 222 
 
 Hsematoxylin 330 
 
 Halogens, free 240 
 
 Haloid acids 116 
 
 — salts no 
 
 estimation of with chromate as an indicator no 
 
 Mohr's method no 
 
 Helianthin 335 
 
 Hot way titrations 68 
 
 Hiibl's number 283 
 
 Hydriodic acid 100, 119 
 
 syrup 120 
 
 Hydrobromic acid 100, 117 
 
 Hydrochloric acid 9^ 
 
 — -^ action of, on permanganate 142 
 
 normal 61 
 
 standardization of 63 
 
 Hydrocyanic acid.. 123 
 
 Hydrogen dioxid 155. 223, 317 
 
 Hydrogen sulphid 205 
 
 Hydroxid and carbonate of alkali mixed 73 
 
 of alkali earths, mixed 87 
 
 — potassium 66 
 
 — potassium normal V.S 9° 
 
 — sodium "7 
 
 normal V.S 93 
 
 Hydroxids, alkali, estimation 64 
 
 . — of alkali earths S4 
 
350 INDEX 
 
 PAOE 
 
 Hydroxids, sodium and potassium mixed yj 
 
 Hypobromite solution for urea estimation 323 
 
 Hypochlorite 222 
 
 Hypophosphite, calcium 166 
 
 Hypophosphites 164 
 
 Hypophosphorous acid gp, 164 
 
 Hyposulphite, sodium 205 
 
 V.S. preparation of 211 
 
 Immiscible solvents 258 
 
 Indicator 16, 59 
 
 Indicators, classification of 24 
 
 — description of individual 328 
 
 — guide for the selection of 26 
 
 Indicator, requirements of a good 26 
 
 — sensitiveness to alkaloids 268 
 
 — theories of 20 
 
 Indirect oxidation, analysis by - 187 
 
 Inorganic acids 94 
 
 Instruments, calibration of 45 
 
 Introduction i 
 
 Inverted cane sugar 288 
 
 lodates 234 
 
 lodeosin 331 
 
 lodid, arsenous 197 
 
 — ferrous syrup of 1 20 
 
 — potassium ' 114 
 
 • — strontium 121 
 
 — zinc - 121 
 
 lodids Ill 
 
 — alkali 231 
 
 lodin, free 216 
 
 — purification of 188 
 
 — tincture 214 
 
 — titrations, use of sodium bicarbonate in '. 193 
 
 Iodine absorption number 283 
 
 — V.S. preparation of 188 
 
 Iodized starch test paper 240 
 
 lodometry 210 
 
 lodometiic estimationsj indirect 217 
 
 Ionization theory 18 
 
 Ions 18 
 
INDEX 3,-)l 
 
 PAOE 
 
 Iron, estimation of by stannous chlorul 245 
 
 — reduced 207 
 
 Javelle's water 222 
 
 Katz's method 261; 
 
 Kebler's- Keller method 261 
 
 Kingzett's method 223 
 
 Knop's hypobromite solution , 324 
 
 Kappeschaar's solution 269 
 
 Kottstorter number 2^6 
 
 Labarraque's solution 222 
 
 Lacmoid 331 
 
 Lacmus 333 
 
 Lactic acid 103 
 
 Lactose 288 
 
 Law of Boyle 309 
 
 — of Charles 309 
 
 Lead acetate 167 
 
 — peroxid 234 
 
 — salts 104 
 
 — subacetate 167 
 
 Lime, chlorid of 220 
 
 — chlorinated 220 
 
 Liter flasks 37 
 
 Lithium carbonate 71 
 
 — citrate 81 
 
 — organic salts of 77 
 
 Litmus 333 
 
 — tincture of 17 
 
 Lloyd's method 263 
 
 Lugol's solution 217 
 
 Lunge's pipette 95 
 
 Magnesium salts 86, 104 
 
 — sulphate, use of in permanganate titrations 143 
 
 Malt extract, diastasic value of 293 
 
 maltose in 291 
 
 Maltose 288 
 
 Mandarin-orange -" 33s 
 
 Manganates 2io> 224 
 
 Manganese dioxid , if'8, 228, 234, 243 
 
352 INDEX 
 
 PAGE 
 
 Manganous sulphate, use of in permanganate titrations 143 
 
 Measuring flask 37 
 
 Meniscus 42 
 
 Mercuric chlorid V.S., estimation of alkali iodids by 130 
 
 — salts 247 
 
 — sulphate, use of, in permanganate titrations 143 
 
 Methyl orange 17, 335 
 
 Mohr's burette 28 
 
 — salt 140 
 
 Multiples of atomic weights xiii 
 
 Neutralization analysis ; 36 
 
 — of acids 88 
 
 Nickel salts 104 
 
 Nitrate, silver V.S 106 
 
 Nitrates,, nitric acid in 315 
 
 — Pelouze method 169 
 
 Nitric acid 99, 119, 315 
 
 Nitrite, amyl 316 
 
 — ethyl 316 
 
 — sodium 314 
 
 Nitrites 162, 311 
 
 Nitrogen dioxid 316 
 
 Nitrometer, the 307 
 
 Nitrous acid 162 
 
 — ether 311 
 
 Normal oxalic acid : - ■ - 60 
 
 — solutions 7 
 
 Orange, methyl 335 
 
 Oils, fats and waxes 275 
 
 — iodin absorption number of 283 
 
 — table showing iodin absorption number of 286 
 
 Organic acids 100 
 
 ■ — salts of alkalies , 76 
 
 of the alkalies, table of fa,ctors for 82 
 
 of alkali earths 77 
 
 of lithium 77 
 
 Oxalic acid 103 
 
 and oxalates 151 
 
 decinormal 61 
 
 normal solution 60 
 
INDEX 353 
 
 PAGE 
 
 Oxalic acid, standardization of permanganate by 138 
 
 Oxid, antimonous 197 
 
 — arsenous ig6 
 
 Oxidation and reduction 133 
 
 — indirect analysis by 187 
 
 Oxy-chlor-diphenyl-quinoxalin 335 
 
 Paranitrophenol 22 
 
 Percentage rules for direct estimations 48 
 
 Permanganate, action of hydrochloric acid on 142 
 
 — potassium V.S 135 
 
 — solutions, empirical 144 
 
 — titration with in presence of hydrochloric acid 143 
 
 — typical analysis with 148 
 
 — volumetric analysis by means of 140 
 
 Peroxid, lead 234 
 
 Peroxids 210, 224 
 
 Phenacetalin -^337 
 
 — in estimating mixed alkali hydroxid and carbonate 74 
 
 Phenol 269 
 
 Phenolphthalein 17, 338 
 
 Phosphoric acid 97 
 
 Pinch cocks 32 
 
 Pipettes 33 
 
 Poirrier blue 339 
 
 use of in cyanogen assay 123 
 
 Poirrier's orange iii 335 
 
 Potassium acetate 80 
 
 — aisenite solution 197 
 
 — bicarbonate 70 
 
 — bi-iodate, advantages of for standardizing V.S 93 
 
 preparation of 93 
 
 — bitartrate 79 
 
 purification of 92 
 
 — bromate 236 
 
 — carbonate 69 
 
 — chlorate 236 
 
 — chromate T.S 18 
 
 — citrate 80 
 
 — cyanid 127 
 
 — dichromate 230 
 
 — ferricyanid T.S 18 
 
354 INDEX 
 
 PAGE 
 
 Potassium hydroxid 66 
 
 normal V.S no 
 
 — iodid 114 
 
 — permanganate V.S i ^r 
 
 — sulphite 1 204 
 
 — sulphocyanate V.S ion 
 
 — tartrate -j-j 
 
 — and sodium hydroxids mixed yc 
 
 tartrate. 
 
 79 
 
 Precipitation analysis 105 
 
 Preparation of normal oxalic acid solution 60 
 
 — of standard acid solutions en 
 
 Prollius' fluid 258 
 
 Puckner's method 262 
 
 Pyrogallol-phthalein 329 
 
 Ramsay's bleaching fluid 222 
 
 Reading of instruments 41 
 
 Reduced, estimation of substances readily 209 
 
 — iron .' 150, 207 
 
 Reducing agents 209 
 
 — sugars 288 
 
 Reduction and oxidation analysis 133 
 
 — methods, involving the use of arsenous acid V.S 238 
 
 stannous chlorid 244 
 
 Reichert number 277 
 
 Reichert-Meissl number 278 
 
 Resazurin 340 
 
 Residual titration ■ 15 
 
 Re-titration 15 
 
 Rochelle salt 79 
 
 Rosolic acid 17, 341 
 
 Rules for direct percentage estimations 48 
 
 — for finding percentage 52 
 
 Salicylate, sodium 81 
 
 Saponification number 276 
 
 Seminormal solutions 10 
 
 Separators 259 
 
 Shaking-out process for alkaloids 258 
 
 Silver alloys 129 
 
 — metallic , 1 29 
 
INDEX 355 
 
 PAGE 
 
 Silver nitrate, assay of by means of sodium chlorid V.S 128 
 
 sulphocyanate 129 
 
 V.S 106 
 
 — salts 127 
 
 Soda, clilorinated 222 
 
 Sodium acetate 81 
 
 — and potassium hydroxids mixed 75 
 
 tartrate i . 79 
 
 — benzoate 81 
 
 — bicarbonate 71 
 
 • use of in titrations with iodin '. 193 
 
 — bisulpiiite 205 
 
 — carbonate (aniiydrous) 71 
 
 (crystallized) 7° 
 
 normal V.S 86 
 
 pure, liow to make 86 
 
 — chlorid 112 
 
 — — preparation of pure 108 
 
 V.S 108 
 
 — hydroxid 67 
 
 V.S 93 
 
 — hyposulphite 205 
 
 V.S., preparation of 211 
 
 — nitrite 3^4 
 
 — salicylate 81 
 
 — sulphite 204 
 
 — tetrazo-diphenyl-naphthonate 3^9 
 
 — thiosulphate 2°S 
 
 V.S., preparation of 211 
 
 Spirit of ammonia 68 
 
 — of nitrous ether - 3" 
 
 Squibb's hypobromite solution 3^4 
 
 Standard solutions 7 
 
 — temperature ^3 
 
 Stannous chlorid, estimation of iron by means of 245 
 
 solution, preparation of 245. 248 
 
 use of in reduction methods 244 
 
 Starch =88 
 
 — as an indicator '92 
 
 — after inversion 291 
 
 — inversion by diastase 292 
 
 — iodized, test paper 240 
 
356 INDEX 
 
 PA.GE 
 
 Starch solution 102 
 
 Stating results ,■■, 
 
 Strontium iodid 121 
 
 — soluble salts of 86 
 
 Subacetate, lead i(,y 
 
 Sugars 287 
 
 Sugar in urine 289 
 
 Sulphid, hydrogen 205 
 
 Sulphids, insoluble 206 
 
 Sulphuric acid g(, 
 
 normal 64 
 
 Sulphites, estimation of alkalies in presence of 75 
 
 Sulphite, potassium 204 
 
 — sodium 204 
 
 Sulphites 200 
 
 Sulphocyanate, potassium V.S 109 
 
 — V.S., assay of silver nitrate by 1 29 
 
 Sulphurous acid 100, 200, 202 
 
 Syrup of hy driodic acid 1 20 
 
 Table, alcoholometric 304 
 
 — for correction of volume for the temperature 41 
 
 — of elements xii 
 
 — of factors of organic salts of the alkalies 82 
 
 — of normal factors for acids, alkalies and alkali earths 53 
 
 — or normal factors for oxids, etc 55 
 
 — of multiples xiii 
 
 — of quantities for direct percentage estimations 102 
 
 — of substances estimated by standard iodin solution 209 
 
 precipitation 131 
 
 permanganate and dichromate 186 
 
 — showing color changes of indicators 254 
 
 factors for alkaloids 254 
 
 iodin absorption number of oils 286 
 
 Tartar, cream of 79 
 
 — emetic 200 
 
 Tartaric acid loi 
 
 Tartrate, antimony and potassium 200 
 
 — potassium. 77 
 
 — sodium and potassium 79 
 
 Temperature, standards of 13 
 
 Test mixer 38 
 
INDEX 357 
 
 PAGE 
 
 Tetra-iodo-fluorescein 331 
 
 Theories of indicators 20 
 
 Theory, chromophoric 20 
 
 — ionization of indicators 20 
 
 — of Ostwald 20 
 
 Thiosulphate sodium 205 
 
 V.S., estimations involving use of 210 
 
 preparation of 211 
 
 — standardization of iodin with 189 
 
 — V.S., standardization of by dichromate 212 
 
 of by potassium bi-iodate 214 
 
 of by iodin 211 
 
 of by permanganate 216 
 
 Tin, Lowenthal's method 176 
 
 Titrate, to 14 
 
 Titrated solution 7 
 
 Titration, influence of concentration of V.S 40 
 
 of rate of speed ,. 40 
 
 of temperature 40 
 
 — residual ' 15 
 
 Titer 7 
 
 Tropaeolin D 335 
 
 — (O O) 341 
 
 Turmeric i? 
 
 — paper 342 
 
 — tincture 34^ 
 
 Urea apparatus, Squibb's 325 
 
 — estimation of by Doremus' ureometer 323 
 
 Ureometer, Doremus' 323 
 
 — Hinds-Doremus' 3^5 
 
 Urine, sugar in 289 
 
 Valence '2 
 
 Vegetable drugs, assaying of 257 
 
 Vielhaber's method for cyanogen 1 24 
 
 Violet, anthracene 329 
 
 'Volatile acids, vi^eighing of 95 
 
 — fatty acid, value 277 
 
 — solvents, influence of in alkaloidal assays 268 
 
 Volhard's solution 109 
 
 Volume strength of hydrogen dioxid 157 
 
358 INDEX 
 
 PAGE 
 
 Volumetric method, the -^ 2 
 
 — or standard solutions 7 
 
 Wilson's bleaching fluid 222 
 
 Zinc iodid 121 
 
 — salts 104 
 
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