BOUGHT WITH THE INCOME FROM THE SAGE ENDOWMENT FUND THE GIFT OP Sftenrg W, Sage X891 Cornell University Library arV19361 Short course in inorganic qualitative ch .. 3 1924 031 284 155 olin.anx Cornell University Library The original of tliis book is in tine Cornell University Library. There are no known copyright restrictions in the United States on the use of the text. http://www.archive.org/details/cu31924031284155 A SHORT COURSE IN INORGANIC QUALITATIVE ANALYSIS. FOB ENOINEEBINQ STUDENTS. BT J. S. C. WELLS, Ph.D., Instructor in Analytical Chemistry, Columtiia Universitl/. FIRST EDITION. FIRST THOUSAND. NEW YORK: JOHN WILEY & SONS. London: CHAPMAN & HALL, Limited. 1898. Copyright, 1898, BT J. S. C. WELLS. ROBERT DRCMHOND, ELBOtROTYFXR ANP FBINTEB, MBW XOSOL. PREFACE. The object of this book is to give a short but thorough course in inorganic qualitative analysis for the use of students who have only a limited time to devote to the subject. Experience has shown that Fresenius' manual, while invaluable as a work of reference, is too complicated and voluminous for the beginner, the great mass of information there given confusing him, and he becomes discouraged at his inability to assimilate it. In preparing the present work the idea has been, while still following the general plan of Fresenius, to give only that which seemed essen- tial to a clear understanding of the subject and to make it as concise as possible. For this reason only the more important reactions of the diiferent metals and acids have been given, and the separa- tions are presented in the form of schemes accom- IV PREFACE. panied by explanatory notes, and tables of scheme reactions. The latter have been found of much benefit in helping the student to understand the various reactions taking place in an analysis, as they show at a glance the effect produced by each reagent used. In the notes the aim has been to call attention to likely sources of error, and to give any expla- nations as to conditions, etc., that seemed neces- sary. The separations used differ in some cases from those of Presenius, particularly the separation of the first four groups in presence of phosphates, and that of the third and fourth with potassium hydrate when chromium is present ; also the scheme for the detection of the acids. All of these methods have been thoroughly tested in the labo- ratory by large numbers of students, and found to be very reliable and less complicated than those in general use. Chemical equations are freely used to explain the various reactions taking place in the individual tests as well as in the separations. In the introduc- tion a short explanation is given of these equa- tions and how to write them ; also the definitions of certain terms. This, of course, belongs more PBEFACE. V properly to general chemistry, but experience has proved the average student to be woefully igno- rant in this respect, and a brief exposition of the subject seemed necessary. No attempt has been made to give full description of apparatus, it being considered better for the teacher to do that in the class-room, and he can at the same time perform any of the more difficult tests that may seem to require such demonstration. There are some who consider it simply a waste of time to begin with the individual reactions of each metal before taking up the separations, but it certainly seems more logical to first study the reactions of the different elements group by group, and then proceed to separate them, rather than to reverse this order and begin with the separations. By restricting the tests to be per- formed to those actually used in the analysis the number made will not be very large. If time permits these may be increased later on by mak- ing use of other methods of analysis, thus bring- ing in new reactions, but this should not be attempted until the student is thoroughly familiar with at least one good method of separation for all the bases and acids. The work can be made VI PREFACE. more interesting by giving out for analysis as many natural or commercial products as possible, such as ores, alloys, etc. This can be done almost from the first by beginning with such minerals as calcite, dolomite, etc., and as the work progresses those containing a greater number of bases can be given, care being taken, however, to give only those compounds that are easily soluble in acids, and whose acid radicals do not interfere with the analysis. By thus making practical application of his knowledge the in- terest of the student is excited, and his work becomes a pleasure rather than a task. J. S. e. Wells. Columbia TJkivbrbitt, April, 1898. CONTENTS. SECTION I. PAGE iNTBODrCTION: Grouping of the Metals 3 Acids— Bases — Salts 5 Chemical Equations 9 Reactions of the Metals: Group I 19 " II 33 " III 53 " IV 64 '• V 110 " VI 143 SECTION II. The Acids: Group 1 188 II 318 " III 343 SECTION III. Complete Analysis : Preliminary Examination 367 Dissolving Solid Substances 374 Insoluble Substances 378 Metals ok Allots 379 Silicates 283 Insoluble Cyanides 383 Beaoentb 384 vii ABBREVIATIONS. Ppt. = precipitate. Fill. = filtrate, Sol. = solutioi ^^-<^ placed under a formula means that the compound repri sented is thrown down as a precipitate; when placed above, that it given off as a gas. QUALITATIYE ANALYSIS. INTEOPUGTIQN. Theee" are two principal divisions of analytical cHemistry, viz., qualitative analysis and quantita- tive analysis. " The former, as its name implies, is an analysis or "separation to "determine the quality of an unknown substance, that is, the constituents con- tained in it. The object of quantitative' analysis, on the other hand^ is to determine the quantity of each of the various constituents present. - For an intelligent study of qualitative analysis it iis very necessary t¥at the student have some knowledge of general chemistry. In his work in the laboratory, order, neatness, and absolute clean- liness in regard to apparatus are very essential to success. Skill in manipulation comes naturally to some, and all may acquire it to a greater or less degree by practice. The success of an analysis also depends "very 2 INTBODUCTION. greatly on the care and forethought exercised by the student in performing the various operations that are called for in the course of his work. He should make it a rule never to attempt any test or separation until perfectly familiar with the con- ditions necessary for a successful performance of the operation in hand. Before adding any re- agent, let him stop and ask himself the question, " Why do I add this, and what is it expected to do ? " If he can answer this question, then let him use it in the way best suited to produce the desired result; if not, he should go no further until he understands fully the reasons for its use. This is a common fault with very many students ; they perform a certain operation in a certain way because their text-book says so, but further than that they know nothing about it, and are perfectly satisfied if by chance they get cor- rect results. If they would only remember that the object in making all these tests and separa- tions is to make them familiar with certain chem- ical facts and their application, they would per- haps realize that the first and most important point in any analysis is a thorough knowledge of the principles involved. Correct results will soon follow when this has been gained. Section" I. THE METALS AND THEIR REACTIONS. It has been found most convenient in analysis to treat the bases and acids separately. In inorganic chemistry the bases comprise the metals, and the acids the non-metallic elements (with a few exceptions). Experiment has shown that certain bases or ~ acids act in the same manner when brought in contact with some particular substance of known composition. By making use of these substances, or reagents, as they are called, we are enabled to separate the metals and acids into groups. These groups can then be subdivided and separated into their several constituents. This grouping of the different elements is of great advantage, for it is obvious that if we add a group reagent and get no precipitate, the group of substances affected by this particular reagent cannot be present, and no further testing for them is necessary. 4 SEPARATION INTO GROUPS. The classification of the bases adopted in this book is as follows • First Group (Px)tassium, Sodium, and Ammo- nium). — These nietals form very few insoluble compounds, and are not precipitated by any.- of the reagents used to throw down the metals of the other groups. Second Group (Barium, Strontium, Calcium, and Magnesium). — This group is distinguished from the first by the insolubility of its normal carbon- ates and phosphates, as well as other salts, and from the succeediri'g groups by the fact that it is not precipitated by either hydrosiilphuric acid or ammonium sulphide. Thwd Group (Aluminium,' Chromium, and Titanium). — These metals are not precipitated by hydrosulphuric acid from acid solution, and with ammonium sulphide in neutral solutions yield precipitates of hydroxide. Fou/rth Group (Iron, Manganese, Zinc, Cobalt, and Nickel). — The metals of this group are not pre- cipitated by hydrosulphuric acid from solutions acid with the mineral acids. Ammonium sulphide in neutral or alkaline solutions precipitates them as sulphides. Fifth Group (Silver, Lead, Mercury, Bismuth, ACIDS — BASES— SALTS. 5 Copper^ and Cadmium). — Precipitated as sul- phides from alkaline, .neutral, or acid solutions by hydrosulphuric acid. The precipitate is insolu- ble in ammonium sulphide. Sixth Group (Tin, Antimony, Arsenic, Gold, and Platinum). — Precipitated by hydrosulphuric acid from acid solutions. The precipitate is solu- ble in ammonium sulphide, which distinguishes this group from the third, fourth, and fifth groups. "We might at once proceed to the separation of the metals into the different groups just men- tioned, but this would be of no advantage until we know something of the individual character- istics of the various members of each group. Until these have been studied we are in no posi- tion to distinguish the several metals from each other. Before taking up this portion of our work it will be well, however, to give a little time to the explanation of certain terms that are frequently used, such as base, acid, etc.,* and to the study of chemical equations. If we examine the two compounds hydrochloric acid (HCl) and potassium hydroxide (KOH) we find that they differ veiy decidedly from each other. The former has an acid or sour taste, and changes the color of many vegetable substances. The lat- * For a more complete explanation of these terms see Remsen'a 6 ACIDS — BASES — SALTS. ter has an alkaline taste, and restores the color of vegetable substances that have been changed hy the acid. Their characteristics in every way are opposed one to the other, and if they are brought together they neutralize each other, yielding a compound possessing none of the characteristics of either the acid or the hydroxide. These two compounds serve to represent the two great classes of compounds called acids and bases. Besides the two Just mentioned, there are many others whose characteristics are equally distinct and which are easily placed in one class or the other. There are some substances, however, whose prop- erties are not so decided and which act as acids towards some bases and like bases towards some acids. Like hydrochloric acid, all acids contain hydro- gen, which hydrogen is replaceable by a base. Some acids contain besides this hydrogen a cer- tain part that is not so replaceable : such, for ex- ample, is acetic acid, H(C2H302), in which three of the hydrogen atoms cannot be replaced by metals. In acids containing oxygen the acid hydrogen is supposed to be in combination with one of the oxygen atoms as hydroxyl (OH^, and it is only ACIDS— BASES— SALTS. 7 the hydrogen so combined that is replaceable by bases. In sulphuric acid we have two hydroxyls and two replaceable hydrogen (S02(OH)2 or H2SO4), in phosphoric acid three, in silicic acid four, etc. Acids containing one replaceable hydrogen are called monobasic, those having two are said to be dibasic. There are besides these tribasic and tetrabasic acids. Bases. — As already stated, the bases have prop- erties just the reverse of the acids. Among the inorganic compounds they usually consist of hy- droxyl in combination with a metal (except NH4OH), as KOH, NaOH, Ba(0H)3, etc. The chief characteristic is their power of com- bining with acids to form neutral compounds. Salts, — When an acid and a base act on each other they neutralize each other's properties to a greater or less extent. In the case already given of hydrochloric acid and potassium hydroxide we have a complete neutralization and the formation of a neutral substance possessing none of the char- acteristics of either the acid or the base. This substance is called a salt. Its formation may be represented by this equation : KOH + HCl = KCl + H2O. 8 ACIDS — BASES— SALTS, In the case of a bibasic acid, sucli as sulphuric, we may get either a neutral or an acid salt : 2K0H + H,S04 = KSO, + 2H,0 ; ROH + H,S04 = KHSO4 + H2O. The first K2SO4 is called a neutra,! or normal salt. It is a salt formed by the displacement of all of the acid hydrogen of an acid by an equiv- alent amount of base. The second KHSO4 is called an acid salt because it still contains hydro- gen that may be replaced by a base. Tri- and tetrabasic acids also form acid salts. It is well to remember that normal salts are not always neutral to test-paper. A strong acid in combination with a weak base gives a salt that turns blue litmus red, and a strong base in combination with a weak acid has the re- verse effect, turning red litmus blue. Besides normal and acid salts, there is still another class called basic salts. These are the reverse of the acid salts — that is, they contain a greater amount of base than there is acid to neu- tralize. Magnesium affords a good example of such a salt. The normal carbonate has the com- position MgCOg, but besides this it forms basic carbonates, as 4MgC03Mg(OH)2. This compound CHiEMICAL REACTIONS. 9 still contains basic hydroxyl, and for this reason it is called a basic salt. We have already had occasion to use chemical formulas and equations, and a little time will now be given for a brief expla,nation of their signifi- cance., CHEMICAL EEACTIONS. All .chemical changes may be expressed by means . of reactions written in . the . form of an equation,. e.g.: - - BaCl3 + Na3S04 = BaS04+2NaCl. Now let us examine this equation and see what it is intended to express by it. In the first place it is a kind of shorthand for chemical names ; instead of writing out the words " barium chloride " we denote it much more easily by using the chemical symbol belonging to each element contained in the compound. Thus in the example ^iven Ba stands • for barium and CI for chlorine. Perhaps you may ask why we write it BaClj and not simply BaCl. This brings us to another important property of symbols, viz., that they not only represent the elements themselves, but they also represent their atomic weights, as compared with the atom of 10 CHEMICAL EEACTIONS. hydrogen which is taken as the unit. Now in the . example under discussion it has been found by quantitative analysis that it contains 137 parts by weight of barium and 71 parts of chlorine. The weight of the barium atom has been found to be 137 times that of the hydrogen atom, hence in this compound we have an amount of barium equal to one atom. The chlorine atom has been found to be 35.5 times as heavy as that of hydrogen, hence if BaCla contains 71 parts of chlorine to one barium atom (137) it must contain 71 -r- 35.5 (weight of 1 atom of chlorine) equals 2 atoms of chlorine. This we denote by writing the figure 2 at the lower right-hand corner of the symbol of the ele- ment ; hence we see that BaCla means in the first place the chemical compound barium chloride; secondly, that it represents definite weights of the constituent elements, and also the number of atoms of each element in the molecule, or at least their relative proportions. When a number is written before the symbol representing a molecule, as 2NaCl, it means two molecules. If we should perform the chemical operation expressed by the equation given, we would find CHEMICAL REACTIONS. 11 that 208 (Ba 137 + 01,71) parts by weight of barium chloride and 142 (Nag 46 + S 32 + O4 64) parts of sodium sulphate would produce 233 (Ba 137 + S 32 + O464) parts of barium sulphate and 117 (2(Na 23 + CI 35.5)) parts of sodium chloride. It should be remembered in all equations repre- senting chemical change, that the number of atoms on one side of the equation must be exactly equal to those on the other ; nothing can be gained or lost. Having gained some idea as to what an equa- tion means, let us see if they differ from each other in any important respects. In the equation cited it is evident that we have a simple interchange of the elements contained in the two substances used, the barium replacing the sodium and the sodium the barium. Such a trans- position is called metathesis, and to this class of equations belong a very large number of chemical reactions. It is a general rule, that if we mix two solu- tions capable of forming by exchange or transpo- sition a compound insoluble in the mixture, such insoluble compound will be produced and precipi- tated. In the case just given, although BaCl^ and Na2S04 are both soluble in water, yet when we .12 CHEMICAL REACTIONS. mix. them BaS04 is precipitated, because it is in- soluble in water. Other forms of chemical change are thosa of combination or synthesis, as * - ' • -H8 + Ci;=2HCl, 2C+02 = 2CO; tjiose. of dissociation or analysis, as CaCOa (on ignition) = CaO H- COg , 2AUCI3 " =3Au + 3Cl2; • '' and those of oxidation and reduction, the latter two representing a very important series of reac- tions. Oxidation, strictly speaking, would, m^ean an increase in the quantity of oxygen contained in a body, but the term is often used when oxygen .takes no part in the work, . as when FeCJa is changed to FejClft by means of. chlorine : 2FeGla+ei2-=FeaGl6. Although oxygen does not enter into the reaction, we .say the FeCl^ has been oxidized — meaning that it has been changed from a salt corresponding to FeO to one corresponding to FcgOg. * Generally in tliese reactions there -is dissociation as -well as com- bination. In tie example given the molecules of hydrogen, chlorine, and oxygen must split up into separate atoms -before combining. -CHEMlCi;L REACTIONS. l3 ""As an example of an Dxidatidn equation let us tatetlie one showing the oxidation of FeS04 to Pe2(S04)3 by KgMnA* and HaSO^: .■.-_-«Fej(K(>i)2 + JC^n2084--8HaS04 = , - &Fr2(S04)!iJ-K2S04 + 2MnSOc+.8H-A „ - ■ In ord,er to write such an equation as this, it is necessary ta know, first, how much oxygen is ^needed to change the body from the lower oxide to the higher; -second, how much oxygen we can get from each molecule of our oxidizing agent , and what are the. by-products formed by the reduc- tion or decomposition of the latter ; also what are the other products, if, any, that are formed' by the chemical- changes going on. In order to deter- mine the amount of oxygen necessary, we will first examine the composition" of the body to be oxidized,^ FeS04.' .-; -Ferrous sulphate probably contains two atoms ^f iron in the molecule, as shown in- the following graphic formulae : ^ -.j„ rs- Ferr^ou^- m " ~* Ferrous Fe=0 ^^^^^ -Fe = SO ^,"^P^^*1 * Also written KMnOj. 14 CHEMICAL REACTIONS. The composition of ferric sulphate, the product of the oxidation, is shown by the following formulae : „/F^ = Ferric „„/1'« = so, j,^^„ ^Fe = O ""'*'' '^Fe = SO. »"'?'>•*«• On comparing the formula for ferrous oxide with that for ferric oxide, we see that in the re- action every molecule of the former has taken up one more atom of oxygen in order to become fer- ric oxide, or, in other words, the ferric oxide con- tains one more atom of oxygen in the molecule than does the ferrous ; hence every atom of oxygen yielded by the oxidizing agent will oxidize one molecule of ferrous oxide (Fe^Oa) to ferric oxide (Fe^Os). Next let us see in what way the permanganate acts with the sulphuric acid. It has been found that it is decomposed or reduced, when in pres- ence of an oxidizable substance, as shown in the following equation : Ks,Mn,08 + 3H,S04 = 2MnS0, + KgSO^ + 50 + 3HA From this we see that every molecule of the KgMUijOg will yield five atoms of oxygen, free to CHEMICAL EEACTIONS. 15 enter into combination with the iron, and we have already determined that each molecule of ferrous oxide (Fefiz) requires one atom of oxygen to change it to ferric oxide (FegOg) ; hence five atoms of oxygen will oxidize five molecules of the FcaOa to FcaOg. Next, how much H2SO4 will be needed besides that already contained in the fer- rous sulphate. Ferric oxide when it combines with H3SO4 does so in the following manner : FeA + 3H2SO4 = Fe,(S04)3 + SH^O. Hence it is evident that for each molecule of Fe^Oa we shall need three of HgSO^ ; but as there are already two molecules present in the ferrous sulphate, we shall actually need but one more for every molecule of iron oxidized ; and for five, the amount oxidized by one molecule of KgMngOg, five H8SO4 will be required — making eight in all, with the three needed for combination with the potassium and manganese of the permanganate. We have now determined the quantities of each reagent taking part in the reaction, and also the quantities of the products. We will now take another reaction, in which at first sight the action of the oxidizing agent is not so plain. If we heat chromic hydroxide with 16 CHEMICAL KEACTiOJS'S. a solution of sodium carbonate and broffli'ne, the clirdmium will be oxidized to chromic acid, al- though, of course, the bromine itself contains no oxygen. Let us first write down the substances taking part in the reaction and the products formed : Cr2(OH)6 + Na^COg + Br = 2Na,Cr04 + NaBr + CO^ + H^O. The substance to be oxidized, Cr2(OH)6, con- sists of CrgOj + 3H3O ; the result af the "oxida- tion,. NagCrO^, consists of NagO + CrOj. As, the molecule of. Cr3(OH)6 contains two atoms of chro_- mium, we. must, if all the chromium is oxidized, produce two molecules of Na2Cr04, .In Na3Cr04 the CrOg is the acid anhydride, and it has formed the salt by acting on NagCOa as fol- lows : . Na^COs + CrOa = Na^CrO^ + COj. We see, therefore, that the product of the oxi- dation is really CrOg and that the Na2Cr04 results from the action of the CrOg on NaaCOg. Now^ if we start with a molecule of CrgOg and obtain as a result of .the reaction two molecules of CrOg , it is evident, as shown by the following formulae, CHEMICAL BEAOTION. 17 that tte CrgOg lias taken up three more atoms of oxygen in the change to 2Cr03; hence for every molecule of Gr^O^ oxidized we must have three more atoms of oxygen. How does the bromine furnish it ? It has been found that Br in alkaline solution acts as follows when oxidizable matter is present : SNagCOg + SBrg = 6NaBr + SCO^ + 30. 6 parts of bromine and 3 of sodium carbonate will thus give us sufficient oxygen for the oxidation of one molecule of Cr8(0H)e to 2Cr03. The 2Cr03 formed then combines with more of the sqdium carbonate to form sodium chromate as shown already. We will then need, besides the three molecules of carbonate that react with the bromine, two more to combine with the 2Cr03, making five in all. Hence the complete equation will be Cr2(OH)6 + SNa^COg + SBr^ = 2Na2Cr04 + 6NaBr + SCO^ + SH^O. An equation representing oxidation generally 2'epresents reduction as well — reduction meaning just the reverse of oxidation. In the case just considered, of the oxidation of ferrous salts by 18 CHEMICAL REACTIOIf. permanganate, the latter is reduced — that is, loses oxygen and becomes manganous sulphate — while the iron salt is oxidized. Sulphurous acid is a/ strong reducing agent owing to the facility with which it takes up oxygen and becomes sulphuric acid. This is shown in its action on ferric salts; thus : Fe,(S04)3 + SO3 + 2H,0 = 2FeS04 + 2H2SO4. Stannous chloride (SnClg) is another active re- ducing agent. When added in excess to a solu- tion of HgClg it reduces the latter to metallic mer- cury; at the same time it is oxidized to SnCl4; thus : > 2HgCl2 + 2SnCl2 = 2Hg + 2SnCl4. Many more examples showing oxidation and re- duction might be given, but sufficient have been shown to serve as types of all. FIRST GROUP. POTASSIUM, SODIUM, AND AMMONIUM. Characteristics of the Group. — The chief char, acteristic of this group, analytically, is the fact that all its salts, with very few exceptions, are soluble in water. The normal carbonates of the group are the only ones of all the bases that are soluble in this menstruum. The solubility of the sulphides also serves to distinguish this group from all of the other groups except the second. The solubility of the hydrates is of great analytical importance, for owing to their exceedingly strong basic properties we are enabled to precipitate nearly all of the other metals as hydroxides, by the addition of one of these hydrates to a solution of the metal. Solutions of the hydrates and of the carbonates turn red litmus blue and turmeric paper brown.* * Litmus is a weak vegetable acid, dissolving in water to a red liquid. Its salts are blue. Litmus paper is colored with the red solution of the acid or the blue solution of one of its salts. 19 20 POTASSIUM. POTASSIUM. At. wt. 39.1 ; sp. gr. 0.865. Potassium is a white lustrous metal. It oxi- dizes very rapidly on exposure to the air, and if thrown on water it takes fire spontaneously, burning with a purple flame, and the water is found to have an alkaline reaction. 2K + 2H,0 = 2K0H + H^. Potassium should always be kept covered with petroleum to protect it from the air. Potassium hydroxide and its salts are not vola- tile at a dull red heat. As already stated, very few of them are insoluble, the only ones being the platinichloride, acid tartrate, silicofluoride, and potassium cobaltic nitrite. The first and second are the ones generally used as tests for potassium. Potassium PlatinicMoride (^K^PtCl^.-^li hy- drochloroplatinic acid be added to neutral or acid solutions of potassium salts, not too dilute, a yel- low crystalline precipitate of the platinichloride is formed. Very dilute solutions are not precipitated. The test is best made by adding hydrochloropla- tinic acid to the solution to be tested in sufficient quantity to give a yellow color, the solution is then evaporated on the water-bath nearly to dry- POTASSIUM. 21 ness, and alcoliol added ; if any potassium is pres- ent it will be found as a yellow crystalline precip- itate. In case sodium salts are present care must be taken not to carry the evaporation to complete dryness, for sodium forms with the platinic salt a corresponding compound containing water of crys- tallization, which is soluble in alcohol so long 'as it retains this water, but if it be driven out by carrying the evaporation too far, then the sodium platinichloride also remains as a yellow residue. 2KC1 + HaPtCls = K^PtCle + 2HC1. The precipitate is soluble in 110 parts of cold and in 19 of boiling water. In alcohol it is very much more insoluble, requiring 12,000 parts of absolute alcohol to dissolve it. On ignition it leaves -a residue of potassium chloride, and platinum. Hydrogen Potassium Tartrate {KHC^H^O^. — If tartaric acid be added to neutral or alkaline solutions (in the latter case the reagent must be added until the solution becomes acid) of potas- sium salts, a white crystalline precipitate of the acid tartrate is thrown down (a). The precipi- tate is soluble in the mineral acids (c) and in alkalis {h~). 22 POTASSIUM. (a) KCl + B.,CJI,0, = KHQH^Oe + HCl ; (b) KHC.H^Oe + KOH = K^C^HA + H,0 ; (c) KHC^H^Oe + HCl = KCl + HAHA- Equations a and c are apparently contradictory, but this is explained by the fact that in the test a the tartaric acid is added in excess (that is, more than sufficient to combine with the potassium), and in c the hydrochloric acid predominates. How- ever, owing to the solubility of the precipitate in hydrochloric acid, the reaction as given in a is never complete, some of the potassium always re- maining in solution. For this reason it is better to use as the precipitant a solution of hydrogen sodium tartrate, which forms in the reaction so- dium chloride instead of hydrochloric acid. The sodium chloride has very little, if any, solvent action on the precipitate : KCl + NaHC^H^Oe = KHC^H^Og + NaCl. Whichever precipitant is used, the solution should be concentrated and cold. Vigorous shak- ing helps the precipitation very much. At 10° C. one part of the precipitate requires 250 parts of water to dissolve it, but it is soluble in 15 parts of boiling water. SODIUM. 23 Flaine Test. — If a drop of a potassium salt (best the chloride or sulphate) on the loop of a platinum wire be held in the outer flame of the Bunsen burner the flame will be colored violet. The presence of sodium salts prevents the color from being seen. By interposing between the flame and the eye a piece of deep cobalt-blue glass or an indigo prism the yellow sodium flame is cut off and the violet of the potassium is visible. SODIUM (Na). At. Wt. 23 ; sp. gr. .972. Sodium is a soft silver-white metal greatly re- sembling potassium in all its properties. The properties of the hydroxide and its salts closely resemble the corresponding potassium compounds, with very few exceptions. The sodium salts are even more soluble, there being only one that is at all insoluble in water, viz., the pyroantimonate. Sodium Pyroantirrionate. — If a freshly pre- pared solution of potassium pyroantimonate* be added to a moderately concentrated solution of a sodium salt having a neuti'al or slightly alkaline reaction, a white crystalline precipitate of sodium pyroantimonate is thrown down more or less * Prepared by dissolving a little of the solid salt in hot water and filtering from any insoluble residue. 24 SODIUM. quickly, depending on the concentration of the liquid. By making the test in a watch-glass and stirring thoroughly with a glass rod the forma- tion of the precipitate is much hastened. Acid solutions should always be evaporated to dryness, if possible, to remove the free acid, or if this cannot be done they should be neutralized with potassium hydroxide or carbonate. Free acid decomposes potassium pyroantimonate, giv- ing a white amorphous precipitate of metanti- monic acid (HSbOg). In making this test no other salts than those of potassium and sodium should be present. In dilute solutions the precipitate will not form for a Jong time, often as much as 12 hours being required. Sodium pyroantimonate is soluble in 300 parts of boiling water. The following equation repre- sents its formation : 2NaCl + KJIShO, = NagHaSbaOj + 2KC1. Flame Test. — Sodium salts impart to the Bun- sen flame an intense yellow color, even when pres- ent in very minute quantity. The presence of potassium salts does not interfere with this test. AMMONIUM. 25 It must be remembered that the merest trace of sodium salt will color the flame yellow. Simply- touching the platinum wire with the fingers is sufficient to give a decided test. If the sodium compound is present in appreciable quantity it gives a color to the flame that does not disappear quickly. AMMONIUM (NH,). Ammonium is known only in combination (e.g., NH4CI). Ammonia, NH3, is a gas at or- dinary temperatures, and is very soluble in water. In this solution it is supposed to exist as the hy- droxide NH4OH. On heating the solution the ammonia is set free again as gas. Ammonia combines with all the acids to form ammonium salts, e. g. : NH3 + HCl = NH4CI ; 2NH3 + H,S04 = (NH4)3S04. All ammonium salts are volatile at a low red heat, which is a distinction between t];iem and sodium and potassium salts, the two latter not being volatilized at that temperature. The solubility of the ammonium salts corre- sponds very closely with those of potassium, the 26 AMMONIUM. only difficultly soluble ones being the platini- chloride, the acid tartrate, and certain mercuram- monium compounds. Liberation of Ammonia Gas. — The most char- acteristic test for ammonium compounds depends on the fact that they easily evolve ammonia (NH3) when treated with suitable reagents. If an ammonium salt be triturated with slacked lime and a few drops of water, or heated with solution of sodium or potassium hydroxide, am- monia gas is set free (a and V), which may be recog- nized by its pungent odor, by turning moistened red litmus paper blue, and by the dense white fumes formed when brought in contact with any volatile acid, as hydrochloric, nitric, or acetic (c). This last test is best made by moistening a glass rod with a drop or two of the acid and holding it over the test-tube or other vessel in which the ex- periment is being made. Hydrochloric acid is the most sensitive in this way, but it must be remem- bered that the concentrated acid itself fumes in moist air. For this reason, acetic acid is more re- liable. Another very delicate test for free ammonia is to moisten a piece of filter-paper with mercurous nitrate and hold this over the test-tube or beaker AMMONIUM. 27 in which the ammonia is being evolved. The paper will become black or gray, owing to the action of the ammonia on the mercurous salt (d). (a) 2NH4Cl+^Ca(OH), = CaCl2 + 2OTC+2H,0; (b) NH,C1 + KOH = KCl + NH3 + H2O ; (c) NH3 + HC1 = NH4C1; id) Hg,(N03), + 2NH3 = NH,Hg,N03 (black) + NH4NO3. Presence of other bases does not interfere with the liberation of ammonia from its compounds by the hydrosAdes of the alkalis or alkaline earths. Ammonium Platinichloride. — Hydrochloropla- tinic acid precipitates ammonium salts under the same conditions as for potassium salts, giving a yellow crystalline precipitate, which is rather more insoluble than the potassium platinichloride. On ignition (best after addition of ammonium ox- alate) it leaves a residue of metallic platinum (distinction from the potassium salt). 2NH4CI + H^PtCle = (NH4),PtCl6 + 2HC1. * Hydrogen Ammonium, Tart/rate. — Hydrogen so- dium tartrate added to concentrated solution of ammonium salts gives a white crystalline precipi- 28 AMMONIUM. tate of hydrogen ammonium tartrate. Tartaric acid precipitates only veiy concentrated solutions. The precipitate is much more soluble than the cor- responding potassium salt. NH4CI + NaHC4H406 = (NH4)HC4H406 + NaCl. Dimercurammoniuin Iodide. — For all ordinary tests the reactions already given are sufficient, but in cases where it is necessary to test for exceed- ingly minute quantities of ammonia (as in potable waters) the Nessler test, as it is called, is the most delicate. This is made in the following manner : To the solution containing free ammonia or an ammonium salt is added a solution of potassium niercuric iodide containing an excess of potassium hydrate (Nessler reagent*). If the ammonia is present in any quantity a brown precipitate is formed, but in very dilute solutions it gives only a brown or yellow color. 2(2KI, Hgis) + NH3 + 3K0H = NHgJ, H,0 + 7KI + 2H,0. *Nessler's reagent is made by dissolving 16 grams of mercuric chloride in about 500 c.c. of water. Add this slowly to a solution of 35 grams of potassium iodide in 200 c.c. of water, until a precipitate begins to form. Tten add 160 grams of potassium hydrate and dilute to one litre. Finally, add strong solution of mercuric chloride until the red precipitate of mercuric iodide just begins to be permanent. Do not filter, but allow any precipitate to settle. AMMONIUM. 29 The presence of salts of the alkalies (except cyanides and sulphides) does not interfere with this test, but care must be taken to see that the solution is neutral or alkaline before adding the reagent. Salts of the alkaline earths should be re- moved by adding just sufficient of a mixture of about 1 part potassium or sodium hydrate and 2 parts of sodium carbonate to completely precipi- tate them. The precipitate is allowed to settle and the clear liquid tested with the Nessler re- agent. SEPAEATION OF FIEST-GEOTJP METALS. We have found by our study of the several metals of the group that potassium and ammo- nium salts are both precipitated by hydrochloro- platinic acid and hydrogen sodium tartrate, so we cannot use these reagents as a distinctive test for either metal when both may be present in the same solution. Again, the test for sodium with pyroantimonate requires that no other salts than those of potassium may be present, and as it adds potassium to the solution, we cannot, of course, test in it for that metal. It is evident, then, that the first step in the sep- 30 SEPARATION. aration will be to determine whether ammonium salts are present, and if found, to remove them by suitable means. The test with slaked lime or potassium or sodium hydrates affords us a ready means of test- ing for ammonia, the test being made on a sepa- rate portion of the mixture. If found, we can make use of the fact that ammonium salts are vol- atile on ignition, by taking a fresh portion of the solution and evaporating to dryness and then igniting carefully until no more fumes of ammo- nia salts are given off. The residue will now con- tain only the sodium and potassium salts, and by dissolving it in a very little water and dividing into two portions we can test one with pyroanti- monate for sodium and the other with hydrochlo- roplatinic acid or hydrogen sodium tartrate for potassium. It is well to test the residue left after driving out the ammonium salts by taking a little of it on a platinum wire and making the flame test. This simple test will often show at once which metal is present; for instance, if it should give a violet flame, this would prove the presence of potassium, and it would be unnecessary to look further for sodium, as its flame would have shown even in SEPARATION. 31 preisence of much potassium. If the residue gave a strong yellow flame that took some time to burn off, that would be sufficient for sodium, and the potassium could then be determined by use of the blue glass or by precipitation with hydrochloro- platinic acid. SECOND GROUP. BAEIUM, STRONTIUM, CALCrCTM, MAaNESIUM. The insolubility of the normal carbonates and phosphates of this group serve to distinguish it from the first group, and the fact that it is not precipitated by hydrosulphuric acid nor ammo- nium sulphide distinguishes it from the succeeding groups. The solubility of the hydroxides in water de- creases in regular order from barium hydroxide, which is easily soluble to magnesium, hydroxide, which is very insoluble. The reverse is true of the sulphates, magnesium sulphate being easily soluble, calcium sulphate difficultly so, strontium sulphate much more insoluble, and barium sul- phate insoluble not only in water but in acids. The hydroxides and salts of this group are all colorless (unless the acid be colored as chromic acid), and are not volatile at a red heat. 38 barium:. 33 BARIUM (Ba). At. wt. 137.4. Barium is found in considerable quantities as barite or heavy spar (BaS04) and as witherite (BaCOg). Almost all of the barium salts are insoluble in water, but with the exception of the sulphate and silicofluoride they are soluble in acids. The chloride and nitrate are nearly insoluble in al- cohol, and insoluble in a mixture of equal parts of alcohol and ether. They are also insoluble in sti'ong hydrochloric or nitric acids ; for this rea- son, if either of these acids be added to a concen- trated solution of a barium salt a white precipitate will form, which, however, dissolves readily on diluting with water. Bariwm Hydroxide {Bai^OH^^. — Sodium and potassium hydrates (free from carbonate) give no precipitate with barium salts imless the solutions are concentrated. Ammonium hydrate causes no precipitate. Bariumi Carbonate {^BaGO^. — The carbonates of potassium, sodixim, and ammonium precipitate barium salts as a white amorphous precipitate. The precipitate is slightly soluble in ammonium chloride. It is easily soluble in all acids. 34 BARIUM. BaCla + (NH4)8C03 = BaC^ + 2NH4CI ; BaCOg + 2HC,H302 = BaCC^HgOg), + CO^. Barivm, Sulphate {BaSO^. — Sulphuric acid or soluble sulphates give, even in very dilute solu- tions of barium salts, a white pulverulent precipi- tate of barium sulphate. The precipitate is insol- uble in water and dilute acids. A solution of calcium sulphate in water serves as a useful reagent to distinguish between barium and strontium salts. With the former it gives an immediate precipitate, unless the solution be very- dilute, but with strontium the precipitate forms only slowly. BaCla + H2SO4 = BaSO, + 2HC1 ; BaCl, + CaS04 = BaS04 + CaClg. Barium Ohromate {BaCrO^. — Solutions of barium salts containing no free mineral acid are precipitated by potassium chromate or dichro- niate, giving a yellow precipitate which is very insoluble in water and in ammonia. It is slightly soluble in ammonia salts, in chromic acid, and STEONTIUM. 35 acetic acids. The solubility in acetic acid in- creases with the strength of the acid. / BaClg + K2Cr04 = BaCr04 + 2KC1. Solulle pliosphatea and oxalates precipitate neu- tral or alkaline solutions of barium salts as white precipitates of phosphate («) and oxalate (h). The phosphate is soluble in hydrochloric, nitric, and acetic acids,, the oxalate in hydrochloric and nitric acids, but not in acetic except when first precipitated. {a) BaCl^ + Na3HP04 = BaHPO^ + 2NaCl ; (5) BaCl^ + (NH4),C,04 = BaCA + 2NH4CI. Flame. — Barium salts (best the chloride or sul- phate) when held in the Bunsen flame on a loop of platinum wire impart to it a yellowish-green color. STRONTIUM (Sr). At. wt. 87.6. Strontium salts in their reactions very closely resemble the corresponding barium compounds. The hydroxide is somewhat less soluble in water than barium hydroxide. The chloride is deliquescent, and is soluble in absolute alcohol 36 STEONTITJM. (distinction from barium). The nitrate is nearly- insoluble in absolute alcohol and in a mixture of absolute alcohol and ether (distinction from calcium). The hydroxides and carhonates of the alkalies give the same reactions with strontium salts as with those of barium. Strontium Sulphate {SrSO^. — Sulphuric acid and soluble sulphates precipitate strontium salts, giving a white precipitate of sulphate. The pre- cipitate is much more soluble than barium sul- phate, both in water and in acids, and for this reason it precipitates only slowly from dilute solutions. With calcium sulphate the precipita- tion is always slow except in very concentrated solutions (distinction from barium). The appli- cation of heat helps this precipitation as well as all others of strontium as sulphate. Soluble phosphates and oxalates precipitate neu- tral or alkaline solutions of strontium salts, yield- ing white precipitates of phosphate and oxalate corresponding to the barium phosphate and oxa- late, and having about the same solubilities. Potassium dichromate does not precipitate strontium salts even from concentrated solutions. Normal chromates in neutral solutions not very CALCITTM. 37 dilute give, on standing for some time, a yellow precipitate of chromate. Solutions acid with acetic acid are not precipitated (distinction from barium). Flame. — Strontium salts (preferably the chlo- ride) impart a crimson color to the Bunsen flame that is very characteristic. Strontium sulphate and salts of fixed acids should be moistened with concentrated hydrochloric acid before briuging them into the flame. « CALCIUM (Ca). At. wt. 40. Calcium is found in enormous quantities in the form of carbonate (limestone, mai'ble, etc.) and as sulphate (gypsum), and is an important constitu- ent of many silicates. Calcium oxide (quicklime) and the hydroxide Ca(0H)2 (slaked lime) are considerably more in- soluble than barium and strontium oxides and hydroxides. Calcium chloride and nitrate are soluble in absolute alcohol and in a mixture of alcohol and ether. The hydroxides of potassium, sodium, and am- monium, and their carbonates, act in much the same^way with calcium salts as with those of barium and strontium. 38 CALCIUM. Calcium carbonate when first precipitated is amorphous, but on standing, or more quickly by heating, changes to the crystalline form. It is easily soluble in acids, even in carbonic, and is also somewhat soluble in ammonium chloride. From its solution in carbonic acid it is precipi- tated on heating. Ca.Cl2 + Na,C03 = CaCOs + 2NaCl; CaCOg + GO, + H^O = CaH,(C03)2 ; CaH,(C03)3 + heat = CaCOs + Ca + HA Calcium Sulphate. — Sulphuric acid and soluble sulphates give with calcium salts white precipi- tates, of sulphate that form more or less quickly, according to the sulphate used and the strength of th'^ solution. Very dilute solutions are not precipitated. Calcium sulphate ig soluble in a large quantity of water, and much more soluble in acids. It is also soluble in strong solution of ammonium sul- phate. Calcium Oxalate. — Solutions of calcium salts containing no free mineral acid are precipitated by ammonium oxalate, yielding a white precipitate MAGNESIUM. 39 of oxalate. In very dilute solutions the precipi- tate only forms slowly. Presence of free am- monia and heat help the precipitation. Calcium oxalate is easily soluble in hydrochloric and nitric acids, and almost insoluble in acetic or oxalic. Of all tests for calcium it is the most delicate. Soluble pJiospJiates give the same reactions with calcium salts as with those of barium and stron- tium. Potaasiwm dichromate and the normal chro- mates do not precipitate calcium salts. Flame. — Calcium salts give a yellowish-red color when brought into the Bunsen flame. MAGNESIUM (Mg). At. wt. 24.3 ; sp. gr. 1.74. Magnesium is a silvery -white metal, which when ignited in the air burns with an intense white flame, forming the oxide MgO. It does not decompose water at ordinary tem- peratures. It is found vddely distributed in na- ture as carbonate and silicate. Hydrochloric and sulphuric acids dissolve it very readily with evolu- tion of hydrogen. The oxide and hydroxide are very insoluble in water, but soluble in acids. 40 MAGNESIUM. Many magnesium salts are soluble in water, and with the exception of the sulphate are easily decomposed on gentle ignition, most of them even by evaporation of their solutions. Magnesium Hydroxide. — The hydroxides of potassium, sodium, barium, strontium, and calcium give, with solutions of magnesium salts, a white precipitate of hydroxide. If ammonium salts are present, the reagent must be added in excess and the solution boiled to insure complete precipita- tion: MgCla + 2K0H = Mg(OH), + 2KC1. Afnifnonium Hydrate. — The action of ammonia on magnesium salts varies, depending on the pres- ence or absence of ammonium salts. Magnesium salts readily unite with those of ammonium to form double salts, and on these compounds am- monia has no action. Of course the addition of ammonia to an acid solution containing magnesium has the same effect, as the first action is a combi- nation of the ammonia and free acid forming an ammonium salt, which then unites with the mag- nesium to form the double salt. When ammonium salts or free acid are not present, ammonia partially precipitates magnesium MAGNESIUM. 41 solutions, giving a white precipitate of tlie hy- droxide : 2MgCl3 + 2NH40H = Mg(OH), + MgCl, , NH4CI + NH,C1. The reason for the precipitation not being com- plete is shown by the equation to be due to a por- tion of the magnesium uniting with the ammonium chloride formed in the reaction aud producing the double chloride MgCl„NH4Cl. The magnesium hydroxide formed in this and the preceding reaction is soluble in ammonium salts, as shown by the following reaction : Mg(OH), + 3NH4CI = MgCl2,NH4Cl + 2NH4OH. Magnesium Carhonate. — Sodium and potassium carbonates give with magnesium solution a white precipitate of basic carbonate, a portion of the magnesium remaining in solution as acid or bicar- bonate. On heating the solution the latter is de- composed and precipitates as normal carbonate : SMgCla + 5Na,C03 + HgO = 4MgC03,Mg(OH)3 + lONaCl + CO^. The presence of ammonium salts interferes with or prevents this precipitation. 42 MAGNESITTM. Ammonium Garhonate in presence of ammo- nium chloride does not precipitate magnesium salts unless the latter are concentrated, and then only after standing some time. In absence of ammonium salts it gives after some time a precipitate of MgCOgSH^O if the ammonium carbonate be present in slight excess ; of a double carbonate, MgCOg, (NH4),C03,4H80, when the reagent is added in large excess. Magnesium Phosphate. — Hydrogen disodium phosphate gives with neutral solutions of mag- nesium salts a white precipitate of phosphate, MgHP04 : MgCla + Na,HPO, = MgHPO, + 2NaCl. The test is made much more delicate by adding ammonium chloride and hydrate to the solution before the addition of the phosphate. Under these conditions a white crystalline precipitate of magnesium ammonium phosphate is thrown down, even from very dilute solutions. The presence of free ammonia in excess helps the precipitation. Shaking or stirring the mixture hastens the pre- cipitation : MgCla + NH4CI + NH4OH + Na2HP04 = MgNH,P04 + 2NaCl + NH^Cl + H,0. MAGNESIUM. 43 The precipitate of MgHP04 and that of Mg]SrH4P04 are both easily soluble in acids, even in acetic. Am.monium Oxalate in dilute solutions of mag- nesium salts causes no precipitate except on stand- ing. In concentrated solutions a precipitate forms more quickly. Presence of ammonium chloride and hydrate interferes with or prevents the pre- cipitation of magnesium as oxalate. SulplmHc Acid and Soluble Sulphates produce no precipitate with magnesium salts. Chromates do not precipitate magnesium solu- tions. Magnesium salts do not give any color to the Bunsen flame. SEPARATIONS. On reviewing the reactions of the metals of the second group we find that three of them, viz., barium, strontium, and calcium, act in much the same way with many of the reagents, and conse- quently the methods of separation are correspond- ingly limited. The metals Just mentioned ai'e readily separated from magnesium by making use of the action of ammonium carbonate in presence 44 SEPAKATIONS. of ammonium chloride, which precipitates them as carbonates, leaving the magnesium in solution. The carbonates may then be separated from each other by taking advantage of the differences in their reactions with chromates and by the differ- ence in solubility of strontium and calcium sul- phates. The accompanying scheme of separation is based on these reactions. Another method, and one which has the advan- tage of giving some idea as to the quantity of each metal present, is to dissolve the three carbonates in acetic acid and precipitate the barium as in the method given in the scheme ; the strontium and calcium are then reprecipitated from the filtrate by adding ammonia and ammonium carbouate. The carbonates are dissolved in nitric acid and evaporated to dryness, and then heated for about fifteen minutes at a temperature of 150° to 180° C. The dry residue should be at once finely pul- verized and then rubbed up with about 10 c.c. of a mixture of equal parts of absolute alcohol and ether. The solution is filtered after a few minutes and the residue washed several times with small quantities of the alcohol and ether mixture. The strontium nitrate remains insoluble while the calcium nitrate goes into the solution. The SEPARATIONS. 45 residue may be dissolved in a little water and the presence of strontium confirmed by testing with calcium sulphate. The solution containing the calcium nitrate is tested by the addition of a few drops of dilute sulphuric acid, when the formation of a precipitate in any considerable quantity proves the presence of calcium. Should the pre- cipitate be only trifling, it might possibly be due to a trace of strontium, and it should be tested in the flame or by some other confirmatory test. The Sulphates of Barium, St7'ontium, and Cal- cium, are so insoluble in both water and acids (this is especially the case with barium and strontium sulphates) that some special method has to be adopted in order to get them into solution. This is best done by fusing the dry sulphates with four or five parts of sodium potassium car- bonate in a platinum crucible. The insoluble sulphates are thus converted into carbonates, and the sulphuric acid combines with the sodium and potassium as sulphates soluble in water : BaSO, + I^a^C03 = BaCOs + Na^SO^. By treating the fused mass with boiling water until thoroughly disintegrated the sodium and po- tassium sulphates, as well as the excess of flux, are 46 SEPARATIONS. dissolved, and on filtering and washing the residue (the washing should be continued until the wash- ings give no precipitate with barium chloride) the carbonates of barium, strontium, and calcium are obtained free from sulphate, and can be at once dissolved in acetic acid and separated in the usual way. SEPARATIONS. 47 r SCHEME I.— SEPARATION OF 1° AND 2° GROUPS. Filtrate SB. Add a small quantity of NH4CI, and then NH^OH until alkaline, then (NHjjjCOa in sUght excess; warm and Alter. Wash the ppt. Note 1. Ppt. 35. BaCOa + SrCOj + CaCOj. Dissolve in the least possible quan- tity of hot, dil. HCaH.Oj, (36), and test a small portion of the S(j1. for Ba, by addition of CaSOj. An im- mediate ppt. shows Ba. If no ppt. forms at once heat gently and allow to stand 10 min. The appearance of a ppt. on standing will indicate Sr. If Ba has been found by the above test, add to the main part of the sol., KgOr04* in slight excess; allow to stand a few minutes, filter and wash. (If Ba is not found as above, omit this treatment, and proceed to test for Sr and Ca, as below. Part 1° and 8». (Notes.) FiLTKATE 35. MgCl.j, NH/cI -I- KCl + NaCl 4- traces of Ba and Ca 4- NH4 Salts. Divide into 2 parts. Part l".— Add a few drops of (NH4)2S04, and of (NH4).jCj04; a white ppt. shows ti-aces of BaSO* -1- CaCaO^ (43). Filter. To the filtrate add NH4OH and NajHPOj. A white, crystalline ppt. shows MgNH4P04 (44), and proves Mg. (Note 3.) Part 2°. — Evaporate to dryness, and ignite, to expel all NH4 salts. Dissolve residue (45) in a little water, and filter. Evaporate the filtrate to very small bulk, and make the flame- test for Na. Then add a few drops of HjPtClj to the solution, and stir with a glass rod. A yellow, crystal- line ppt. shows (46) KjPtCla, and proves K. Addition of alcohol in- creases the delicacy of this test. (Note 4.) Ppt. 37. BaCrOi (yellow). Filtrate 37. Sr(CjHjOa)j -f Ca(CaH30a)j. Add NH4OH to alkaline reaction, then (NH4)2C03 in slight excess, heat, filter, and wash. Ppt. 38. SrCOa -t- CaCOa. Dissolve in hot dil. HCjHaOj. Divide solution into two parts. FiLT. 38 NH( salts. Eeject. Part. 1°. Add a solution of CaS04, warm and allow to stand some time. A white ppt. = SrS04. Moisten ppt. with HCl (cone.) and test in the flame. Crimson flame proves Sr. Part 2°. Make alkaline with NH4OH, then add a concentrated solution of (NH4)2S04 boil for- some minutes and filter. (Note 5.) Ppt. 41. SrS04H-CaS04. (?) Filtrate 41. Add a few drops of HCaHaOj & (NH.laCaO,. A white ppt.=CaCa04— proves Ca. (Note 6.) * Prepared from KaCraOy by adding NH4OH until the color changes to yellow, but not sufficient to make the solution alkaline. NOTES TO SCHEME. l^OTE l.-'^The amount of ammonium chloride added must be sufficient to form a double chloride with the magnesium present, so that the latter will not be precipitated by the ammonia or am- monium carbonate. A large excess of the chlo- ride is to be avoided, as the carbonates of barium and calcium are perceptibly soluble in it. When the solution under examination is a filtrate from the separation of the 3° and 4° groups, it is not necessary to add any ammonium chloride at this point because it has already been added before precipitating those groups. After the addition of the ammonia and ammonium carbonate the mix- ture should be heated gently (do not boil) for some time, and until the precipitate, if any, settles in the crystalline form. If the solution is a filtrate from the 3° and 4° groups it is sometimes bi'own or almost black, due generally to an imperfect separation of nickel in 48 NOTES TO SCHEME. 49 the preceding groups. . In sucli cases the solution should be boiled until no more ammonia is given off, then allowed to cool, and finally made slightly acid with acetic acid, and filtered. The precipi- tate is tested for nickel, and the filtrate, after be- ing made alkaline with ammonia, for the 1° and 2° groups. Ammonia is always tested for in the original solution. The test is made, as already described under the reactions for ammonia, by heating with an excess of sodium or potassium hydrate. The formation of a precipitate on the addition of the alkali does not interfere with the test, but care must be taken to add sufficient of the hydrate to make the solution distinctly alkaline. Note 2. — The precipitate of the carbonates should be dissolved in a little hot acetic acid and the solution obtained evaporated nearly to dryness to drive off any excess of acid. It is then diluted with a small quantity of water and tested as given in the scheme. The presence of a little free acetic acid is necessary to prevent any strontium being thrown down as chromate, but an excess is injurious, as barium chromate is somewhat soluble in it. Note 3. — If barium and calcium have not been 50 NOTES TO SCHEME. found in Ppt. 35, the ammonium sulphate and oxalate should be added to separate portions of the filtrate, and if either or both give a precipitate they should be filtered on separate filters and the precipitates, after moistening with a drop of con- centrated hydrochloric acid, tested on platinum wire in the Bunsen flame for baiium and calcium. The ammonium sulphate, of course, precipitates the barium, and the oxalate the calcium and per- haps barium. Note 4. — If magnesium has been found in part 1°, it is well to add some ammonium oxalate to part 2° before evaporation, as it helps to decom- pose the magnesium salt. The ignition must be continued until all fumes have ceased. Should no precipitate form on addition of the hydrochloroplatinic acid, the solution must be evaporated nearly to dryness and alcohol added. Potassium, if present, will then surely show. In using the flame test for sodium the student must learn to distinguish between a mere trace and any appreciable quantity. Note 5. — This treatment with ammonium sul- phate does not eifect a complete separation of the strontium and calcium. Some of the calcium gen- erally remains with the strontium, and a little of NOTES TO SCHEME. 51 the strontium may dissolve and go into the solu- tion with the calcium. Note 6. — The solution is made slightly acid with acetic acid so as to prevent the precipitation of any traces of strontium that may have dissolved in the ammonium sulphate. If the solution is very dilute, the precipitate of calcium oxalate will only form slowly. THIRD GROUP. ALXJMrNIUM, CHROMIUM, TITAKItrM. The oxides and hydroxides of this group are very insoluble in water. The precipitated hy- droxides (except metatitanic acid) are easily solu- ble in acids ; but the oxides dissolve with more or less difficulty, and after ignition are almost in- soluble. The third-group metals do not form sulphides in the wet way, and for this reason are not pre- cipitated as sulphide by either hydrosulphuric acid (HgS) or alkaline sulphides. Ammonium sulphide precipitates them from neutral solutions as hydroxide. This reaction distinguishes them from the metals of the fourth, fifth, and sixth groups. ALUMINIUM (Al). At. wt. 27.1 ; sp. gr. 2.58. Aluminium is found in enormous quantities in combination with silica, as clay, feldspar, etc. It also occurs as the oxide and fluoride. It is a 53 ALUMINIUM. 53 moderately hard tin-white metal, very malleable and ductile, and may be tigbly polished. It fuses at 700° C. It is not oxidized by the air, and is unacted on by hydrosulphuric acid or ammonium sulphide. Nitric acid has scarcely any effect on it, but it is soluble in hydrochloric and sulphuric acids, and in sodium and potassium hydrates : 2 Al + 6HC1 = Allele + sSi ; 2A1 + 2K0H = K^AlgO, + SHj. Aluminium forms the oxide ALOg and a cor- responding series of salts. The oxide dissolves with difficulty in dilute acids, but more easily when the acid is concentrated and hot. When fused with sodium or potassium disulphate it is converted into the sulphate, which is easily solu- ble in water : AlA + 3(KaS04, SO3) = Al2(S04)3 + 3K2SO4. When fused with sodium or potassium hydrate or carbonate (if sodium or potassium hydrate is used the fusion must be made in a silver or nickel crucible) it forms au aluminate soluble in water or acids : AlgOa + 2K0H - K2AI3O, + HjjO. 54 ALtTMiNiirir. The aluminium salts are almost all colorless. The haloid salts and the normal nitrate, sulphate, and acetate are all soluble in water. On ignition the soluble oxygen salts are decomposed, leaving a residue of oxide (AI2O3). Many aluminium salts combine with those of the metals to form double salts, among the most important of which are the alums, such as K,S0,,A1,(S04)3,24H30. Most of the aluminium compounds insoluble in water (with the exception of certain minerals) are dissolved by hydrochloric acid. Those that are insoluble in this acid may be got into a soluble form by fusion with alkaline hydrate, carbonate, or disulphate, as already ex- plained. Aluminium Hydroxide {Al^^OH)^). — ^If am- monium sulphide be added to a solution of an aluminium salt a white precipitate of the hydrox- ide is thrown down, which is insoluble in an ex- cess of the precipitant : Al2(SO,)3 + 3(NH,),S + 6H,0 = M(0H)6 + 3(NH4),S04 + 3h5. Potassium and Sodium Hydroxides give with aluminium salts a precipitate of the hydroxide ALUMlNltJM. 55 contaiaing alkali and basic salt (a). The pre- cipitate is easily soluble in excess of the reagent (3), and is not reprecipitated on boiling (distinc- tion from chromium) ; but on the addition of am- monium chloride, especially on heating, it repre- cipitates as the hydroxide (o). («) Al2(S04)3 + 6K0H = M(0H)6 + 3K2SO4 ; (h) A1,(0H)6 + 2K0H = K^Aip^ + ^H^O ; (c) K2AI2O, + 2NH4CI + 4H3O X = M(0H)6 + 2KC1 + 2NH,0H. Ammonium Hyd/roxide also gives a precipitate of the hydroxide which is insoluble in a slight ex- cess of the precipitant, but perceptibly soluble in a large excess. From this solution in excess of the ammonia it is reprecipitated on boiling, as this drives out the excess of ammonia. The Alkali Garhonates give a precipitate of hydroxide and basic carbonate. The precipitate is slightly soluble in the carbonates of potassium and sodium, but much less soluble in ammonium carbonate. Ba/rivm Ca/rhonate completely precipitates solu- 56 OHEOMItrM. tions of aluminium salts, even in the cold, giving a precipitate consisting of hydroxide and basic salt. The barium carbonate should be used in a very finely divided state. The presence of non-volatile organic acids, such as tartaric, citric, etc., and of sugar or similar or- ganic matter, interferes w^ith or completely pre- vents the precipitation of aluminixam as hydroxide. AluminiiLm Phosphate. — With neutral solu- tions of aluminium salts or those acid only with acetic acid sodium phosphate gives a white floc- culent precipitate of phosphate. The precipitate is soluble in the mineral acids and in sodium or potassium hydroxide, but diffi- cultly so in ammonia. Ammonium chloride reprecipitates it from its solution in the alkalies : Ala(S0,)3 + 2Na,HP04 + SN-aC^HaO, = M(PO,), + 3Na,S04 + 2HC2H30,. Akbyninium Basic Acetate. — If a moderate ex- cess of sodium or ammonium acetate be added to a dilute neutral or slightly acid solution of an aluminium salt and the liquid boiled, a white pre- cipitate of basic aluminium acetate will be thrown down : CHROMIUM. 57 A1,(S0,)3 + eNaCAO, + 2H,0 = Al2(0H)s(C,H A)4 + 3Na,S04 + 2HC2H3O0. The precipitate is easily soluble in hydrochloric and sulphuric acids. CHROMIUM (Cr). At. wt. 52.14; sp. gr. 6.8-7.3. Chromium in the metallic state is rarely met with. It is sometimes alloyed with steel in the proportion of 0.5% to 0.75^, for the purpose of hardening the former. It is readily soluble in chlorine or hydrochloric acid, and in sulphuric acid on warming. Chromium occurs chiefly as oxide in com- bination with oxide of iron as chrome-ironstone FeO,CrA- Three oxides of chromium are known — CrgOg, CrOgjCrOg. The monoxide has not been obtained in the anhydrous state, although its hydroxide is known. Chromic oxide (CrgOg) is green, the hydroxide a bluish green. The common salts of chromium are those of this oxide. They are all of a violet or green color, and many of them are soluble in water. Most of those not soluble in water are soluble in hydrochloric acid. 58 CHEOMItrM. The oxygen salts of volatile acids are decom- posed on ignition, leaving a residue of chromic oxide (CrgOg) which is almost insoluble in acids. Ohromium, Hydroxide. — Potassium and sodium hydroxides precipitate chromic salts, giving bluish- green precipitates of the hydroxide containing al- kali (a). The precipitate is easily soluble in an excess of the reagent, yielding an emerald-green solution (J)). From this solution on long boiling the hydroxide is reprecipitated (6). Addition of ammonium chloride effects the same result (d). {a) Cr^CSOOa + 6K0H = Cr,(0H)6 + SK^SO^ ; (h) Cr,(0H)6 + 2K0H = KA2O4 + 4H2O ; (c) KjCr^O^ + 4H3O = Cr3(OH)6 + 2K0H ; {d) K^Cr.O^ + 2NH4CI + 4H30 = Cr3(OH)6 + 2KC1 -f 2NH,0H. Ammonia and Ammonium Sulphide also give precipitates of chromic hydroxide which are slightly soluble in excess of ammonia. On boil- ing the precipitation is complete. The Alkaline Carhonates give precipitates of basic carbonate, which dissolve with difficulty in an excess of the reagent. CHROMItTM. 59 Barium CarJxytiate precipitates chromium as a basic carbonate. The precipitation takes place in the cold, but requires long digestion to make it complete. Sodiwn Phosphate in presence of sodium ace- tate precipitates neutral or slightly acid solutions as green chromic phosphate, Cr3(P04)2. The presence of tartaric, citric, and oxalic acids, sugar, etc., interferes more or less with all of the precipitations of chromium that have been men- tioned. Oxidation to Chromic Acid. — The most charac- teristic reactions for chromium depend on its oxi-> dation to chromate. The reaction is indicated by a change in color from green or violet to yellow or red. If any one of the precipitates already mentioned or any compound of chromium be fused (best in a platinum crucible) with four or five parts of a mixture of sodium carbonate and potas- sium chlorate in equal proportions, the chromium will be oxidized to chromate, NagCr04. By dissolving the fused mass, which, if chro- mium is present is of a yellow color, in water, acidifying with acetic acid, and adding lead ace- tate, a yellow precipitate of lead chromate will be obtained: 60 TITANITTM. Cr2(OH)6 + 2Na,C03 + KCIO3 = 2Na,Cr04 + KCl + 200^ + 3H^ ; Na,Cr04 + Pb(CaH302)2 = PbCrOa- 2NaC,H30,. A similar oxidation is effected by treating tlie solution of chromium hydroxide in potassium hy- droxide (KjCrgO^) with bromine. In performing the operation the solution must be kept alkaline and the mixture warmed for some time and finally heated to boiling. The color changes from bluish green to yellow. If the solution be now acidified with acetic acid and a drop or two of lead acetate added, a yeUow precipitate of lead chromate will be formed : KA2O, + 8K0H + SBrg = 2K2Cr04 + 6KBr + 4H30. Other oxidizing agents, as lead or sodium diox- ides, or potassium permanganate, in alkaline solu- tion, effect the same oxidation. In a Bead of Borax or Sodmm Metaphosphate in either the oxidizing or reducing flame, chro- mium compounds give an emeiald-green color. TITANIUM (Ti). At. wt. 48.1; sp. gr. 3.58. Titanium is a comparatively rare metal, and is not found in the metallic state. It occurs as the TITANIUM. 61 dioxide, TiOg, in several minerals, and is frequently met with in magnetic iron ores, and for this rea- son it is necessary to know something about its properties. The metal is easily soluble in acids, even in acetic. It forms three oxides, TiO, Ti^Og, and TiOg. Salts of the last two oxides are known. TiO^ , besides acting as a basic oxide, also acts as a feeble acid. Titanium dioxide after ignition is insoluble in acids, except in concentrated sulphuric and hydro- fluoric. Fusion of the dioxide with potassium disulphate (the fusion should be continued for a considerable time, half an hour or more, at a low heat) yields a mass that is completely soluble in a large quantity of cold water (hot water must not be used, as metatitanic acid is precipitated from hot solutions). Titanic oxide may also be readily brought into solution by fusing with potassium hydrogen fluo- ride and dissolving the fusion in dilute hydro- chloric acid. Solutions of titanium dioxide in hydrochloric or sulphuric acids are precipitated by the alkalies, the alkaline carbonates, ammonium sulphide, and barium carbonate as titanic acid, HgTiOg, a white 62 TITANIUM. flocculent precipitate insoluble in excess of the reagents. When thrown down from cold solution the precipitate is soluble in hydrochloric and sul- phuric acids : Ti(S0,)2 + 4K0H =2^X103 + 2K3S04 + HA Sodium Phosphate precipitates solutions of titanium salts as a white precipitate of phosphate (TiOHP04). The precipitate forms even in solu- tions containing much hydrochloric acid. Potassium Ferrocyanide added to acid solutions gives a bright reddish-yellow precipitate. Sodium Thiosulphate precipitates nearly neutral solutions on boiling as titanic acid, HgTiOs (sepa- ration from iron). Precipitation on Boiling. — Solutions of tita- nium in hydrochloric or sulphuric acids, not con- taining much free acid, on dilution with consider- able water and boiling give a white precipitate of metatitanic acid insoluble in dilute acids. The solution best suited for this reaction is one in sul- phuric acid containing only a slight excess of the acid. In order to insure complete precipitation the boiling must be continued for five or six hours, with renewal from time to time of the water lost by evaporation. TITANIUM. 63 Hydrogen Peroxide gives with titanium solu- tions containing no fluoride an orange-yellow color. When iron is present a solution of the sulphate should be used. Metallic Zinc or Tin added to acid solutions of titanic salts give a violet or blue color to the solu- tion. The color often does not appear for some time, and the test is not very delicate. SEPAKATION. The separation of the metals of the third group will be discussed later on in connection with those of the fourth group. FOURTH GROUP. IRON, ZINC, MANGANESE, COBALT, NICKEL. The metals of this group are distinguished from those of the fifth and sixth by the fact that they are not precipitated by hydrosulphuric acid (HgS) from solutions acid with the mineral acids, and from those of the third group by their precipita- tion as sulphide and not hydroxide by ammonium sulphide. The fact that they are precipitated as sulphides also distinguishes them from the metals of the first and second groups. IRON (Fe). At. wt. 56; sp. gr. 7.8-8.2. Iron is found in small quantities in the metallic state in some volcanic rocks and in meteorites. Compounds of iron are very widely distributed, and occur in enormous quantities particularly as oxide and sulphide. Pure iron is a lustrous, grayish- white metal. It 64 IRON. 65 is hard, malleable, and ductile, fusing with diffi- culty, and is attracted by the magnet. In moist air or oxygen it is slowly oxidized to the hydrated sesquioxide (rust). When ignited in air or oxy- gen it forms the black magnetic oxide Fe304, or a mixture of this with ferric oxide, FcgOa. Iron unites directly with the halogens, and with sul- phur, carbon, boron, silicon, phosphorus, and arsenic. It also unites with many metals to form alloys. Iron is readily soluble in dilute acids. In hydro- chloric and dilute sulphuric acids it dissolves with evolution of hydrogen (a) in hot concentrated sulphuric acid, sulphurous anhydride is liberated (h) and ferric sulphate formed. With nitric acid the reaction varies according to temperature and strength of acid used. With cold dilute acid the products are ferrous nitrate and ammonium nitrate, and with hot moderately dilute acid fer- ric nitrate is formed (c, d). {a) re + H,S04 = FeS04 + H2; {h) 2Fe + 6H3SO4 = Fe,(S04)3 + SSO^ + eH^O ; (c) 4Fe+10HNO3 = 4Fe(N03), + NH4NO3 + 3H3O ; (d) 2Fe + 8HNO3 = Fe,(N03)6 + 2N0 + 4H,0. 66 IKON. Three oxides of iron are known : ferrous oxide, FeO (also called protoxide) ; ferroso-ferric oxide, Fe304 (magnetic or black oxide) ; and ferric oxide, FejOg (peroxide, sesquioxide, or red oxide). The ferrous and ferric oxides form two distinct series of salts having very different properties. FBRKOTJS COMPOUNDS. Ferrous oxide is black, the hydroxide white. The latter absorbs oxygen very rapidly, and quickly changes in color, assuming first a greenish shade and finally becoming reddish brown. The oxide and hydroxide are both easily soluble in hydrochloric, nitric, and sulphuric acids. As al- ready explained, ferrous salts are also formed by dissolving metallic iron in hydrochloric or dilute sulphuric or nitric acids (see equations a and c). Ferrous salts in the anhydrous state are white, in the hydrated green. The solutions unless concen- trated have scarcely any color. These solutions absorb oxygen and change gradually to feme salt with precipitation of basic compounds. Oxidizing agents, as chlorine, bromine, or nitric acid, quickly oxidize ferrous solutions to ferric : IKON. ' 67 6FeCl, + 8HNO3 = 2Fe,Cl6 + Fe3(N03)6 + 2N0 +4H,0. Ferrous salts of volatile oxygen acids are decom- posed on ignition, leaving a residue of ferric oxide. Ferrous Sulpliide (^Fe8^. — Hydrosulphuric acid in alkaline, and ammonium sulphide in neu- tral or alkaline, solutions of ferrous salts give black precipitates of ferrous sulpMde, which are easily soluble in hydrochloric, nitric, and sul- phuric acids. The presence of ammonium chlo- ride aids the precipitation : FeSO^ + (NH4)2S = FeS + (NH4),S04 ; FeS + 2HC1 = FeCla + ITs. Ferrous Hydroxide {Fe^OH)^. — The hydrox- ides of potassium, sodium, and ammonium precipi- tate ferrous solutions as hydroxide, Fe(OH)g. The presence of ammonium salts interferes with the precipitation by potassium or sodium hydrox- ide, and prevents the precipitation by ammonia. These alkaline solutions however quickly oxidize, 68 IRON. and the iron then precipitates as ferroso-ferric or ferric hydroxide. ferrous Carbonate {FeCO^. — The soluble car- bonates give a white precipitate of ferrous carbon- ate, wh;ch rapidly oxidizes to the reddish-brown ferric hydroxide. Barium carbonate does not precipitate ferrous salts in the cold, except solutions of the sulphate. Ferrous Ferrocyanide. — Potassium ferrocyan- ide in neutral or slightly acid solutions produces a bluish-white precipitate of potassium ferrous ferrocyanide, KgFe(FeC6N6), which on exposure to the air or by treatment with oxidizing agents changes to Prussian blue. Ferrous Ferricyanide. — Potassiiim ferricyanide in neutral or slightly acid solutions produces a deep blue precipitate of ferrous ferricyanide, Fe3(FeC6N6)a (Turnbull's blue) : 3FeS0, + Ke(FeCeNe), = Fe3(Fe CeNe),+ 3K,S0,. Of all the tests for ferrous salts this is the most delicate and characteristic, the reaction taking place even in very dilute solutions. IRON. 69 FERRIC COMPOUNDS. Ferric oxide is red or brownish red, and dis- solves with more or less difficulty in acids even on heating. It is most easily gotten into solution by digesting with hot strong hydrochloric acid, or by fusion with potassium disulphate. Ferric hydroxide is reddish brown, and is easily soluble in acids. Ferric oxide and its hydrates are found very widely distributed, and in enormous quan- tities, and constitute the most important ores of iron. Ferroso-ferric oxide, Fe304, is soluble in hot hydrochloric acid, the solution containing fermus and ferric salt. The normal ferric salts in the anhydrous state are white, the basic ones are red or yellow. Neu- tral ferric solutions are brownish yellow, the acid solutions yellow, and the color is visible even when much diluted. Salts of volatile oxygen acids are decomposed on ignition. Ferrous Sulphide. — Hydrosulphuric acid added to acid solutions of ferric salts produces a white precipitate of sulphur, and the ferric compound is reduced to feiTous salt : Fe^Cl, + H,S = SFeCls + 2HC1 + S. 70 IKON. Ammonium sulphide in neutral or alkaline solution causes the same reduction (a), and the ferrous salt is then precipitated as sulphide (^). (a) Fe^Cle + (NH4),S = 2FeCla + 2NH4CI + S ;• (b) 2FeCl, + 2(NH,),S = 2FeS + 4NH4CI. The properties of this sulphide have already been described. Ferric Hydroxide (^Fe^^ OH)^. — Potassium, sodium, and ammonium hydroxides give reddish- brown pi'ecipitates of the hydroxide, which are Insoluble in an excess of the precipitant and in ammonium salts, but easily soluble in acids : Fe^Cle + 6K0H = Fe3(OH)6 + 6KC1. The Soluble Carbonates also give precipitates consisting principally of hydroxide, insoluble in excess of sodium and potassium carbonate. Barium Carbonate even in cold solution pre- cipitates the iron completely as hydroxide and basic salt. Non-volatile organic acids or sugar interfere with or prevent the precipitation of ferric salts as hydroxide. lEOK. 71 Ferric Acetate. — If to a neuti'al or slightly acid soliition of a ferric salt an excess of sodium or ammonium acetate be added, the color changes to a reddish brown. On diluting and boiling the solution a reddish ■ brown precipitate of basic acetate is thrown down. The precipitate is easily soluble in hydrochloric acid : Fe^Cls + eNaCgHaOa + 2R^0 = Fe,(OH),(CaH30a)4 + 6NaCl + 2RG.^,0,. Ferric Pliosphate. — Soluble phosphates in neu- tral or slightly acid solutions to which a slight excess of a soluble acetate has been added give yellowish-white precipitates of ferric phosphate, which are insoluble in acetic but soluble in the stronger acids : Fe^Cle + 2Na2HP04 + BNaC^HgOs = Y%^{VO,\ + 6NaCl + 2HC2H3O2. Ferric Ferrocyanide. — Potassium ferrocyanide produces even in very dilute solutions a deep blue precipitate of ferric ferrocyanide, Fe4(FeC|iN6)3 (Prussian blue). 'J'S IKON. The test is best made in a solution slightly acid with hydrochloric acid : 2Fe,Cl6 + SK^FeCeNe) = Fe,(FeC6N6)3 + 12KC1. Potassium Ferricyanide produces no precipitate with a ferric salt, but changes the color of the solution to a dark brown (distinction from ferrous salts). Ferric Thiocyanate. — Potassium thiocyanate (sulphocyanide) gives a deep red color of ferric thiocyanate, Fe2(CNS)6, even in excessively dilute solutions. The reagent should be added in mod- erate excess, and the solution be acid with hy- drochloric acid. Certain acids, as acetic, tartaric, oxalic, phosphoric, hydrofluoric, arsenic, etc., inter- fere with the reaction, and when any one of them is present the hydrochloric acid must be added in large excess. A similar color is produced by nitric acid if the acid be strong or the solution hot. By adding a little alcohol and warming, the color, if due to nitric acid, will disappear. This test is the most characteristic and delicate for ferric salts. Ferrous salts give no color with the thiocyanate. ZINC. 73 ZINC (Zn). At. wt. 65.4 ; sp. gr. 6.9-7.2. Zinc is a bluish- white lustrous metal. At or- dinary temperatures commercial zinc is brittle, but becomes malleable at temperatures between 100° C. and 150° C. ; at 210 C. it becomes so brittle that it may be finely pulverized. It melts at 419° C, and boils at about 950° C, and may be easily dis- tilled at a bright red heat, the vapors given off burning in the air with a bright white light and the formation of zinc oxide, ZnO. When exposed to, the air zinc becomes gradually coated with a thin film of basic carbonate. Zinc is found in considerable quantity, the principal ores being the carbonate, silicate, and sulphide. Pure zinc is scarcely at all soluble in acids, but the presence of any impurity even in small amount causes it to dissolve easily. With dilute hydrochloric and sulphuric acids it evolves hydro- gen (a). With veiy dilute nitric acid ammo- nium nitrate or nitrous oxide is formed (h and c), and in stronger acid nitric oxide (d). It is also soluble in sodium or potassium hy- droxides (e). 74 ziisrc. (a) Zn + H,S04 = ZnSO^ + H^ ; (b) 4Zn+10HNO3 = 4Zii(N03)2 + NH4NO3 + 31-1,0 ; (c) 4Zn + IOHNO3 = 4Zn(N03)3 + So + 5H,0 ; (d) 3Zn + 8HNO3 = 3Zn(N03),+ 2N0 + 4H,0 ; (e) Zn + 2K0H = K^ZnO, +1^^ Zinc forms one oxide, ZnO, white, and the hy- droxide, Zn(0H)2. Both are easily soluble in acids. Zinc salts are colorless unless the acid be col- ored. Some are soluble in water, the rest in acids. Zinc jSulpJiide (Z71S). — From neutral solutions of zinc salts, even those of the stronger acids, hy- drosulphuric acid precipitates the greater part of the zinc as a white sulphide. From solutions of the acetate the precipitation is complete, even in presence of free acetic acid. For this reason if a moderate excess of sodium acetate be added to a neutral or slightly acid solution of a zinc salt in one of the stronger acids and H^S be added, the Avhole of the zinc will be precipitated as sulphide : ZnCl, + H3S + 2NaC,H303 = ZnS + 2NaCl -f 2HC2H3O2. ziKO. 75 This reaction is particularly useful as a separa- tion from manganese, which is not precipitated under these conditions. Ammonmm Sulphide in neutral of alkaline solutions produces the same white precipitate of sulphide. Presence of ammonium chloride aids the precipitation. Zinc Sulphide is easily soluble in hydrochloric, nitric, and sulphuric acids, and insoluble in sodium, ammonium, or potassium hydroxides. Zinc Hydroxide {Zn{OH')^. — Sodium and po- tassium hydroxides produce white precipitates of zinc hydroxide, which are easily soluble in an ex- cess of the precipitant. These alkaline solutions if concentrated are not precipitated on boiling, but when dilute the greater part of the zinc pre- cipitates : ZnClo + 2K0H = Zn(0H)2 + 2KC1 ; ZnOHa + 2K0H = K^ZnOs + 2H3O. Ammonium. Hydroxide added to neutral or slightly acid solutions also precipitates the hy- droxide, which is easily soluble in excess of the ammonia. Presence of ammonium salts or the 76 ZINC. presence of much free acid in the original solution (forming ammonium salt on addition of ammonia) interferes with or prevents the precipitation. If the afkaline solution be boiled, the zinc par- tiall}' precipitates if the solution be a concentrated one; if dilute the precipitation is complete. ZiiiG Carbonate. — The soluble carbonates give white precipitates of basic carbonate, insoluble in excess of sodium or potassium carbonate, and sol- uble iu excess of ammonium carbonate. Presence of ammonium salts prevents these precipitations either wholly or partially according to the quan- tity present. Bariwm Carbonate does not precipitate zinc salts from a cold solution, with the exception of the sulphate. Zinc Cyanide. — Potassium cyanide throws down white zinc cyanide, which is easily soluble in excess of the reagent : ZnCl, + 2KCy = ZnCy^ + 2KC1 ; ZnCyg + 2KCy = ZnCy2,2KCy. From this solution, if the excess of potassium cyanide is not great, the zinc may be precipitated MANGANESE. 77 as sulphide by the addition of sodium or potas- sium sulphide. Hydrosulphuric acid and ammonium sulphide act slowly and imperfectly. Color Test. — If to a solution of a zinc salt a few drops of cobalt nitrate (sufficient only to give a faint pink color) be added, the solution then made alkaline with sodium carbonate, boibd, filtered, and washed, and the precipitate ignited on a plat- inum capsule, a green mass will be obtained. The color shows more distinctly if the ignited precipitate be pulverized. MANGANESE (Mn). At. wt. 55 ; sp. gr. 6.85-8. Manganese is a whitish gray, lustrous metal, very hard and brittle. It is very difficultly fusi- ble, its melting-point being about 1900° C. It oxidizes very easily in the air, and decom- poses water with evolution of hydrogen. Manganese is easily soluble in all acids. With copper it forms a valuable alloy resem- bling bronze, and alloyed with iron it forms spie- geleisen, or ferromanganese, used in the manufac- ture of steel. Manganese forms a number of different oxides — MnO, Mn304, Mn^Os, MnOg, and probably 78 MANGANESE. MnOg and M118O7. The oxides MnOj and MngO- are the anhydrides of manganic, H2M11O4, and permanganic acids, HgMnaOg. Manganic acid has not been separated, but many of its salts are known. The oxides MnO, MD3O4, and MngOg are basic oxides, and dissolve in hydrochloric or sul- phuric acid, forming manganous or manganic salts. If the acid is hot the solutions contain only manganous salt. MnO;^ acts both as a basic and an acid oxide: MnO + 2HC1 = MnCl^ + H^O ; Mn.Oa + 6HC1 (hot) = 2MnCl, + Cl^ + SH^O ; MnO, + 4HC1 (hot) = MnC], + Cl^ + 2HaO ; 2Mn03 + 2H2SO4 (cone, hot) = 2MnS04 + 2H30 + 02. The stable salts of manganese are those cor- responding to manganous oxide, MnO. They are colorless or pink. Some are soluble in water, the rest in acids. Manganous Sulphide (MnS). — Ammonium sulphide in neutral or alkaline solutions precipi- tates manganese as a flesh-colored sulphide. The precipitate is readily soluble in acids, even in MANGANESE. 79 acetic. On exposure to the air it turns brown, ow- ing to oxidation. Presence of ammonium chloride helps the precipitation : MnClg + (NH4)2S = MnS + 2NH,C1. Mangmious Hydroxide (^Mn(^OII)^. — Potas- sium and sodium hydroxides give precipitates of manganous hydroxide which are white at first but quickly change to brown and finally to a dark-brown, almost black color, due to the pre- cipitate oxidizing to a higher oxide, probably MnCla + 2K0H = lln(OH), + 2KC1. AmmoniuTYi Hydroxide added to neutral solu- tions causes a partial precipitation of the man- ganese as hydroxide, a portion remaining in solution as a double salt of manganese and am- monium. Solutions containing much free acid or ammonium salt are not precipitated by ammonia owing to the formation of this double salt : 2MnCl,+2NH40H=Mn(OH)2+MnC], , 2NH4CI. Manganous hydroxide is soluble in ammonium 80 MAJsIGANESE. cliloride, but after oxidation the precipitate is in- soluble : Mn(OH)2+4NH,Cl=MnCl3,2NH,Cl+2NH,OH. The alkaline solution on standing absoi'bs 0x3'^- gen, and a brown precipitate is thrown down. Manganous hydroxide in presence of free alkali is oxidized to MnOg, HgO by bromine or hydrogen dioxide : Mn(OH)2 + 2KOH + Biv, "^ = MnOa , HgO + 2KBr + H^O. Bromine also precipitates solutions of man- ganous salts containing sodium acetate, as dioxide. Heating assists the reaction : MnCla + 4NaC,H30, + Br, + 3H,0 = MnO, , HgO + 2NaCl + 2NaBr + 4HC3H3O,. Manganous Carhonate. — The soluble carbonates precipitate manganous salts as the normal carbon- ate or as a mixture of carbonate and hydroxide. The precipitate is insoluble in an excess of the precipitant. Barium Carbonate produces no precipitate ex- cept in solution of the sulphate. MANGANESE. 81 Sodium, Manganate {Nac^MnO^. — If any com- pound of manganese be fused in a platinum cap- sule with a mixture of sodium potassium carbon- ate and an oxidizing agent such as sodium nitrate, the fused mass on cooling will have a characteris- tic bluish-green color, due to the formation of sodium manganate : 3Mn(OH)3 + Na^COa + 4NaN03 = 3Na3Mn04 + CO3 + 4N0 + SHaO. The same result may be obtained by fusing the substance in a bead of sodium potassium carbonate held in a loop of platinum wire, in the oxidizing flame of the blowpipe or Bunsen burner ; Mn(OH)3+Na2C03+03=Na,Mn04+COa+HaO. Permanganic Acid {Hc^Mn^O^. — If a few drops of a solution of a manganous salt, which must be free from chlorine or chloride, aie boiled with a large excess of concentrated nitric acid, and then some lead dioxide* add-jd and the solution *The lead dioxide should be tested in tlie same way to see if it contains any manganese, as tlie latter is sometimes found in it as an impurity. 82 COBALT. diluted, the liquid will have a purple color from the permanganic acid formed in the reaction : 2MnS0, + 2PbO, + 2HNO3 = H^MngOs + 2PbS04 + 2^. Bead Test. — If any compound of manganese be fused in a borax or sodium metaphosphate bead in the oxidizing flame of the blowpipe or Bunsen burner, an amethyst-colored bead is obtained. In the reducing flame the color is destroyed. COBALT (Co). At. wt. 58.9 ; sp. gr. 8.96. Cobalt is steel-gray, lustrous, hard and mallea- ble, and ductile at a red heat. It is slightly mag- netic even at a full red heat. In the air at oidi- nary temperatures it does not oxidize. In hot dilute hydrochloric or sulphuric acid it'dissolves wdth evolution of hydrogen, but the best solvent is dilute nitric acid. The solutions contain co- baltous salts. Cobalt forms three distinct oxides, CoO, C03O4, and CogOg. With some few exceptions the stably salts of cobalt are those corresponding to cobalt- ous oxide, CoO. COBALT. 83 The higher oxides are soluble in hydrochloric acid with evolution of chlorine and formation of cobaltous salts. The cobaltous salts in solution or in the crystal- line state are generally red or pink, and the anhy- drous salts are mostly blue. A solution of co- baltous chloride on evaporating nearly to dryness changes to blue, but on addition of water the pink color is restored. CohaltouH Sul/pliide (CoS). — Hydrosulphuric acid added to neutral cobaltous salts of the min- eral acids causes a partial precipitation of the metal as sulphide. Presence of free strong acid prevents the precipitation. Neutral solution of the acetate is almost wholly precipitated, but pres- ence of free acetic acid prevents the precipitation, or nearly so. If, however, an excess of an alkali acetate be present, then HgS precipitates the cobalt completely, especially on heating, even in presence of free acetic acid. Ammonium 8%iJ/phide added to neutral or alka- line solutions completely precipitates the metal as sulphide. Presence of ammonium chloride aids the reaction. Cobalt sulphide is insoluble in alkalies, and nearly so in acetic acid. In hydrochloric acid it 84 COBALT. dissolves with difficulty, and when precipitated from hot solutions it is almost insoluble in that acid. Nitric acid and aqua regia dissolve it on heat- ing: CoCla + (^Bi)S = CoS + 2NH4CI. Cobaltous Hydroxide (Joi^OH).^. — Potassium and sodium hydroxides give at first a blue precip- itate of basic salt, which on boiling is converted into the pale red hydroxide. The precipitate is insoluble in an excess of the dilute alkali, but sol- uble in acids and in ammonia and ammonium car- bonate : C0CI2 + 2K0H = Co(OHX + 2KC1. Ammonia in neutral solutions gives the same precipitate as sodium or potassium hydroxide, but it is easily soluble in an excess of ammonia, form- ing a brownish-yellow liquid. In acid solutions or those containing ammonium salts, ammonia gives no precipitate. The precipitation of cobalt as hydroxide is interfered with or prevented by the presence of non-volatile organic acids or sugar. COBALT. 85 Oohaltous Cao-honate. — Sodium and potassium carbonates give precipitates of basic carbonate wMcb are insoluble in excess of the reagent. Barium Carbonate in tbe cold precipitates only a solution of the sulphate. Cobaltous Cyanide. — Potassium cyanide pro- duces a brownish-white precipitate of cobaltous cyanide which is very easily soluble in an excess of the precipitant, with formation of a double cy- anide : CoCL + 2KCN = Co(CN), + 2KC1 ; Co (CN% + 4KCN = Co(CN)o,4KCN. From this solution the cobaltous cyanide is re- precipitated on the addition of hydrochloric acid : Co(CN)5„ 4KCN + 4HC1 = Co(CN), + 4KC1 + 4HCy. If to a solution of the double cyanide contain- ing some excess of KCN a few drops of hydro- chloric acid be added (so as to libei-ate some hy- drocyanic acid) and the solution boiled, or better if it be mixed with some potassium hydroxide and chlorine or bromine added to the cold solu- 86 COBALT. tion, the cobaltous double cyanide is changed to po- tassium cobalticyanide, KbCo^(CN)i3, from whicli solution acids do not reprecipitate the cobaltous cyanide. The reaction with hydrochloric acid yielding free hydrocyanic is perhaps represented by the following equation : 2(Co(CN)3, 4KCN) + 2HCN =K6Co2(CN)i3 4- 2KCN + 5, and with bromine or chlorine by this : 2(Co(CN%, 4KCN) + Br, = K6Co,(CN)i2+2KBr. Nickel forms no corresponding compound. Potassium Cobaltic Nitrite. — If a moderate ex- cess of potassium nitrite (it is well to add it in the solid form so as to avoid diluting the solution) be added to a cobaltous solution that has been made acid with acetic acid, the cobalt will be pre- cipitated as a yellow precipitate of potassium co- baltic nitrite : 2Co(N03), + lOKNO, + 4HNOo = KeCo.CNO,),, + 4KNO3 + 2N0 -f 2Hp. NICKEL. 87 The more concentrated the solution the more quickly the reaction takes place. With dilute solutions the precipitation is very slow, often re- quiring as much as 24 hours or more before it is complete. Even with strong solutions the test should be allowed to stand for at least 12 hours. Nickel salts are not precipitated by potassium nitrite when cobalt and nickel only are present in the solution. In the Borax Bead cobalt compounds give a beautiful deep blue color in either the oxidizing or reducing flame of the Bunsen burner or blow- pipe. The test is exceedingly delicate and very characteristic. NICKEL (Ni). At. wt. 58.7; sp. gr. 8.97-9.26. Nickel is a lustrous white metal with a yellow- ish-gray tinge. It is hard, malleable, ductile and very tenacious, and slightly magnetic. It fuses with difficulty, and does not oxidize in the air at ordin- ary temperatures. On heating itdissolves slowly in hydrochloric and dilute sulphuric acids, but is easily soluble in dilute nitric acid. In concentrated nitric acid it is passive. In these solutions the nickel is in the form of nickelous salt. Nickel is found in the metallic state in some 88 NICKEL. meteorites. In its principal ores it exists in com- bination with arsenic or sulphur. There are thi-ee oxides — NiO, Ni304, and NigOs. The ordinary salts of nickel are those correspond- ing to nickelous oxide, NiO. This oxide and its hydroxide, ^(OH)^, are green; both are readily soluble in hydrochloric, nitric, and sulphuric acids. The higher oxides dissolve in hydrochloric acid to nickelous salt, with evolution of chlorine. In the anhydrous condition nickelous salts are usually of a yellow color, and in the hydrated state they are green. Nickelous Sul/pMde {NiS'). — Hydrosulphuric acid acts in the same manner with nickelous salts as with those of cobalt. Anvmoniuin Stdphide in neutral or alkaline solutions produces a black precipitate of nickelous sulphide, which is slightly soluble in ammonium sulphide, particularly in pi'esence of much free ammonia. This alkaline solution of the nickel sulphide is brown. By neutralizing the solution with acetic acid and warming the nickel sulphide is precipi- tated. The presence of ammonium chloride or acetate aids the precipitation of nickel as sul- phide. Nickel sulphide is very slightly soluble NICKEL. 89 in acetic acid, and dissolves with difficulty in di- lute hydrocliloiic. It is easily soluble in nitric acid and in aqua regia on heating : ■ NiCla + (NH,),S = NiS + 2NH4CI. Nickeloiis Hydroxide {Ni(^OH)^. — Potassium and sodium hydroxides give a light green precipi- tate of nickelous hydroxide, v?hich is insoluble in an excess of the precipitant : NiCla + 2K0H = M(0H)2 + 2KC1. Ammovda in neutral solution causes a slight turbidity, but on the further addition of the re- agent this dissolves to a blue liquid containing a combination of nickelous salt and ammonia. In presence of ammonium salt or free acid ammonia gives no precipitate. Nickelous Oarhonate. — Sodium and potassium carbonates give precipitates of green basic carbon- ate of nickel, insoluble in excess of the reagent. Barium Carhonate in the cold precipitates only solutions of the sulphate of nickel. Nickelous Cyanide {NiCy^. — Potassium cya- nide produces a green precipitate of nickelous 90 NICKEL. cyanide, which is readily soluble in an excess of the reagent. From this solution the nickelous cyanide is reprecipitated by acidifying the liquid with hydrochloric or sulphuric acids : NiCla + 2KCN = Ni(CN% + 2KC1 ; Ni(CN)2 + 2KCN = Ni(CN)„ 2KCN' ; Ni(CN)„ 2KCN + 2HC1 = Ni(CN)3 + 2KC1 + 2HCN. If the solution of the double cyanide be made strongly alkaline with potassium hydroxide and chlorine or bromine be added (care must be taken to keep the solution alkaline), the nickel will be precipitated as the black hydroxide, ]Sria(0H)6 : 2(Ni(CN)8, 2KCN) + 6K0H + Br, = Nio(OH)6 + 2KBr + 8KCN. This reaction enables us to separate nickel from cobalt ; the latter it will be remembered gives no precipitate under the same conditions, but changes to cobalticyanide, which remains in solution. NICKEL. 91 Bead Test. — When fused in the borax bead in the oxidizing flame nickel compounds give a red- dish-brown color. With sodium metaphosphate the bead is yellow or reddish yellow in both flames. SEPARATIONS — THIRD AND FOURTH GROUPS. If we examine the reactions of tte metals of tte third and fourth groups with ammonia in presence of ammonium chloride, we find that those of the third group and ferric salts of the fourth are pre- cipitated by this reagent as hydroxides, insoluble in an excess of the reagent. This would seem to offer an excellent method for separating these metals from the remaining ones of the fourth group, but it has been found in practice that these hydroxides if present in considerable quantity always carry down some of the other metals as well, so much so in fact that if the latter were present in small quantity only they might be completely precipitated. By dissolving the precipitate in hydrochloric acid and reprecipitating with ammonia it is possi- ble to remove them to some extent, but not en- tirely. When very exact results are not required, 93 SEPARATIONS— THIRD AND FOURTH GROUPS. 93 or when the hydroxides are present in small quan- tity, this separation may be used to advantage. The reaction with barium carbonate in a solu- tion of the chlorides affords one of the most accu- I'ate methods for separating the third-group and ferric salts from the fourth-group metals. The test should be made on a cold solution that is neutral or only slightly acid, and a moderate ex- cess of the finely divided carbonate added (freshly precipitated barium carbonate acts most rapidly). The mixture must be allowed to stand for some time and until no more carbon dioxide is evolved. When this plan of separation is adopted the cobalt and nickel are generally separated first as sulphides. The method of procedure is given in Scheme Number III. The hydroxides of the third group and iron may be separated in various ways. If we make use of their different solubilities in excess of so- dium or potassium hydroxide, we can separate the aluminium and chromium hydroxides, which are soluble, from that of iron (and titanium), which is insoluble. The chromium and aluminium can then be sep- arated from each other by oxidizing the chromium 94 SEPAKATIONS — THIRD AND FOURTH GROUPS. in the alkaline solution with bromine or hydrogen dioxide and testing the resulting solution as given in Scheme II. Another and perhaps better plan is to fuse the hydroxides of aluminium, chrcmium, and iron with three to four parts of a mixture of sodium potassium carbonate and potassium chlo- rate in about equal proportions. On treating the fusion with water the iron is left as oxide (Fe^Og) and perhaps some of the alumina. The solu- tion contains the chromium as chi'omate and the greater part of the alumina as Na^Al204. One portion of the solution is acidified with acetic acid and lead acetate added, when the formation of a yellow precipitate shows the pi-esence of chro- mium. The remainder of the solution is acidified with hydrochloric acid and then made slightly alkaline with ammonia or ammonium cai'bonate. A white precipitate indicates alumina. Cobalt and nickel are generally removed at the beginning of an analysis by treating the precipi- tate produced by ammonium sulphide with dilute hydrochloric acid, when, owing to the insolubility of their sulphides in that acid, they are left as a black residue. Their separation from each other may be ac- complished in various ways. The one used in SEPAKATIONS — THIRD AND FOURTH GROUPS. 95 the scheme with bromine in a solution of the cyanides is accurate and rapid. The precipitation of the cobalt with potassium nitrite, and of the nickel in the filtrate with potassium hydroxide, gives excellent results, but the test requires a long time before the precipitation is complete. Zinc and manganese may be separated as given in the scheme, or, taking advantage of the differ- ence in solubility of their sulphides in acetic acid, we can effect the same result by treating a solu- tion of their acetates made acid with acetic acid, with HgS ; the zinc is thrown down as sulphide, and the manganese remains in the solution, from which, after filtering out the zinc sulphide, it may be precipitated as hydroxide with potassium hy- droxide. Solution of salts of the stronger acids may be separated in the same way by adding a moderate excess of sodium acetate and acetic acid before precipitating with HaS. The separation of the metals of the third and fourth groups is somewhat complicated in the presence of phos- phates, silicates, organic matter, etc., and therefore requires some modification of the ordinary meth- ods of analysis. In Scheme IV a method is given in which the presence of these acids is provided for. 96 NOTES TO SCHEME II. Note 1. — If the solution to be analyzed is a filtrate from the separation of the fifth and sixth groups, the Jl.S contained in it should be expelled by boiling. After this has been done, ferrous salts if present (determined by testing a few drops of the solution with potassium ferricya- nide) are oxidized to ferric by the addition of a little concentrated nitiic acid, and boiling. A test for non- volatile organic matter must also be made at this point, unless it has already been tested for in the preliminary examination in the glass tube (see complete analysis). The test is made by evaporating a small quantity of the liquid, which should be made acid (best with sul- phuric acid) if not so already, to dryness, and igniting gently. A black carbonaceous residue proves the presence of organic matter. In pres- ence of much nitric acid or nitrate the blackening may be only transient, due to the carbon being quickly oxidized by the oxygen of the nitric acid or nitrate. In cases of this kind more or less deflagration will be noticed. Should organic matter be found, the main portion of the solution must be evaporated to dryness (best after addition of some concentrated nitric acid) and ignited at a temperature just suificient to thoroughly carbon- SCHEME II.— THE SEPARATION OF Al, Cu, Ti, Co, Ni, Zn, Mn, Pk BY KOII AND BROMINE. (Note 1.) To the solution of the metals add NII4CI in niodeiale (iimutity, then NILOII until sligiilly alkaline, and finally (NHjijS until solution smells distinctly of that reagent; stir well and heat gently for some time, lilter and wash the precipitate. FlLTUATE 1°. Contains the alkalies and alkaline earths. Prf.CII'IT.M'E 1°. Treat with cold dilnle IK'l (sp. gr. 1.0'2) in modrralc exc(!ss; liller and Wiish the residue thoroughly. (Note 0.) Rksidue. (CoS and NiS.) Test in bora.x bead. Blue bead ^ Co. (Note 3.) Brown bead = Ni. If Co is found to be present, treat the residue of sulpliides as follows : Place ill a small evaporating dish, add aqua-regia and l)oil, evaiioraling nearly to dryness; dilute willi a liltle water; filter, if Jieccssary; nearly neutralize the solution with KOII; add KCy in moderate ex- cess — make strongly alkaline witli KOH, and add Br, and warm, taking care to keep the solulion .strongly alkaline. (Note 4.) Filter, and wa.sh precipitate thoroughly with boil inir water. FlI.TUATE. Cobalt may be deter- mined in this solulion, if necessary, by acidifying with HNO3, and then add- ing (Hg).j(NO,!.j ; filter, ami ignite the precipitate, and test residue in borax bead ; blue = Co. Ppi (Ni..(01I).. Black. Test in borax bead — brown bead = Ni. PlLTltATH 2°. Evaporate nearly to dryness to expel H.jS. and the excess of IICl;* dilute with a little water, filter, if necessar}', from any separated sulphur; add strong solulion of KOH in considerable excess, and then a few dro|)s ol bromine (it is better to use i>ure bromine ratlicr than bromine water); heat gently for a few minutes, lilter, aiul wash precipitate with boiling water. (Note 5.) Per. ;!. Test a portion in Na.jC'Oj; head in oxidizing flame, green bead = Wn. (Note (i.) Dissolve some of the preci])- itate in HCI; test a portion of this witli NIhCyS for iron (Note 7); red color = Fe''. Test remaiiuler of solution with Zn or Sn — violet color = Ti. (Note 8.) PlI/ritATE 0°. Divide solution into three parls. Pass II.jS .gas into the solulion (not to satur.-i- lion) — a while precipi- lale, insoluble in KOII = ZnS. (Note 9.) Acidify with IICl, then make very faintly alka- line with' NH,OH — a while lloccnieut precipi- tate = A1,(0H)„. (Note 10.) Acidify with acetic acid and add a few drops of plumbic acetate— j'el- low precipitate =:PbCr04 = (!r. (Note 11.) * If the color of the solution at tliis noiiit lioos not indicate the prespnce of chronihnn, tin* treatiniMit with bromine maybe omitted, but i[i this case iron {if pre.sent) ninst be o.^iilized by boiling: the solution witii a little HNOj before the aildition ol the KOti. A solution containing one part of chroininni in ten thousand of water shows a disl inct bluish-green color; consequently a colorless solution 'oulil not contain more tliaii a liuce of that melill. To follow p. 95. NOTES TO SCHEME II. 97 ize tlie organic matter present. No attempt should be made to burn off the carbon, as the high tem- perature necessary would make the oxides of iron, alumina, and chromium very insoluble. After the residue has been completely carbonized, add a small quantity of strong hydrochloric acid and a little water, and boil for some time, and filter. Tie filtrate is now ready to be tested with ammonia, ammonium sulphide, etc. These preliminary operations completed, a small portion of the solution is taken and ammonium chloride and a slight excess of ammonia added. Note carefully whether the latter causes any pre- cipitate, and if so what color it is, then add ammo- nium sulphide and see if it produces any further precipitate and of what color. (A light-colored precipitate would indicate the absence of metals yielding black sulphides, viz., iron, cobalt, and nickel.) Should ammonia have produced a precipitate, and the original solution have an acid reaction (neutral solutions cannot contain phosphates), take another small portion of the liquid from which the H^S has been expelled, and the iron oxidized, and add a moderate quantity of tartaric acid, then make alkaline with ammonia. If the 98 NOTES TO SCHEME II. latter still causes a precipitate it indicates the presence of phosphates, etc., of the second group, and -the solution must be analyzed according to Scheme IV. The tartaric acid in this test is used to prevent the precipitation of iron, alumina, and chromium as hydroxides or phosphates. If in the test last given ammonia produces no precipitate, the mixture is tp be analyzed by Schemes II or IIL If in the first test ammonia gave no precipitate, metals of the third group and iron must be absent, and no further testing for them is necessary. Should ammonium sulphide also have produced no precipitate, the fourth group cannot be present, and the rest of the solution is tested at once for the second group. Our preliminary tests completed, and the results indicating the absence of phosphates, etc., and the presence of third and fourth-group metals, we will return to the treatment of the main portion of our solution. A moderate quantity of ammonium chloride is added to form a double salt with the manganese (and with magnesium if second-group metals are present), and to aid the precipitation of the fourth- group metals as sulphides. If the original solution NOTES TO SCHEME II. 99 is acid it may not be necessary to add any ammo- nium chloride, as suflBcient may be formed when the solution is neutralized with ammonia. One will have to judge as to whether this is the case or no by noting the quantity of ammonia required to make the solution neutral. After the addition of the ammonium chloride or its formation in the solution, enough ammonia is added to make the solution faintly alkaline. An excess is to be avoided, as it not only dissolves appreciable quantities of aluminium and chro- mium hydroxides, but also helps to carry nickel sulphide into solution. The ammonium sulphide should be added until the solution after shaking smells distinctly of that reagent. Thorough stirring or shaking, and heating the liquid gently for some time, aids the precipitation and causes the precipitate to collect in large par- ticles that are more easily filtered. The precipita- tion is best made in a small flask that can be corked, and the mixture in it thoroughly shaken. The precipitate should be washed with water to which a few drops of ammonium sulphide have been added. Note 2. — It is better to remove the precipitate 100 NOTES TO SCHEME II. from the filter-paper and treat it in a beaker or test-tube with, the acid. It should be well shaken, so that every particle of the precipitate will come in contact with the solvent. In case the amount of the precipitate is small, it may be removed from the. filter by making a hole in the point of the latter and washing the precipitate through with a little water. If the treatment with acid leaves no black resi- due, cobalt and nickel cannot be present. A small quantity of iron sulphide is sometimes left undissolved in the residue. Note 3. — The test for cobalt with the borax bead is so delicate, even in the presence of large quantities of nickel, that no other test for it is generally required, Note 4. — In the treatment with potassium cyanide care must be taken not to use too great an excess ; a little more than enough to dissolve the precipitate first formed by the cyanide is sufficient. It must be remembered that cobalt and nickel cyanides dissolve in potassium cyanide very easily, and if there is any portion of the precipitate that does not go into solution readily it should be filtered out, as it is probably a little ferric hydrox- ide formed from a small amount of ferrous sul- NOTES TO SCHEME II. 101 phide left undissolved by the hydrochloric acid in the first treatment of the sulphides. The solution must be kept decidedly alkaline with KOH, and a moderate quantity of bromine added. Note 5. — After the solution has been evapo- rated to expel the excess of acid, it must be diluted with very little water, so as to keep it concentrated, and the solution of potassium hydroxide should be strong, and added in con- siderable excess in order to insure solution of the zinc. A few drops of pure bromine (not bromine water) are then added if the color indicates chro- mium, and the solution gently heated for a few minutes to make sul-e of oxidizing the latter ; the liquid is then boiled for a short time and filtered. In making this separation care must be taken to keep the mixture strongly alkaline. Note 6. — This precipitate may be tested for manganese by fusing it in a platinum capsule with a mixture of sodium potassium carbonate and sodium nitrate, when the green manganate will be formed. Note 7. — The state of oxidation of the iron is of course not shown by this test. In order to determine it, the original solution must be tested before any oxidizing or reducing agents have been 102 NOTES TO SCHEME II. added, with potassium ferricyanide for ferrous salts and the thiocyanate for ferric. Note 8. — Titanium is so seldom met with except in the analysis of some iron ores and fur- nace products, that its presence is rarely considered in an ordinary analysis. When required to test for titanium it is better to fuse some of the original material with potassium disulphate, dis- solve the fusion in cold water, and test with hydrogen dioxide, or metallic zinc, or any other characteristic test. Note 9. — If this solution be saturated with HjS, aluminium if present would be precipitated as hydroxide. It can, however, be easily distin- guished from zinc sulphide by adding some more potassium hydroxide, in which it is readily soluble, while the zinc sulphide is insoluble. Note 10. — An excess of ammonia must not be added. Heating the solution helps the precipita- tion of the aluminium hydroxide. The precipitate is insoluble in ammonium chloride (distinction from zinc hydroxide). Since alumina and silica are common impurities in sodium or potassium hydroxides, it is always well to take a quantity of the alkali equal to that used in the test for alumina, and treat it with NOTES TO SCHEME II. 103 hydrochloric acid and ammonia in precisely the same way, and if any precipitate is obtained com- pare it with that from the regular test. If much less, it is safe to conclude that the substance under examination contains alumina. Note 11. — Only a drop or two of the lead acetate should be added, so as to avoid the pre- cipitation of lead chloride or bromide, as any con- siderable quantities of these white precipitates might hide small amounts of the yellow chromate. Scheme IV. THE SEPARATION AND DETECTION OF Al, Qr, Fb, Co, Ni, Mn, Zn, Ba, Ca. Sk, and Ma IN THE PRESENCE OF PHOSPHORIC, OXALIC, BORIC, SILICIC, OR HYDRO- FLUORIC ACID. If a solution containing the above-mentioned acids and bases is made alkaline with ammonium hydrate the following is likely to occur : HsPO. ppts. Al Cr Fe(ic) Co Ni Mn Zn Ba Sr Ca Mg HsBOs " " " " " " " " " " " " HsSlOa " " " " " " " " " " " " TTT^ K I* t( <( " '( (< •( <( (( (( it HaCsOi " chiefly " " " If the test with tartaric acid and ammonia (see Note 1, Scheme II) has indicated the presence of phosphate, etc., a few preliminary tests should be made before proceeding further. Several small portions of the solution are taken and tested as follows : I. Make slightly alkaline with ammonia, filter out precipitate and dissolve it in a small quantity of dilute nitric acid, and add an excess of ammo- nium molybdate. A yellow precipitate proves the 104 SCHEME III.— SEPARATIOK OF GROUPS III AND IV, IN THE ABSENCE OF PHOSPHATES, BORATES, OXALATES, FLUORIDES, ETC. (BaCO, METHOD). Filtrate 3. Boil out HjS, add NH.Cl, NH4OH, and (NH,),S in slight excess; warm and filter; wash with water contaiuiug a little (NH))jS. NiS + CoS + FeS + MnS + ZuS + Cr,{OH), Precipitate 22. Al2(OH)o. Dissolve in cold dilute HCl (sp. gr, 1.03) and wash residue thoroughly. Filtrate 23. 1°, 2° groups. Residue 23.* NiS + CoS (black). Test in borax bead in O. flame of the blowpipe. A blue bead shows Co; a violet-brown bead shows Ni. If the bead is blue, dissolve the residue in aqua regia, evaporate to small bulk to expel excess of acid, dilute with water, and nearly neutralize with KOH. (If a ppt. forms, it must be redissolved by a drop of HCl.) Add KCy until ppt. which first forms is dissolved by the excess; add KOH in excess, and then Br. water ; warm gently, and keep strongly alkaline with KOH ; filter. Wash thoroughly with boiling water. Ppt. 24. Ni.,(OH)« (black). Test in borax bead with blowpipe in O. flame. A violet-brown bead proves Ni. Filtrate 24. KeCOjCyi-r. Evaporate to dryness, and confirm Co io O.liame, of the blowpipe. Blue bead proves Co. Filtrate 23.+ MuCl, + FeCU + ZnCU + Or^CA., + A1,C1„ + HCl + H»S. Boil, to expel HjS ; add a little cone. HNO3 , and boil, to oxide FeO to FcaOs ; evaporate to small bulk, to expel excess of acid; cool, nearly neutralize with NajCOa, and pour into a small flask. Dilute the solution; add BaCOs suspended in H2O, in excess, and fill up the flask; then cork loosely and allow to stand some time; filter and wash with hot water. Dissolve in dil and wash thoroughly. Ppt. 26. Fe2(0H)„ + AU(OH)„ + Cr,(OH)o + BaCO,. HCl ; dilute the solution and add NHjOH in slight excess. Filter Filtrate. BaClj. Reject. Ppt. 26. Fe»(OH)o + A1.,(0H). + Cr,(OH),. Dry ppt. and fuse in a platiniim capsule with I! parts of NaKCOs and 1 part of KCIO3 ; digest the fused mass with HjO and boil until disin- tegrated. Filter and wash. Filtrate 26. ZuClj + MnClj + BaCl.,. Heat to boiling; add slight excess of dil. HiSOj ; allow to stand, filter, and wash. Ppt. 33. BaSO^. Reject. Filtrate 33. MiiCU + ZnCIj. Add excess of KOH; digest and filter. * A Method far the Separation of Niand Co Without the UseofKGN. Residue 37. FeaO., -f AI5O3. Dissolve in HCl, and add a drop of NH,CNS. Blood-red color shows Pe-2(CNS)o , and proves Fe. Dissolve the mixed sulphides in aqua regia and boil to expel excess of acid. Take a small part of this solution and add a slight excess of NHjOH; then add a few drops of pure Br and allow to stand for several minutes. Note the color of the solution: a violet tinge shows the proba ble presence of Ni in some quantity. Add to the .solution a considerable excess of KOH and filter, Wash the ppt. of Ni(OH)j and test for Ni in the O. flame of the blowpipe, with borax bead. t Test for FeO and FesOs in the original solution as follows: Acidify a small part of the sol. with HCl; filter, if necessary, and divide into 2 parts. To the first part add a "drop of NH.CNS; a blood-red color shows the presence of FejOs. To the second part add a drop of KsFeCy.; a blue ppt. shows the presence of FeO. Filtkate 37. KsAUO. -f KjCrOi. Acidify with HC2H3O2, aiul heat. Divide solution into 2 parts. Part 1 (30). Add PNCaHjOj)! . a yellow ppt.; shows PliCrOi , and proves (^r. Part 2 (39). Add excess of (NH4)2C03. and boil. A white flocculcnt ppt. shows A12(0H)b , and proves Al. Or, instead of the above, add Na3HP04. A white, flocouleut ppt. shows Ali(P04).j, and proves Al. S2. < c T3 Q.S CJi n a; 5; 5 p- CO CO 9 ^ • o S PI "^ B ? O 3 P ^ > P 0= cc TofoUowp. 103. SCHEME IV. 105 presence of phosphoric acid. (Arsenic acid, which acts like phosphoric, will have been removed with the fifth and sixth group metals as sulphide.) II. Oxalates are seldom met with in an analy- sis. They may be detected by heating the solu- tion (the more concentrated the better) with some manganese dioxide (free from carbonate) and a little dilute sulphuric acid. If present, there will be an effervescence and evolution of carbon diox- ide (CO3), which may be recognized by its action on lime-water (^see Carbonic Acid). III. Silicic acid when present is generally re- moved before . beginning the tests for the metals by evaporating an acid solution to dryness and taking up the residue with acid and water (see Silicic Acid). The same test may be applied here if neces- sary. Borates and fluorides are usually held in solu- tion (or sufficiently so) by the ammonium chloride present, so that they do not interfere to any great extent, and no modification of the method of an- alysis is required. We are now ready to proceed with the separa- tion, supposing the presence of phosphates, etc., to have been shown by the tests just made. 106 SCHEME IV. Boil out HgS, if present. Add a few drops of HNO3 and boil (Note 1). If HAO* or organic matter is present, evaporate to dryness and ignite gently. If these are not present, but silicic acid is, then acidify with HCl and evaporate to dry- ness, but do not ignite (Note 2). Treat the residue with HCl (cone), dilute with H^O, and boil ; it dissolves wholly or leaves a white residue of SiO^. Filter. If neither oxalic nor silicic acids are present the solution after boiling with nitric acid to oxidize the iron is treated according to filtrate 1°: Besidde 1° SiOo. Filtrate 1." Nearly neutralize with NH4OH, then add a moderate excess of NHjCaHjO, and a little HCaHjOa, dilute largely, boil for about 5 min., and filter. Wash the ppt. with boil- ing water. (Note 3.) Ppt. 2°. Fe, Al, and Cr as phosphates and basic acetates. Dry and mix with NajCOj + KClOs, and fuse in platinum capsule. Dis- solve the fusion in boiling water, and filter. (Note 4.) Filtrate 2°. Metals of the 1°, 2", and 4° Groups, ex- cept iron. Analyze in the same way as any ordinary mixture of the first four groups. Residue 8°. FejO,,AI,0,, Feo(PO,)j, and Alj(Vo.\. Solution 3°. NagCrOi, NagAlgOt, etc. (Divide in 3 parts.) 1° acidify with HCjHjOj (if any precipitate forms, it shows AIPO4) : filter, and to the filtrate add Pb(Cj,Hs03)3, a yellow ppt. = PbCr04 ; proves Or. 2° acidify with HCl; then add (NH4),C03, and bofl. A white, floccutent precipitate proves Al. NOTES.— PHOSPHATE SEPARATION. Note 1. — The nitric acid is added to oxidize ferrous salts to ferric. In case , organic matter is present it i^ well to use quite an excess, as it aids in the decomposition of the organic substances. As already stated, ignition, when necessary, should always be conducted at as low a temperature as possible, just sufficient to completely carbonize the organic material present. A high temperature decomposes iron, aluminium, and chromium salts, and makes their oxides, which are left by the ignition, very difficultly soluble in acids. Note 2. — The evaporation for silica must be continued until the dry mass on stirring smells no longer of acid. The temperature used should not exceed 110°- 120° C. Note 3. — The neutralization with ammonia should be continued until the precipitate which forms on its addition dissolves very slowly. A moderate excess of ammonium acetate should 107 108 NOTES. — PHOSPHATE 8EPAKATI0W. then be added. If iron or aluminium salts are present, a precipitate of their phosphates may- then be thrown down. If sufficient iron is pres- ent to precipitate all the phosphoric acid, the solution, after the addition of the ammonium ace- tate, will have a red color, due to ferric acetate. If the solution, however, remains colorless, ferric chloride must be added, drop at a time, until the red color appears or until a few drops of the liquid give a red precipitate with ammonia. Any large excess of iron is to be avoided, but enough to give a reddish tinge to the liquid is necessary to insure complete precipitation of the phosphoric acid. The ferric acetate is precipitated on boiling the dilute solution as a basic acetate, and any ferric phosphate that it may have been held in so- lution is thrown down at the same time. A small amount of free acetic acid is necessary in the reaction to prevent any zinc or manganese pre- cipitating. The solution should not be boiled more than four or five minutes, and it must be quite dilute. The filtration should be carried on rapidly, the liquid being kept hot, and the precipitate washed with boiling water. NOTES. — PHOSPHATE SEPARATIOlSr. 109 Note 4. — Although, aluminium phosphate is not completely decomposed by fusion with sodium carbonate, still sufficient of it goes into solution to give a satisfactory test for aluminium. Another method that could be used would be to treat the precipitate with potassium hydroxide. The aluminium and chromium basic acetates and phosphates are soluble in this reagent, and go into solution. The chromium is then oxidized in the alkaline solution with bromine, and the chromium determined in one portion by acidify- ing with acetic acid, filtering if necessary, and adding lead acetate. The alumina is determined in the other part by acidifying with hydrochloric acid and then making alkaline with ammonia. METALS OF THE FIFTH GEOUR The common metals of this group are : Silver (Ag), Lead (Pb), Mercury (Hg), Bismuth (Bi), Copper (Cu), and Cadmium (Cd). The chief characteristics of the group are the precipitation of the metals by H^S from acid solutions, and the insolubility of these precipi- tates in alkaline sulphides (note exceptions to latter under mercury and copper). HgS or alkaline sulphides also precipitate neu- tral or alkaline solutions in the same way. The group is subdivided into two divisions, the first comprising the metals yielding chlorides insoluble in dilute hydrochloric or nitric acids, viz., silver, lead, and mercurous salts. Lead is often found in both divisions, owing to the fact that its chloride is not very insoluble in water, and even more soluble in solutions of certain salts. HO SILVER. Ill We will now proceed to a study of the indi- vidual metals of the group, and when that has been completed the separation of the group from the other groups, and the detection of its different metals, will be explained. SILVER (Ag). At. wt. io8; sp. gr. 10.53. A white, lustrous metal, very malleable, fuses at 954° C. It is not oxidized by water or air at any temperature; is easily attacked by CI, Br, or I, and is quickly tarnished by H^S or sul- phides. The proper solvent is nitric acid, in which it is readily soluble. 'It is also soluble in hot strong sulphuric acid. 6Ag + 8HNO3 = 6 AgNOa + 2N0 + 4H2O ; 2Ag+2llSOi = AgSO, +SO2 +2H2O. Silver Ohloride (AgCi). — Hydrochloiic acid or soluble chlorides when added to solutions of sil- ver salts give a white curdy precipitate of silver chloride (a) which blackens on exposure to light. The precipitate is insoluble in dilute hydro- chloric and nitric acids, but is somewhat soluble in the strong acids. It is easily soluble in am- 112 SILVER. monia (5), from whicli solution it reprecipitates on acidifying with nitric acid (c). Also soluble in potassium cyanide as double cyanide (d), and in sodium thiosulphate (hyposulphite) (JSaS^O^) (e). Soluble to a considerable extent in concen- trated solutions of chlorides of the alkalies. The double cyanide solution is largely used in electro- plating, and the reaction with thiosulphate is made constant use of in photography, also in the extraction of silver from its ores. (a) AgNOs + HCl = AgCl + HNO3 ; (b) 2AgCl + 3NH,0H = 2AgCl,3NH3 + SH^O ; (c) 2AgCl,3NH3+ 3HN03= 2AgCl + 3NH4NO3 ; {d) AgCl + 2KCy = AgCy,KCy + KCl; {e) 2AgCl + 3Na2SA = Ag3S303,2Na2S203 + 2NaCl. Silver chloride may be reduced to metal by adding to it a piece of zinc and a little dilute acid, as sulphuric : 2AgCl + Zn = Ag + ZnCl2. SILVER. 113 Fusion of the chloride with sodium carbonate effects the same result. Silver Bromide (^AgBr). — Soluble bromides precipitate silver salts, giving a yellowish-white precipitate of bromide, insoluble in dilute nitric acid and less easily soluble in ammonia than the chloride. Soluble in potassium cyanide and thio- sulphate, the same as the chloride. Decomposed by concentrated hydrochloric acid, with formation of AgCl. Silmer Iodide {AgT). — Soluble iodides give a yellow precipitate of silver iodide, soluble in ex- cess of the reagent,' forming a double iodide, and in potassium cyanide and thiosulphate. It is almost, insoluble in ammonia. Insoluble in dilute acids. Silver Cyanide {AgCy). — HCy or soluble cy- anides give white precipitates of silver cyanide, easily soluble in excess of reagent. Insoluble in dilute nitric acid, and decomposed by the boiling concentrated acid. Soluble in ammonia and thio- sulphate. Decomposed on ignition into metallic silver, cyanogen gas, and paracyanogen. Silver Sxilplvide {AgoS). — HgS or soluble sul- phides precipitate silver salts, giving a black pre- cipitate of silver sulphide, insoluble in cold dilute 114 SILVER. acids, ill alkaline sulphides, and in potassium cya- nide. Easily soluble in boiling dilute nitric acid, with separation of sulphur: 2 AgNOs + H^S = Ag,S + 2HNO3 ; SAgsS + 8HNO3 = 6 AgNOs + 3S+ 2N0 + dH^O. Silver Oxide (Ag^O). — Sodium or potassium hydrates precipitate solutions of silver salts, yield- ing brown silver oxide, insoluble in excess of the reagent : 2AgN03 + 2K0H = Ag30+2KN03+HA Ammonium hydrate gives the same precipitate in neutral solutions, easily soluble in excess of ammonia (Ag20,2NH3). Silver oxide is decomposed by heat into metallic silver and oxygen. Silver can easily be obtained in the metallic state from its solutions by the addition of one of the more electropositive metals, as Cu, Zn, Hg, Fe etc., also by many reducing agents, as FeSOi j SOg , etc. LEAD. 115 LEAD (Pb). At.wt. 207; sp. gr. 114. Lead is a soft, malleable metal, of a gray color, fusing at 335° C. It is insoluble in dilute sul- phuric and hydrochloric acids, but dissolves easily in dilute nitric acid : 3Pb + 8HNO3 = 3Pb(N03)2+2NO+4H30. Water containing air acts on lead forming a hy- drate, which soon changes to a basic carbonate. This action is increased by the presence of nitrog- enous organic matter, also by nitrates and nitrites. Lead forms several oxides (PbO, Pb304 , PbO^, etc.). On ignition the higher oxides are all re- duced to PbO, which is the only oxide forming stable salts. Lead oxide (PbO) is a yellow substance soluble in nitric and acetic acids, also in sodium and potas- sium hydrates. All the stable salts of lead are those of this oxide. Triplumbic tetroxide, or red lead (Pb304), is, as the name implies, of a red or scarlet color. It is largely used as a pigment. With nitric acid it gives lead nitrate, Pb(!N'03)2 , soluble, and leaves a brown residue of lead diox- 116 LEAD. ide (PbOs). Hydrochloric acid slowly dissolves it, with evolution of chlorine. Reducing agents, as oxalic or tartaric acids, sugar, or alcohol, cause it to dissolve completely in nitric acid. Lead dioxide, often called puce or peroxide, is of a brown color. It is insoluble in nitric acid ex- cept when reducing agents are present. With hydrochloric acid it yields lead chloride and chlor- ine. Lead Chloride {PhCl^. — In solutions of lead salts, not too dilute, hydrochloric acid or soluble chlorides give a white precipitate of lead chloride, only slightly soluble in cold water, but soluble in hot. In dilute hydrochloric acid it is more insol- uble than in water. In strong hydrochloric or nitric acid it is much more soluble than in water : Pb(N03)a + 2HC1 = PbCls + 2HNO3. If ammonia be added to lead dhloride it changes it to a basic chloride, PbCl2,3PbO,4H20, which is very insoluble in water. Lead Bromide {PlBr^. — Soluble bromides precipitate lead salts as bromide, white, solubility in water about the same as the chloride. LEAD. 117 Lead Iodide {Phl^. — Soluble iodides precipi- tate lead iodide, yellow, more insoluble in water than tbe chloride or bromide. Soluble in excess of the reagent, forming double iodides (as KI,Pbl2 or 4KI,PbI,). Lead Sulphide^ PbS. — H^S or soluble sul- phides give with lead salts a black precipitate of PbS (a). In presence of much hydrochloric acid the precipitate is sometimes red, owing to the formation of a compound of lead chloride and sulphide (as PbClg,2PbS) ; on diluting the solu- tion with water and adding an excess of HgS, the red precipitate is concerted into the black sul- phide. Lead sulphide is easily soluble in hot dilute nitric acid, yielding lead nitrate (h). With hot concentrated acid it gives lead sulphate (c). (a) Pb(N03)2 -f H,S = PbS + 2HNO3 ; Q>) SPbS+SHNOs =3Pb(N03)3+2NO+3S+4H20 (c) 3PbS+8HN03=3PbS04 + 8NO + 4H20. Lead Sulphate {PhSO^. — Sulphuric acid and soluble sulphates precipitate lead salts as sulphate 118 lEAD. (a), wMte, insoluble in water, and in dilute acids, particularly dilute sulphuric. Concentrated bydrochloric dissolves it as lead chloride. It is soluble in boiling ammonium acetate and in the fixed alkalies. Also soluble in warm (68° C.) sodium thiosulphate (J)) (distinc- tion from barium sulphate). Presence of much free nitric acid interferes with the precipitation of lead as sulphate. («) Pb(N03)2+H2S04 = PbS04 + 2HNO3 ; (i)PbS04 + 3Na2SA '~^ =2Na,SA,PbSA + Na^S04. Lead Hydroxide (^Pb(^OH)z). — Sodium or po- tassium hydrates precipitate lead solutions as hy- droxide, white, soluble in excess of the reagent : Pb(N03)2 + 2K0H = Pb(OH), + 2KNO3 ; Pb(0H)2 + 2K0H = KgPbO, + 2H,0. Ammonia gives precipitates of basic salts, in- soluble in excess of ammonia. With solutions of the acetate not too concen- trated, ammonia gives no precipitate owing to LEAD. 119 the formation of a soluble tribasic acetate (PbaO^CC^HaO,).). Lead Carbonate. — Soluble carbonates give white precipitates of basic lead carbonates, varia- ble in composition. Lead Ohromate {PhCrO^. — Solutions of lead salts are precipitated by soluble chromates, giving a precipitate of lead chromate, yellow, soluble in sodium or potassium hydrates, insoluble in dilute nitric and acetic acids : Pb(N03 )a + KjCrO^ = PbCrO^ + 2KNO3. In addition to the compounds spoken of, the phosphate, oxalate, cyanide (there is no double cyanide), and sulphite are insoluble in water. Lead salts are easily reduced to the metallic state by the more electropositive metals, such as zinc, iron, or magnesium. If any lead compound is heated on charcoal in the reducing flame of the blowpipe, the lead is reduced to the metallic state {a), and a portion volatilizes and forms a coating of lead oxide on the coal. This coating is dark yellow when hot, light yellow on cooling. The addition of sodium carbonate to the sub- 120 MERCURY. stance before the fusion on cliarcoal assists the reaction very materially (5). (a) PbS04 + C = Pb + C02 + S02. (b) 2PbS04+2NasC03+5C=2Pb+2Na2S+7C02. If, instead of making the fusion on charcoal, the mixture of NagCOg with the lead copapound be fused in a porcelain crucible, the reaction is different, the lead being obtained as oxide : PbS04 + Na^COa = PbO + Na^S04 + CO^. MERCURY (Hg). At. wt 200 ; sp. gr. 13.59. A white lustrous metal; liquid at temperatures between -40° C. and 360° C. At the latter temperature it boils and volatil izes. It is slightly volatile at ordinaiy tempera- tures. When in a very finely divided state the metal appears as a gray or black pulverulent pow- der, without metallic lustre. It is not oxidized by the air. The best solvent is nitric acid. It dissolves easily in dilute acid on heating, or in the cold if lower oxides of nitrogen are present. With cold dilute acid the product is mercurous MEECtJET. 121 nitrate («), with hot acid in excess mercuric nitrate ((5»). It is also soluble in strong sulphuric acid on heating (c) and in chlorine. (a) 6Hg + 8HN03= B}Ig,(NO,%+ 2N0 + 4H,0; (b) 3Hg + 8HN03= 3Hg(N03),+ 2N0 + 4H2O; (c) Hg + 2H2S04= HgSO,+ SO2 + 2H2O. Hydrochloric acid does not dissolve it, A very characteristic property of mercury is its power of forming amalgams with many other metals. This is made use of on a very large scale for the extraction of gold and silver from their ores. To a few drops of mercury in a watch-glass add a little finely divided lead or zinc, and notice that the lead or zinc dissolves in the mercury, and that the latter becomes pasty. By igniting the amal- gam gently the mercury volatilizes and the metal with which it was combined is left as a residue. Mercury forms two well-defined oxides, mercur- ous oxide (HggO), black, and mercuric oxide (HgO), red, and two corresponding series of salts. 122 MEEOUET. Mercurous Salts. With the exception of the nitrate (Hg2(N03)2), the mercurous salts are insoluble or difficultly so in water. The oxygen salts are volatile on ignition, under- going decomposition at the same time. The chlo- ride and bromide volatilize unaltered. Hg^Oli,. — Hydrochloric acid or soluble chlorides give with mercurous salts a white precipitate of HggClg (calomel). This precipitate is insoluble in cold hydrochloric acid, but on boiling is slowly changed to mercuric chloride (HgClj) and metallic mercury (Hg). Chlorine or nitric- acid quickly change it to mercuric salt, soluble in water. KOH or NaOH convert it into mercurous oxide (Hg^O), black. NH4OH also turns it black, but the product is a mixture of a mercuric ammonium compound and metallic mercury, insoluble in excess of ammonia (distinction from silver) : Hg.(N03),+ 2HC1 = Hg£l,+2HN03 ; Hg2Cl2+2NH40H white = NHaHgCl + Hg +NH4CI+ 2H2O. black MEECOEY. 123 BSr or soluble hromides precipitate mercurous solutions, giving a precipitate of mercurous bro- mide (HggBra), yellowish white; insoluble in di- lute HNO3. Soluble iodides give a greenish-yellow precip- itate of Hggia, which is glowly decomposed by excess of the reagent to Hglg ,2KI + Hg. Sulphides. — HgS or soluble sulphides precipi- tate from mercurous solutions a mixture of HgS + Hg (black). The metallic mercury may be dissolved out by boiling with dilute HNO3 ; the HgS remains insoluble. Oxide. — KOH or NaOH gives black precipi- tates of mercurous oxide (HgaO) when added to mercurous solutions. The precipitate is in- soluble in excess of the alkali. NH4OH also gives a black precipitate of mer- curic ammonium salt and metallic mercury, in- soluble in excess of ammonia. Eeducing agents, as SnClg , FeSOi , etc., quickly reduce mercurous solutions, giving a precipitate of metallic mercury (gray or black). 124 MERCtTET. SECOND DIVISION NOT PEECIPITATED BY HYDEOCHLOEIC ACID. Mercwric Salts. With the exception of the chloride, most of the mercuric salts require for their solution in water the presence of some free acid ; otherwise they decompose, giving a precipitate of basic salt. Like the mercurous salts they are all volatile on ignition. HgS. — If HgS or ammonium sulphide be added gradually to mercuric solutions, a white precipitate is first formed, which changes, on fur- ther addition of the reagent, to yellow, then brown, and finally, when the reagent is in excess, to a black precipitate. The lighter-colored precipitates consist of com- binations of the mercuric salt with HgS, in vary- ing proportions (a). The final, black, precipitate is mercuric sulphide (V). ^ {a) SHgCl^ + 2H,S = HgCli,2HgS +^C1 ; (h) HgCla + H^S =HgS + 2HCl. The precipitate is insoluble in nitric acid, even on boiling (a distinction from all other sulphides MERCURY. 125 of the group). On long boiling with strong nitric acid it is changed in color from black to white, the change being due to the nitric acid dis- placing some of the sulphur of the sulphide, and forming a compound similar in composition to the one given by HjjS when not added in excess (a) (Hg(N03)2,2HgS). This, like HgS, is in- soluble in nitric acid, but is soluble in chlorine. HgS is readily soluble in aqua regia (i.e., chlo- rine). It is slightly soluble in boiling dilute hydrochloric, much more so in the concentrated acid. Sodium or potassium sulphides, particu- larly in presence of free alkali, dissolve it, with the formation of double salts. Oxide. — KOH or NaOH added to neutral or slightly acid solutions of mercuric salts give at first a reddish-brown precipitate of basic salt, which changes on further addition of the alkali to yellow mercuric oxide (a). The precipitate is insoluble in excess. In presence of ammonia salts the precipitate is white (V) (a) HgCl^ + 2K0H = HgO + 2KC1 + H^O ; yellow (b) HgCl2,NH4Cl + 2KOH = (NHaHg)Cl + 2KC1 + 2H2O. 126 MEECTTRT. Ammonia produces similar precipitates to the one given in (b) unless the solution contains much free acid, in which case no precipitate is formed : HgCl, + 2NH,0H = NH^HgCl + NH,C1+ 2H,0. Bedtbction to Metal. — Stannous chloride when added to mercuric solutions gives at first a white precipitate of mercurous chloride, which on further addition of the reagent changes to metal- lic mercury, gray or black: 2HgCl2 + SnCl^ = Hg^Cla + SnCl, ; white HgaClg + SnCla = 2Hg + SnCl^ . black Other reducing agents, such as FeS04, SO2, Na2S;j03, etc., also reduce either mercurous or mercuric solutions to the metallic state. If a piece of bright, clean copper be placed in a neutral or slightly acid solution containing either mercurous or mercuric salt, the copper will soon be covered with a gray coating of mercury. If this deposit be gently rubbed with a piece of filter- paper or cloth, it becomes bright and lustrous like polished silver. On gently warming, the coating volatilizes and the color of the copper reappears. BISMUTH. 127 If any dry salt of mercury, either mercurous or mercuric, be mixed with dry sodium carbonate and heated in a glass tube closed at one end, metallic mercury will be volatilized and condense in the cold upper portions of the tube as a metal- lic mirror. BISMUTH (Bi). At. wt. 210; sp. gr. 9.9. Bismuth is a hard, brittle metal, of a white color with a reddish lustre. Fuses at 264° C. It is very slightly oxidized by the air at ordinary temperatures, but very rapidly at a red heat. It is almost insoluble in hydrochloric acid; boiling sulphuric converts it into sulphate, but its best solvent is nitric acid : 2Bi + 8HN03= 2Bi(N03)3 + 2N0 + 4H2O. There are several different oxides of bismuth, but the only one forming stable salts is the oxide Bi,03. Precipitation hy Water. — Bismuth salts require for their solution in water the presence of some free acid, otherwise they decompose into basic salts Avhich precipitate. This action is very char- acteristic and important, and is made much use of in analysis. 128 BISMUTH. If to a solution of bismuth chloride an excess of water be added, the greater part of the bismuth is precipitated as an oxychloride (a). With solution of the nitrate, the precipitate is variable, depend- ing on the quantity of water added (b and c ). In this reaction a portion of the bismuth is always held in solution by the nitric acid liberated. With bismuth chloride the precipitation is com- plete if excess of free acid has been removed by evaporation or neutralization before making the test. In solutions of the nitrate the delicacy of the reaction is greatly increased if to the solution, which should be only slightly acid and concen- trated, a solution of a chloride, as sodium chloride, be added. The precipitate obtained is the oxy- chloride and not the basic nitrate (d). {a) BiCla + H3O = BiOCl + 2HC1 ; {b) Bi(N03)3 + 2H,0 = Bi (OH)2N03 + 2HNO3 ; (6") 2Bi(O H)3N0 3 =(BiO),(OH)N03 + HNO3 + H,0 ; {d) Bi(N03)3 + NaCl + H20 = BiOCl + NaNOg + 2HNO3. The above precipitations are prevented by cer- tain organic acids, as acetic, citric, etc. BISMUTH. 129 The precipitates are easily soluble in HCl and HNO3 , but not in tartaric, the latter being a dis- tinction from antimony. Bismuth Sulphide (^Bi^S^). — H^S or soluble sulphides precipitate black bismuth sulphide from solutions of bismuth salts. The precipitate is in- soluble in cold dilute acids and alkaline sulphides. Soluble in hot dilute nitric acid. 2Bi(N03)3+ 3H2S = Bi,S3 + 6HNO3; Bi2S3 +8HN03= 2Bi(N03)3+ 2N0 + 4H,0 + 3S. Bismuth Hydroxide {Bi(^OII)^. — Potassium, sodium, and ammonium hydrates give with solu- tions of bismuth salts a white precipitate of bis- muth hydroxide, insoluble in excess of the reagents. The precipitation by ammonia is a distinction from copper and cadmium : Bi(N03)3 + 3K0H = Bi(0H)3 + 3KNO3 Tartaric acid, citric acid, and certain organic substances prevent the precipitation. Bismuth Dioxide {BioO^. — If a solution of potassium stannite (made by adding KOH to stannous chloride until the precipitate which first forms has all dissolved) containing an excess of alkali be added to a solution of bismuth salt, a 130 BISMUTH. black precipitate of bismuth dioxide is formed. The reaction is delicate and, in absence of mer- cury salts, characteristic : 2Bi(N03)3 + K^SnOa + 6K0H = Bi^O, + KgSnOa + 6KNO,3. Bismuth Oarhonate {{BiO')iCO^. — The alka- line carbonates precipitate bismuth salts, giving a basic carbonate. Bistnutli Ohromate. — Soluble chromates precip- itate bismuth solutions, giving yellow precipitates varying in composition according to conditions. With potassium dichromate it is (BiO)^^^©^ : 2Bi(N03)3 + SKA'aO, + 4H2O = {BiO\Cv^O, + 6KNO3 + 4H,Cr04. Bismuth chromate is easily soluble in nitric acid, and insoluble in sodium or potassium hy- drates (distinction from lead). Reduction to Metal. — Bismuth is obtained fi'om its solutions in the metallic state by addition of iron, zinc, etc., in the same way as for lead. With sodium carbonate on charcoal, before the blowpipe the reaction is very similar to lead except that the bead is brittle. COPPER. 131 COPPER (Cu). At. wt. 63.6 ; sp. gr. 8 9. Copper has a charactei'istic red color, is very malleable and ductile, and is one of the best con- ductors of heat and electricity. Fuses at about 1100° C. It is contained in many important alloys, as bronze (Cu and Sn), brass (Cu and Zn), German silver (Cu Zn Ni), etc. Copper dissolves easily in nitric acid (a), and is also soluble in hot concentrated sulphuric acid (i). Hydrochloric acid has very little action on it. {a) 3CU+8HNO3 = 3Cu(N03)3 +2N0 +4H3O ; {h) Cu + 2HaS04 = CuSO^ + SO, + 2HgO. Copper forms two oxides, the cuprous (CugO) and cupric (CuO), and two corresponding series of salts. The cuprous salts are seldom met with, are unstable, and easily changed to cupric salt and metallic copper. They ai'e nearly all insoluble in water. Ouprio Salts. — The ordinary copper salts all belong to this class. In the form of crystals or in solution they are of a green or blue color. The anhydrous salts are white. Coppei- Sulphide. — H3S and soluble sulphides 132 COPPER. precipitate cupric salts, giving a black precipitate of sulphide (2CuS,Cu2S). This precipitate is easily soluble in hot dilute nitric acid, slightly soluble in ammonium sulphide, soluble in potas- sium cyanide (distinction from cadmium and all the other sulphides of the group), insoluble iu boiling dilute sulphuric acid (another distinction from cadmium). 3(2CuS,Cu2S) + 32HNO3 = 12Cu(N03)2 + 8Nb + 9S + 16H,0 ; Sodium thiosulphate (NajS^O,,) in hot solutions of copper salts gives a black precipitate of cu- prous sulphide. In solutions strongly acid with HCl this is a separation from cadmium. Owpric Hydroxide (^Cu^OH')^^. — Sodium or potassium hydrates precipitate from ciipric solu- tions the blue hydroxide, insoluble in excess, soluble in ammonia and in acids. On heating the precipitate, while still sus- pended in the liquid, it becomes black, changing to 3CuO,H30. In presence of sufficient tartaric, citric, or ar- senious acids, grape-sugar, etc., the alkalies fail to give a precipitate, a blue solution being obtained. If to the alkaline tartrate solution some grape- COPPEE. 133 sugar be added and the solution boiled, a red or yellow precipitate of cuprous oxide is formed. This reaction is frequently used to determine the presence of grape-sugar in solutions. Ammonium hydrate and carbonate when added to copper solutions gives at first a blue precipi- tate of a basic salt, which readily dissolves on further addition of ammonia to a deep-blue liquid. From copper sulphate the blue solution contains CuS04,4NH3+ HgO; with the chloride the solu- tion contains CaCl3,4NH3 + H3O. Potassium cyanide decolorizes the blue solution, owing to the formation of double cyanide. Oiipric Carhonate. — The fixed alkaline car- bonates precipitate from cupric solutious a green- ish-blue basic carbonate of variable composition. Caproris Oyanide. — Potassium cyanide when added to cupric salts gives a precipitate of cuprous cyanide, or cuproso-cupric cyanide, easily soluble in excess. Cupric Ferrocyanide (C'wai^eC^ye). ^Potassium ferrocyanide, when added to a solution of cupric salt gives a very characteiistic I'eddish-brown precipitate of cupric ferrocyanide : 2CuS0, + K^FtiCye = Cu^FeCy^ + 2K,S04. 134 CADMItTM. The precipitate is not affected by dilute acids, but it is decomposed by alkalies : Cu^FeCye + 4K0H = Cu(OH), + K.FeCye. Reduction to Metal. — Copper is easily precipi- tated in the metallic state, from its solutions, by iron, zinc, etc. A delicate and convenient way to make the test is to put a few drops of a dilute solution of copper in a platinum capsule, acidify witb a drop of acid, and then add a small piece of zinc. The copper will usually be found adher- ing to the platinum, and can be recognized by its characteristic color. Flame Test. — Copper salts, as well as the metal and its alloys, if held on a platinum wire in the flame of a Bunsen burner or blowpipe, color the flame green. Moistening the test with a drop of concentrated hydrochloric acid adds much to its delicacy. With the borax bead in the oxidiz- ing flame copper compounds give a blue color. CADMIUM (Cd). At. wt. II2 ; sp. gr. 8.8. Cadmium is a tin-white malleable metal, fusing at 315° C. It dissolves slowly in hot dilute hydrochloric or sulphuric acids with evolution CABMITJM. 135 of hydrogen. It is more easily soluble in nitric acid, which is its best solvent : Cd + 2HC1 = CdCl, +2H; Cd + BSOi = CdSO, + 2H ; 3Cd + 8HNO3 = 3Cd(N03), + 2N0 + ffl^O. Cadmium resembles zinc in its chemical prop- erties. It forms only one oxide, CdO, and a corre- sponding series of salts. Cadmimn Sulphide (^CdS). — HgS or soluble sulphides give with cadmium salts a yellow j)re- cipitate of sulphide, insoluble in cold dilute acids, alkaline sulphides, or potassium cyanide. Sol- uble in hot dilute nitric acid, also in boiling dilute sulphuric (the latter being a distinction from copper). Cadmium Hydroxide (^Cd(^OII)^. — Potassium and sodium hydrates precipitate cadmium solu- tions, giving a white hydrate, insoluble in excess of the alkali, soluble in acids. With ammonium hydrate the same precipitate is formed, in absence of ammonium salts, which is very easily soluble in excess of ammonia. Cadmium Carbonate {GdCO^. — Potassium, so- dium, and ammonium carbonates precipitate cad- 136 CADMIUM. mi urn carbonate, white, insoluble in excess of the precipitant. Soluble in potassium cyanide. Free ammonia prevents the precipitation. Cadmium Cyanide {GdCy.^. — Potassium cya- nide gives with cadmium salts a white precipitate of cadmium cyanide, easily soluble in excess of leagent, forming double cyanide. From this solu- tion the cadmium is precipitated by HgS as sul- phide : CdCNOa)^ + 2KCy = CdCy, + 2KNO3 ; CdCy3+ 2KCy = CdCy3,2KCy ; CdCys ,2KCy + 2H,S = CdS + 4HC5r + K,S. Blowpipe Test. — Cadmium compounds fnsed on charcoal with dry sodium carbonate are re- duced to metal, which volatilizes and forms a characteristic brown coat on the coal. BECAPITULATION. On reviewing the reactions of the metals of the fifth group we at once see that they differ from all of the metals of the preceding groups in the fact that they are precipitated from solutions acid with the strong acids, by HgS, and it will be SEPARATIONS — FIFTH GROUP. 137 found later on that they differ from the sixth group in the insolubility of their sulphides in alkaline sulphides, the sixth-group sulphides be- ing soluble. A further distinction in regard to three of the metals (Ag, Pb, Hg (ous)) is found in the insolu- bility of their chloiides in dilute hydrochloric acid. This reaction enables us to separate these metals, not only from those of all the other groups, but also from the remaining metals of the fifth. Let us suppose, for example, that we have a solution containing metals of all six groups ; hy- drochloric acid is added to it in slight excess, and the mixture filtered. The precipitate will con- tain silver, lead, and mercurous chlorides, and in the filtrate will be all of the other metals. This filtrate is now acid with hydrochloric acid, owing to the fact that more than sufficient to precipitate the chlorides was used in the first precipitation. If now we add H^S in excess to the solution, the remainder of the fifth group and all of the sixth will yjrecipitate. On filtering, these metals will be separated from those of the first four groups, which will be found in the filtrate. If now we make use of the fact that the fifth- 138 SEPAKATIONS— FIFTH GROTTP. group sulphides are insoluble in ammonium sul- phide, and those of the sixth soluble, we can easily separate one from the other. Doing this, we have a residue containing the second division of the fifth group (Pb (?), Hg (ic), Bi, Cn, Cd) and in the filtrate the sixth group as sulphur salts. Going back now to our precipitate of the chlorides (AgCl,PbCl2,HgaClg) we see, on study- ing their properties, several points in which they differ from each other. Lead chloride, for ex- ample, is soluble in boilings Avater, the others insoluble: if, then, we treat a precipitate of the three with hot watei-, the lead chloride will all be dissolved, and on filtering will be found in the filtrate, where its presence may be proved by add- ing sulphuric acid, which, as we already know, precipitates lead and no other metal of the group. In the residue left after removal of the lead chloride we have only silver and mercurous chlorides, which, you will remember, differ from each other in their action with ammonia, the sil- ver chloride being soluble, and the other not only insoluble, but it changes in color from Avhite to black, which change is characteristic for mer- curous salt. SEPARATIONS — FIFTH GROtTP. l39 From the ammoniacal solution, you know, the silver chloride can be reprecipitated by acidifying with nitric acid. By this series of reactions we have separated these three metals from all of the groups and from each other. Let iis return to our residue of sulphides (PbS (?), HgS, BioSg, CU4S3, and CdS), and see in what way it can be separated. Hot dilute nitric acid will dissolve all of these, with the exception of the mercuric sulphide (HgS); so if we boil the residue with dilute nitric acid, the lead, bismuth, copper, and cadmium sulphides go into solution as nitrates, and the HgS remains insoluble, which after filtering can be dissolved in a little aqua regia {j,.e., CI), and on the solution of mercuric chlo- ride obtained any of the characteiistic tests for mei'cury may be made — prefei'ably the precipita- tion with stannous chloride (SnCljj). In the filtrate containing the nitrates it is al- ways necessary to look for lead, since, owing to its being somewhat soluble in water, a portion is always likely to be found at this point. Making use of the action of sulphuric acid on lead salts, we can readily remove any lead that may be present. Filtering, only bismuth, copper, and cadmium 140 SEPAEATIONS — FIFTH GROUP. are left. The action of ammonia on these salts enables us at once to separate the bismuth as hydroxide, the copper and cadmium going into solution as ammonia compounds — copper, if present, giving the solution its characteristic blue color. As the formation of a white precipitate on the addition of ammonia is not conclusive proof of the presence of bismuth, it is always necessary to dissolve the precipitate in a little dilute hydrochloric acid and make the very char- acteristic water test. The copper and cadmium may be determined in the ammoniacal solution by the reactions with cyanides. If further test for copper besides the blue color of the solution is desired, it is made by acidifying a portion of the solution ^^'ith acetic acid, and then adding potassium ferrocyanide, copper, if present, giving its well-known red ferrocyanide. You perhaps remember that copper sulphide is soluble in potassium cyanide, and that cadmium sul|.iliide is insoluble , if, then, we add sufficient of the cyanide to change the copper and cadmium to double cyanides, and pass HgS into this solu- tion, the cadmium is precipitated as yellow sul- phide. SEPARATIONS — FIFTH GROUP. 141 We might make use of some of the other reac- tions of these metals, as the action of hot dilute sulphuric acid on the sulphides, to effect their separation, but the one with the cyanide is reli- able and easy to perform. In Scheme No. I the plan adopted is that Just outlined. METALS OF THE SIXTH GROUP. TIN, ARSENIC, ANTIMONY, GOLD, AND PLATINUM. These metals, like those of Group V, are pre- cipitated by HjS from acid solutions, but differ from the latter in the solubility of the sulphides in alkaline sulphides, witli formation of soluble thio-salts. The metals of this group may be conveniently divided into two divisions : arsenic, antimony, and tin comprising the first, and gold and plat- inum the second. Those of the first division form oxides, stable at a high temperature, while those of the second are characterized by the easy reduction to metal of all their compounds. The sulphides of gold and platinum are insoluble in either boiling hydrochloric ov nitric acids, which is a distinction from the sulphides of arsenic, antimony, and tin. Thio-salts. — These salts of the sixth group are of so much importance that it will be best to try 143 GKOUP VI. 143 and get a clear conception of their character before proceeding any further. Tin, antimony, and ai'senic form two series of oxides; viz., stannous (SnO), stannic (SnO^), antimonious (Sb40e), antimonic (Sb^Oj), arsenious (AS4O6), and arsenic (As^Oj). These oxides act as the anhydrides of corresponding acids, e.g, arsenious acid, HgAsOg (As^Og + 6H2O = 4H3ASO3), arsenic acid, H3ASO4 (AsgOg + 3HaO = 2H3A8O4). Stannous oxide generally acts as a base, but in a few compounds it plays the part of a weak acid corresponding to HsjSnOg. Stannic oxide is more distinctly acid in its properties, and much more stable, and forms the acid HaSnOa. The two oxides of antimony act in much the same way as tin, only the acid characteristics are more decided. There is a corresponding series of sulphides, viz., SnS, SnSa, SbgSg, SbA, AsgSg, AsaSg. These sulphides, particularly the higher ones, correspond in many ways with the anhydrides of the acids just spoken of, combining with other sulphides to form salts, in the same way that the oxides do ; e-.g., SK^b + AsPs = 2K3ASO4. 144 GROUP VI. Now suppose the oxygen in this equation to be replaced by sulphur, and we get the equation SKaS-fAsaSg^: 2KgAsS4 (potassium thio-arsenate). Antimony sulphide forms a similar compound, (NH4)3SbS4, corresponding to ortho-antiraonic acid {HaSbO^). Stannic oxide combines with oxides of the metals to form stannates, such as K^O + SnOo = KoSnOg. If we treat this equation as we did the previous one and replace all the oxygen with sulphur, we have KaS + SnSg = KgSnSg (potassium thio-stannate). It is evident from this comparison that the thio-salts correspond to the oxygen compounds, su^hur replacing oxygen atom for atom. Many of the thio-salts are much more complex than the examples given, and correspond to the meta- and pyro-salts of the oxygen acids, e.g.^ Na4As2S7 (sodium pyro-thio^rsena^e), corresponding to TIN. 145 TIN (Sn). At. wt. 119; sp. gr. 7.3. Tin is a white, lusti-ous, and malleable metal, fusing at 235° C. Fused in the air, it is rapidly changed to oxide, forming a white powder used for polishing, under the name of putty powder. Its best solvent is hot strong hydrochloric acid, in which it dissolves to stannous chloride (SnCl^) (a). Aqua regia or chlorine dissolves it readily, forming stannic chloride (SUCI4) (b). With nitric acid the action varies accoi'ding to the strength of acid. Strong acid yields metastannic acid (HioSujOij), insoluble in acids and in water (c) ; dilute acid gives either stannous nitrate (Sn(N03)2), stannic nitrate (Sn(]S'03)4), or a mixture of the two, depending on temperature and strength of the acid. In dilute sulphuric acid it is very slowly soluble, but dissolves in concentrated acid to stannic sulphate. (a) Sn+2HCl = SnCl, + 5; (h) Sn + 2C1, = SnCl, ; (c) 15Sn+20HNO3+5H,O=3H,pSn5O,5+ 20NO. 146 TIN. Tin is a constituent of many important alloys, as bronze, solder, etc. ; and it is also largely used in the manufacture of tin-plates, which are iron plates coated with tin. Tin forms two oxides, stannous (SnO) and stannic (SnOg), and corresponding salts. Stannous /Sa/te. — Stannous salts oxidize very easily to stannic, and consequently are strong reducing agents. On exposui-e to air, stannous chloride, for example, oxidizes to stannic oxy- chloride (SuoOClg). Chlorine or strong nitric acid easily effects its oxidation. Many other substances, as ferric salts, mercury salts, etc., act in the same way : SnCl, + FegCle = SFeCl^ + SnCl4. Reactions with mercury and bismuth salts have already been given. Stannous Hydroxide (^Sni^OIiy^. — The alka- line hydrates and carbonates give with solu- tions of stannous salts a white precipitate! of stannous hydrate. The precipitate is soluble in excess of sodium or potassium hydrate and in acids. SnC]^ + 2K0H = Sn(OH), + 2KC1 ; Sn(OH), + 2K0H = KjSnOg + 2HA TIN. 147 Stannous Sulphide (SnS). — Stannous salts yield with HgS a brown precipitate of stannous sulphide (a), insoluble in dilute acids and in colorless ammonium sulphide and in ammonium carbonate. It is soluble in potassium or sodium hydrates and, as already explained, in alkaline polysul- phides, Avith formation of a thio-stannate of the alkali (i). From its solution in alkali or alkaline sulphide it is reprecipitated by acids (c). Stannous sulphide is soluble in hot sti'ong hydrochloric acid, with evolution of li^S ((i). Concentrated nitric acid convei'ts it into meta- stannic acid : (a) SnCl, + H^S = SnS + 2HC1 ; (b) SnS + (NHOsS, = ) As,S3 + 3(NH,)2S = 2(NH4)3AsS3; (c) 2(NH,)3AsS3 + 6HCl -As,S3+6NH,Cl+3Cs ; (d) As,S3 + 3(NH4),S, = 2(NH4)3AsS, + S ; (e) 2(NH4)3AsS4+6HCl =AsA+6NH4C1+3HS; (/)2As,S3 + 2(NH4)3C03 ^~^^ = NH,AsO, + 3NH4ASS2 + 2Ca, ; (g) NH4AsO,+ 3NH4ASS, + 4HC1 = 2As,S3 + 4NH,C1 + 2IL0. 160 ARSENIC. Silver Arsenite {Ag^AsO^. — If silver nitrate be added to a solution of an arsenite, it produces a yellow precipitate of silver arsenite. In aque- ous solutions of arsenious acid no precipitate is formed ; but if dilute ammonia be carefully added, the precipitate forms at once. The precipitate is very easily soluble in nitric acid and in ammonia, and is somewhat soluble in ammonium nitrate : K3ASO3 + SAgNOa = AggAsOg + 3KNO3. Cupric Arsenite {CuHAsO^. — Cupric sul- phate added to a solution of an arsenite produces a yellowish-green precipitate of arsenite, easily soluble in acids and alkalies : K3ASO3 + CUSO4 + H3O = CuHAsOg + KaS04 + KOH. If the solution containing the arsenite be made strongly alkaline with potassium hydrate, and then only a few drops of a weak solution of cupric sulphate be added, and the solution boiled, a red precipitate of cuprous oxide is obtained : K3ASO3 + 2CUSO4 + 4K0H = Cu^O + KaAsO^ + 2K,S04 + ^H^O. ARSENIC. IQI This test is not conclusive for arsenic, since grape-sugar and other organic substances give the same precipitate. It is, however, valuable as a confirmatory test and as a means of distinguish- ing betvi^een arsenious and aisenic acid. As the reaction depends on the oxidation of the arsenite to arsenate, it is evident that the latter could not give the reaction. AESENIO OXIDE. Arsenic oxide is a white amorphous solid, fus- ing at a red heat, and at higher temperatures it volatilizes as arsenious oxide and oxygen. It is slowly soluble in cold water, more easily in hot, going into solution as orthoarsenic acid (H3ASO4). Arsenic acid is very similar in many ways to phosphoric acid. Like the latter it forms pyro and meta acids, and the solubility of its salts is much the same. The only arsenates soluble in water are those of the alkalies. Arsenic acid, like arsenious, is a poison. Arsenic Sulphide. — HgS does not precipitate neutral or alkaline solutions of arsenates, and in acid solutions no precipitate is formed at first, but after long standing one is slowly formed, consisting of a mixture of arsenic and arsen- ious sulphides and sulphur. If the solution be Ib2 AESENIC. moderately acid, and heated to 70° C, and a strong current of H^S be passed into it, the precipitate will be principally ai-senic sulphide. When admissible arsenates can be very readily precipitated as sulphide, by first reducing the arsenate to arsenite by adding sulphurous acid (a), or, what is equivalent, a sulphite to an acid solu- tion (5), warming, and then precipitating with H,S. {a) K3ASO4+ SO3 + H2O = K3ASO3 + ILSO, ; (h) Na,S03 + 2HC1 = SO3 + 2NaCl + H^O It is seen from (a) that sulphuric acid is pro- duced by the reduction of the arsenate, conse- quently in any solution containing metals forming insoluble sulphates, this method could not be used to advantage. The solubilities of arsenic sulphide are analogous to those of arsenious sulphide al- ready given. Silver Arsenate (^A.g^A.sO^). — Silver nitrate added to solutions of arsenic acid or arsenates gives a reddish brown precipitate of arsenate. The precipitate is easily soluble in nitric acid and am- monia, and also somewhat soluble in ammonium nitrate. ARSENIC. 163 Magnesimn Ammonium Arsenate. — If to a solution of arsenic acid or an arsenate soluble in water a mixture of magnesium sulphate, am- monium chloride, and excess of ammonia (known as magnesium mixture) be added, a white crys- talline precipitate of magnesium ammonium ar- senate is obtained. It is easily soluble in acids. If a small portion of the washed precipitate be placed on a watch-glass and dissolved in a drop of dilute nitric acid, silver nitrate added, and the solution obtained neuti-alized very cautiously with ammonia, the characteristic red arsenate of silver is precipitated : K3ASO4 + MgCl3,NH4Cl + NH4OH = MgNH4As04 + 3KCI + NH4OH. Hydrogen Arsenide (AsNs). — If a solution of arsenious or arsenic acid or any of their com- pounds (except sulphides) be treated with zinc and dilute sulphuric or hydrochloric acids in the same manner as given for antimony, gaseous hydi'ogen arsenide will be formed (a), which can be ignited in the same way as the antimony compound, giving the flame a bluish tint, and white fumes of arsenious oxide are formed (b). If a piece of cold porcelain be held in the flame, 164 ARSENIC. a brownisli-black stain of solid hydrogen ar- senide, having a decided lustre, is obtained. The deposit, unlike that of antimony, is soluble in sodium hypochlorite (c). When arsenic and antimony are present together, the use of sodium hypochlorite as a means of separation is not reli- able. A better plan is to add a drop or two of ammonium sulphide to the black stain obtained, evaporate at a gentle heat, and then place the piece of porcelain with the stain, now changed to sulphide, downward, over a small beaker containing fuming hydrochloric acid. If antimony only is present, the orange-colored residue disappears, vol- atilizing as chloride ; but if arsenic be present, the yellow arsenious sulphide remains (Akdekson). If hydrogen arsenide be passed into a solution of silver nitrate, a black precipitate of metallic silver is formed, and the arsenic goes into solu- tion as arsenious oxide (d). After filtering out the silver, the arsenic may be precipitated in the filtrate as silver arsenite, by adding a few drops of silver nitrate, and then neutralizing very care- folly with dilute ammonia (e). (a) H3ASO3 + 3Zn + 6HC1 = A^H3 + ''^ZnCl, + 3HgO; GOLD. 165 (h) 4ASH3 + 60, = As^Oe + 6H2O ; (c) 2ASH3 (solid hydride, or stain) + SNaClO = 2H3ASO4 + 8NaCl ; (d) AsHa + 6AgN0, + SH^O = H3ASO3 + 6Ag + 6HN0, ; (e) H3ASO3 + 6HNO3 + 3 AgN03 + 9NH4OH = Ag3As03 + 9NH4NO3 + 9H,0 Heduction to Metallic Arsenic. — If arseniles, arsenates, or the sulphides are fused with a mix- ture of three pai'ts of sodium carbonate and one of potassium cyanide, the arsenic is reduced to the metallic form. The test is made in a glass tube blown into a small bulb at one end, the mix- ture is introduced into the bulb and heated, the reduced arsenic volatilizes, and condenses as a dark mirror ia the upper part of the tube. If any compound of arsenic be mixed with sodium carbonate, and fused on charcoal in the reducing-flame of the blowpipe, the highly charac- teristic garlic odor will be observed. GOLD (Au). At wt. 197.2 ; sp. gr. 193. Gold is a yellow metal, very lustrous, soft, and exceedingly malleable. Precipitated from solu- 166 GOLD. tion it varies in color from brown to nearly blaci. It fuses only at a high temperature (about 1035° C), and does not oxidize upon ignition in the air. It is insoluble in any of the acids alone, but dissolves easily in aqua regia or any solution con- taining chlorine, yielding a solution of auric chlo- ride (AuClg). Solutions containing free bromine or iodine also dissolve it. The alkaline cyanides in presence of air or oxy- gen dissolve gold, and although the action is slow, it is of great commercial importance, a veiy large amount of gold being obtained from its ores by leaching them with a solution of potassium cyanide : 4Au + 8KCy + O3 + 2H5,0 = 4KAuCy8+ 4K0H. It is not attacked by fusion with acid potas- sium sulphate (a distinction fi'om almost all of the metals). Gold is found very widely distributed in nature, although as a rule in very small quantities. It is almost always found in the metallic state alloyed with more or less silver. Iron and copper pyrites often contain it, and it is a disputed point as to whether the gold is there present as a sul- GOLD. 167 phide or in the metallic state. Tellurium com- bines with gold to form a telluride. Grold forms a considerable number of double salts, as KAuCl^ , KAuCja , etc. Strong ignition decomposes all gold salts, leav- ing a residue containing metallic gold. Oxides. — There are two oxides of gold, aurous, AujO, and auric, AugOa , and two corresponding series of salts. Both of the oxides are dark brown or black, and decompose on ignition into metal and oxygen. They dissolve readily in hydrochloric acid, but not in sulphuric or nitric. Aurous chloride is decomposed by water into auric chloride and metallic gold. Auric chloride is soluble in water, giving a reddish-brown color in concentrated solutions and yellow in dilute, the color being visible even when greatly diluted. Action of the Alhalies. — Potassium hydrate in very concentrated solution gives a brown precipi- tate that dissolves in excess of the reagent, form- ing potassium aurate, KgAugO^. Ammonia gives in concentrated solutions a reddish-yellow precipitate of auric oxide combined with NHg, called " fulminating gold. ] 68 GOLD. Auric oxide is most readily obtained by pre- cipitating with magnesia (MgO) or its hydro- xide. Gold Sul/pMde. — HgS in neutral or acid solu- tions gives a brownish-black precipitate which is paid to be Au^S^ when precipitated from cold solutions, and a mixture of the same with metallic gold and sulphur if the solution is hot. The precipitate is insoluble in hydrochloric and nitric acids even on heating, but is soluble in aqua regia (CI). It is also soluble in the alkaline sulphides, particularly on heating; at least that portion which is present as sulphide is dissolved. Ammonium sulphide also precipitates gold as Au^S^ , soluble in an excess of the reagent. Redihction to the Metallic State. — Gold is easily reduced to the metallic state from its solutions by maoy different reagents, among the most impor- tant being ferrous and stannous salts, oxalic acid, sulphurous acid (in hot solution), and the metals, -particularly zinc. Action of Ferrous Salts. — Ferrous salts when added to solutions containing gold as chloride or bromide give a dark-brown precipitate of me- tallic gold. The solution coutaining the sus- pended precipitate has a characteristic blue-black GOLD. 169 color, especially if the original solution was veiy dilute. 2AUCI3 + 6FeS04 = 2Au + 2Fe,(iiO,% + Fe,Cl,. Oxalic Acid, when added to a solution of the chloride or bromide, which must be free fi'om nitric acid, and contain little or no hydrochloric, gives on warming the solution a precipitate of metallic gold which separates in flakes or is de- posited as a mirror on the sides of the vessel. 2AUCI3 4- 3H AO4 = 2Au + 6C0, + 6HC]. Stannous GMoride containing a little stannic salt gives in dilute acid or neutral solutions of gold a brown or purple -precipitate (" purple of £!assius"). Potassium nitrite also precipitates gold in the metallic state, even from very dilute solutions. MetalliG Zinc is a very valuable precipitant for gold, precipitating it not only from acid solutions, but also from its solution in cyanides : 2KAuCy2 + Zn = 2Au + 2KCy,ZnCy2. Mercury. — Gold dissolves readily in mercury, forming an amalgam from which the gold can be 170 PLATINUM. easily regained by ignition, the mercury vola- tilizing. This is the basis of the well-known amalgamation process for the extraction of gold from its ores. PLATINUM (Pt). At. wt. 194.9 ; sp. gr. 21.5. Platinum when compact is of a steel-gray color, is exceedingly malleable and ductile, and very infusible. Precipitated platinum is black. Its f using-point is about 1 775° C. ; ignition in the air does not alter it. Like gold, it is insoluble in any one of the acids, and also when fused with acid potassium sulphate. It is soluble in aqua regia or chlorine, the solution consisting of hydrochloroplatinic acid (H^PtCle). The alkaline nitrates and their hydroxides with access of air oxidize platinum at a red heat. For this reason fusion with the alkaline hydrates must never be made in platinum vessels. Platinum very readily forms alloys with metals, particularly with easily reducible ones. Platinum forms two oxides — platinous (PtO) and platinic (PtO,). Platinic hy droxide(Pt(OH)4) is easily soluble in dilute acids and in sodium hydroxtde. All of its salts are decomposed on PLATINUM. 171 ignition, yielding a residue of metal. Platinic chloride at a low red heat is converted into plat- inous chloride, and this at a higher temperature gives metallic platinum. Cuprous chloride also reduces platinic chloride or potassium platini- chloride to the platinous salt. Platinic Sulphide {Pt8^. — H„S in cold acid or neutral solutions precipitates the sulphide slowly and incompletely, and only after the solution has been kept saturated with the gas for a long time. In hot solutions the precipitate forms much more rapidly. Platinic sulphide is insoluble in all acids when used alone, but is soluble in aqua regia or chloiine. The alkaline polysulphides in large excess and with the aid of heat dissolve it, but only veiy slowly. From these solutions it is reprecipitated by acids. Precipitation hy Potassium or Ammonium Salts. — Neutral or slightly acid solutions of platinic chloride if not too dilute are precipitated by potassium or ammonium chloride as yellow crystalline precipitates of potassium or ammonium platinichlorides {K^iQ\ or (NH4)3PtCl6). With dilute solutions it is necessary to evajD- orate on the water-bath almost to dryness and treat the residue with dilute alcohol. 172 SEPARATIONS. The potassium precipitate on ignition is d&- composed, yielding a residue of metallic platinum and potassium chloride. The ammonia salt leaves only metallic platinum. Redticing Age?its. — Ferrous sulphate after long boiling produces a precipitate of metallic plati- num. Stannous chloride gives no precipitate, but changes the color of the solution to a dark red, due to the reduction of the platinic salt to plat- inous : PtCl^ + SnCl^ = PtCla + SnCl4. Potassium nitrite gives no immediate precipi- tate, but after standing for some time a yellow precipitate is thrown down, KoPt(NO^)4. Oxalic and sulphurous acids do not precipitate platinum solutions even on boiling. Metallic zinc precipitates platinum very rapidly from its solutions. SEPARATION OF GOLD AND PLATINUM. These metals are very easily separated from each other by adding ammonium chloride to a solution of their chlorides, evaporating on the watei'-bath nearly to dryness, and extracting the SEPARATIONS. 173 residue with a]cohol, whicli dissolves tte am- monium aurichloride, leaving the ammonium pla- tinichloride. To the alcoholic solution contain- ing the gold water is added, and the solution evaporated on the water-bath to expel the alcohol. The gold is then precipitated by ferrous sulphate or oxalic acid. The yellow precipitate containing the platinum on ignition leaves a residue of metallic platinum. In practical work gold and platinum are seldom met with except in alloys (in ores they are de- termined by fire -assay). From these they are separated by treatment with acids or fusion with acid potassium sulphate, which leaves the gold and platinum undissolved. Fusion with sodium- potassium carbonate and nitrate reduces gold and platinum compounds to the metallic state. After removal of the other metals by treatment with acids the gold and platinum are dissolved in aqua regia and then separated from each other, as already given. When present in solution they will, if not re- moved by some of the methods given, be found either wholly or partially with the other metals of the sixth group. Owing to the great difficulty with which platinic sulphide dissolves in ammo- 174 SEPARATIONS. nium, a portion of it at least will be found with the residue of mercuric sulphide left after dis- solving the fifth-group sulphides in nitric acid, as it is insoluble in nitric acid when precipitated from hot solutions. Its presence may be determined in this residue by drying the mixture of the two sulphides and heating them in a glass tube closed at one end. The mercuric sulphide volatilizes, leaving the pla- tinic sulphide or metal, which may be dissolved in aqua regia and tested for platinum. Gold also may be found in the residue with the mercuric sulphide, as it does not always dissolve completely in the ammonium sulphide, and is in- soluble in nitric acid. If present, it will be left with the platinum on treating the residue in the glass tube, and with that will be found in the aqua regia solution. The gold and platinum sulphides that are dis- solved by the ammonium sulphide will, if the Marsh apparatus is used, be found in the metallic state ' in the generator, and after removal of the zinc and tin by hydrochloric acid may be dissolved in aqua regia and tested for as already explained. SCHEME v.— SEPARATION OF GROUPS V. AND VI. Add to the so'.ution HCl iu slight excess, warm gently, filter, aud wash precipitate with cold water. (Note 1). Ppt. 1. AgCl + PbCl 2+ HgjCU. Wash on the filter with boiling water until the PbCla has been dissolved. (Note 2.) Filtrate 2— (PbClj). a white ppt.. Add dil. HjSO, PbSO, , proves Pb. Residue 2''. NHgjCl. NH,CI + Hg. (Black.) Dissolve in aqua regia; heat to expel excess of acid, aud add a few drops of SuClj. A white ppt. (turning dark) shows HgjClj, and proves HgjO. (Note 5.) Residue 2. AgCl f Hg.CU. Pour warm NH4OH over the moist ppt. on the filler. (Note 3.) FlIiTKATE 2''. 2AgCl, SNHs. Boil out excess of NHs, add HNO3 to slightly acid reaction. A while, curdy ppt. shows AgCl, aud proves Ag. (Note 4.) Filtrate 1. Saturate with HjS, filter and wash thoroughly. (Note 6.) Ppt. 3. HgS + PbS + CuS + BijS, + CdS + AsaSs + SbjS, + SnS+ SnS, Remove from the filter, and digest iu a small porcelain dish with a little (NH4)2S, at a gentle heat Filter and wash. (Note 7.) Filtrate 3. 1°, 2% 3°, 4° groups. Residue 4. HgS + PbS + CuS + BUS, + CdS. Remove from the filter, and heat with moderately dil. HNO3 in a porcelain dish. Filter and wash. Residue 5. HgS. (Black.) Dissolve in aqua regia; boil out ex- cess of acid, and add a few drops of SnClj. A white ppt. (turning dark) proves HgO. (Note 8.) Filtkate 5. Pb(N03)2 + Cu(N03)., + Bi(N03)3 + Cd(N03).,. Evaporate nearly to dryness, add dilutesulphuric acid, and filter. (Note 9.) Filtrate 6: CuSO, + Bi2(S04)3 + CdSO,. Add NHiOH in excess; filter aud wash. Ppt. 6; PbSO,. Ppt. 7. Bi(OH),. Dissolve in a few drops of dil. HCl oa the filter, and allow to drop through into a test-tube full of water. A white ppt. shows BiOCl, and proves Bi. (Note 10.) Filtrate 7. until the blue color disappears yellow ppt. shows CdS (11), and proves Cd Test a small portion for Cu, by acid- ifying with HC2H3O1, and adding a few drops of KiFeCye (lOi. A reddish- browu ppt. shows CujFeCye, and proves Cu. If Cu is present, the solution will be colored blue. To detect Cd, add KCy saturate the solution (19) with HjS. A bright (Note 11.) Filtrate 4. (NH4)3AsS4 + (NH03SbS. + (NH^j^SnS, + (NH,)3Sx. Acidify with dil. HCl ; filter and wash. (Note 12.) Ppt. 12. As^Se, SUuSs, SnS, + S. Filtrate 12. NH.Cl. Reject. Remove from filter, aud warm in a porcelain dish with HCl + KCIO3. Boil out all free CI, aud filter. (Notes 13 aud 20.) Filtrate 13. HjAsO^ + SbCl3 + SnCl.. Place in a Marsh apparatus, with Zn -(-dil. HjSOi, and conduct the evolved gases (AsHs -j- SbH3 + Hn) into a dil. solution of AgNOs. Filter and wash. (Note 14.) Residue 18. S. Reject. Filtrate 15. H3ASO3 -I- AgNO, + HNO3. Add a few drops of AgNOs, and then very dilute ammonia. Yellow ppt. shows AgsAsOs (20), and proves As. (Note 15.) Or add dil. HCl, shake, and filter. Ppt.: AgCl. Reject. Filtrate : HsAsOs. Saturate with H3S. Yellow ppt. shows AsjSa, and proves As. (Note 16.) Ppt. 15 from AgNO,. Ag -I- AgsSb. Boil with HCl -f- H-^CHiOs (16), and filter. Pass HjS through the solu- tion. An orange ppt. shows SbjSs (17), and proves 8b. (Note 17.) Sesidue in the OeneraUrr. Sn -I- Sb -I- Zn. Rinse out into a porcelain dish, wash thoroughly, and boil with cone. HCl. Filter and wash. (Note 18.) Filtrate (18): SnCU -|- ZnCl,. Boil off excess of acid, and add sol. of HgClj. White ppt. (turning dark) shows HgjCla (29), and proves Sn. Residue (18)— (Sb). Dissolve in aqua regia, and boil out excess of acid. Pass HjS into solu- tion. Orange ppt. shows SbaSs, and proves Sb. (Note 19. ) To follow p. 174. NOTES TO SCHEME V. Note 1. — The addition of hydrochloric acid should be made with care, for it is veiy impor- tant that the filtrate be acid, yet a large excess of acid is not admissible, since it prevents the com- plete precipitation of the metals by H^S. It must also be remembered that bismuth and anti- mony salts may be precipitated on the addition of hydrochloric acid, and that a moderate excess of dilute acid is necessary to dissolve them. If on washing the precipitate the washings come through turbid, add a few drops of HCl to the wash-water. If the original solution before the addition of the HCl has an alkaline reaction, it should be re- membered that on adding acid to such a solution we may obtain a precipitate of any of those metals whose hydroxides are soluble in alkali. The pre- cipitate in this case, however, will redissolve on the further addition of acid (except lead). In the case of metals of the sixth group dissolved in alkalies or alkaline sulphides the precipitate is insoluble in moderate excess of acid, and it usu- ally shows a characteristic color. If the original 175 176 NOTES TO SCHEME V. solution is alkaline, and hydi-ocliloric acid gives a colored precipitate, insoluble in an excess of the dilute acid, the precipitate should be filtered off, and after washing, digested with (NH4)jjSp. to dissolve any of the sixth group metals, which are to be tested for in the usual way in the filtrate. The residue insoluble in the (^£[4)28^, is dissolved in hot dilute nitric acid and again tested with hydrochloric acid. Note 2. — It is important that all the lead chlo- ride should be washed out before testing for the silver and mercury, as it interferes with the latter tests. Note 3. — Add the ammonia in small portions at a time, otherwise much of the silver may be left in the residue, owing to a reaction between the metallic mercury formed by the action of ammonia on mercurous chloride and the silver chloride whereby metallic silver is precipitated : 2Hg + 2AgCl = 2Ag + Hg,Cl,. Note 4. — The object in boiling out the excess of ammonia is to avoid the formation of much ammonium nitrate, since this salt has some solvent action on silver chloride. Note 5. — Dissolve this' residue in as little aqua NOTES TO SCHEME V. - 177 regia as possible, then evaporate carefully until all free chlorine is expelled, dilute with a little water, and if there is any residue, filter, and test the residue (probably AgCl) for silver. The filtrate is tested for mercury with a few drops of stannous chloride (SnCl^). Note 6. — As stated in Note 1, it is very impor- tant that the solution to be precipitated by H^S should have a proper quantity of free acid (best HCl) : if there is not enough present, there is danger of zinc precipitating ; and if the solution is too acid, the metals of the fifth and sixth groups are not completely thrown down. It is always a simple matter to determine roughly the quantity of free acid in this solution, by taking a little in a test-tube, and adding solution of sodium carbonate until a permanent precipitate is formed. The amount of sodium carbonate used of course indicates the quantity of free acid. When a solu- tion is found to be strongly acid, it is, as a rule, the best plan to evaporate to small bulk, and then dilute with water. In case bismuth or antimony were present, the addition of water may cause a precipitate; but this can be disregarded, as the HjS will convert it all into sulphide. Instead of removing the excess of acid by evap- 178 NOTES TO SCHEME V. oration, the solution is sometimes largely diluted with water, and then precipitated with H2S. In cases where only the fifth and sixth groups are to be determined, this plan might be used to advan- tage, but not when all of the groups have to be tested for. Neutralization of the free acid by alkali is not advisable. The best method of conducting the precipita- tion is to pass a rapid current of HjS gas into the solution until, on shaking, it smells strongly of HjjS. The precipitate of sulphides should be very thoroughly washed with water containing HgS. When washed with pure water only, certain sulphides are liable to assume the colloidal condi- tion, and pass through the filter, giving a dark- colored filtrate. The filtrate from the HgS precipitate should always be tested, to see if the precipitation has been complete ; either by adding to a portion a large excess of HgS water, or, what amounts to the same thing, pass in H^S gas, and if this gives no more precipitate, add a large excess of water, and again saturate with the gas. The water is added, to so dilute the free acid present that it will not prevent the complete precipitation of the metals as sulphides. Of course, if any precipitate NOTES TO SCHEME V. 179 is obtained by these tests, then the whole of the filtrate must be treated in the same way. Whichever way the filtrate is tested, it should always be heated to 60° or 70° C. for some time, in order to precipitate any arsenic acid that may be present. Note 7. — The ti*eatment of the sulphides with (NH4)8S^ is best made by removing the precipi- tate from the filter-paper, placing it in a small beaker, adding sufficient of the ammonium sul- phide to completely cover the precipitate, warm- ing for a few minutes, filtering by decantation, and repeating the treatment of the residue two or three times with small quantities of the (NH4)2Sa.. The final residue must be washed very thor- oughly with water containing a few drops of ammonium sulphide, so as to wash out all chlo- rides ; otherwise the subsequent treatment with nitric acid would form aqua regia, and dissolve the mercuric sulphide. Note 8. — -The residue at this point is not always black, even when mercury is present, since the mercuric sulphide is sometimes changed by the boiling with nitric acid to the, light-colored compound 2HgS,Hg(N03)2. This is tested in the same way as the black sulphide, by dissolving 180 NOTES TO SCHEME V. in chlorine, and precipitating with stannous chlo- ride. It should be remembered that this residue some- times contains lead as sulphate, particularly if the nitric acid used was strong. The best way to test for it is on charcoal with the blowpipe. Tin in the form of metastannic acid is also sometimes found here, owing to the treatment of the sulphides with (NH,)jSa. having been incomplete. Note 9. — Evaporate the solution carefully, al- most to dryness, so as to expel the excess of nitric acid, then before adding sulphuric acid to the whole of it, test a small portion by adding a con- siderable excess of dilute sulphuric acid, allow to stand a few minutes, and if a precipitate forms, add the acid to the remainder of the solution. If no precipitate forms in the first test, do not add any acid to the rest of the solution, but proceed to test for bismuth with ammonia. It is advisable to test a few drops of the con- centrated solution for silver by diluting with water and adding hydrochloric acid. Note 10. — Success in making this test depends on having very little free hydrochloric acid and a large excess of water. Another good way to make the test is to dis- NOTES TO SCHEME V. 181 solve the precipitate on the filter paper m a little dilute hydrochloric acid, allowing the solution to ran through into a small porcelain dish, evaporate until only a drop or two of liquid remains, and pour this into a large test tube full of water. The fact of a precipitate forming on the addi- tion of ammonia in excess is not proof of the presence of bismuth, for if lead or mercury had not been completely removed from the solution they would be precipitated by that reagent. Note 11. — If copper is not present, as shown by the test with ferrocyanide, do not add potassium cyanide to the remainder of the solution, but make it slightly acid with dilute hydrochloric acid, and pass in HjS gas to precipitate the cadmium. If the precipitate for cadmium be dark colored, it is best to confirm it on charcoal with the blow- pipe. (See Cadmium.) Note 12. — Add only enough acid to make solu- tion slightly acid. An excess of acid might dis- solve sulphide of tin. Note 13. — Place the sulphides in a small porce- lain dish, add enough concentrated hydrochloric to cover the precipitate, and heat to boiling, adding from time to time a small crystal of potas- sium chlorate. If antimony or tin sulphides 182 NOTES TO SCHEME V. only are present, they will dissolve in the hot hydrochloric alone, so it is always well to boil a few moments before adding any chlorate, for if everything dissolves (except a little sulphur) no arsenic can be present, and no further test for it is necessary. Continue heating until everything except the separated sulphur has dissolved, and all free chlo- rine has been expelled. Note 14. — The reagents used in this test should always be tested for arsenic and antimony, as these are impurities frequently met with in hy- drochloric and sulphuric acids and in zinc. The surest way to test their purity is to put some of the zinc and acid in a Marsh apparatus, and test the evolved gas for arsenic and antimony by the mirror test on cold porcelain. (See Arsenic and Antimony.) ' The solution supposed to contain arsenic and antimony should always be added in small quan- tities at a time, otherwise the reaction may be- come so violent as to spoil the test. Note 15. — The silver nitrate is added so as to make sure of some being present, since that added in the first place to the solution may have been all precipitated by the arsenic and antimony. NOTES TO SCHEME V. 183 Very dilute ammonia should be used, and it is best to add it so that it forms a layer over the silver solution. This may be done by pouring it very carefully dov^rn the side of the test-tube. At the junction of the two liquids a yellow ring of silver arsenite will be formed if arsenic be present. Note 16. — A small amount of an orange-col- ored precipitate is sometimes obtained here, due to the presence of a little antimony in the solu- tion. Note 17. — When possible, rinse precipitate off the filter into a test-tube with a little water, add a small piece of tartaric acid, and boil. Then add a few drops of hydrochloric acid and filter. If the precipitate given by HgS is not the proper color (orange), it may be easily tested for antimony by dissolving it in a few drops of boil- ing concentrated hyrochloric, and placing the solu- tion in a platinum capsule with a piece of zinc, when, if antimony is present, the characteristic black stairi will make its appearance. Note 18. — The tin in this residue is generally in the foi'm of loose spongy particles, from which the zinc can easily be removed. Antimony, if present, is there as black flakes or powder. 184 NOTES TO SCHEME V. It is best to remove the zinc before dissolving the tin. This is readily done by picking out the hard lumps of zinc with a pair of pincers. Note 19. — It is always well to test in this residue for antimony, if it has not already been found in precipitate 15. If the current of 'hydrogen in the Marsh appa- ratus has been weak, a large amount of the anti- mony is often left here in the metallic state. Note 20. — Another plan for separating these sulphides is as follows : Rinse the precipitate into a beaker, add a lump of ammonium carbonate, and warm gently for a few minutes, filter, and wash. Rbbidtib SiiS.,SbjS.. Dissolve In bot concentrated hydro- chloric acid. Place solution in platinum capsule ■with a piece of zinc. Antimony, if present, gives the vyell-known hlack stain, and the tin is found as a spongy mass. Remove the undissolved zinc, and boil the residue with a little strong hy- drochloric aSid. The tin dissolves as stannous chloride, and after filtering may be tested with meicuric chloride. A further test for antimony may be made by dissolving the residue, if ahy, left by hydrochloric acid, in chlorine, and precipitating with H28. Filtrate (NH4)sAsS.+(NH4)3As04. Acidify with hydrochlo- ric acid, filter, and wash. Ppt. AsiSs. Heat with a little con- centrated nitric acid until dissolved. Test a poi'tion of the solution for arsenic acid with ammonium molyb- date. Another portion test with magnesia mixture. NOTES TO SCHEME V. IPf) This method is not quite exact, as some of the antimony and tin sulphides are liable to go into solution, and a little of the arsenic may be left in the residue. Still it is a rapid method, and for ordinary work sufficiently accurate. Sectioi^ II. THE ACIDS. The detection of the acids cannot be effected by methods similar to those used for the bases. They do not form distinct groups as the metals do, which can be separated from each other by pre- cipitation and filtration. In almost all cases they are determined by spe- cial tests. Although we cannot divide and sepa- rate them group from group, still it has been found convenient to classify them to some extent — this classification depending on their deportment with barium ehloride and silver nitrate ; those acids giving a precipitate with barium chloride in neu- tral solution forming one division or group, and those precipitated by silver nitrate from solutions acid with nitric acid constituting the second divi- sion. Besides these there are two acids (nitiic and chloric) that are not precipitated by any re- agent, and for this reason are put in a group by themselves. The classification just given is for the inorganic 186 THE ACIDS. 1^7 acids only, and they are the only ones that will be considered in the following pages. The presence of organic acids can easily be de- termined by evaporating to dryness and igniting a neutral solution. If present, they leave a car- bonaceous residue. First Group. — Precipitated hy harium cliloride from neuPi'ol solution : Chromic acid (anhydride, CrOg), sulphuric acid (£[^804), phosphoric acid (H3PO4), boric or boracic acid (H3BO3), hydroflu- oric acid (HF), carbonic acid (CO3, anhydiide), and silicic acid (1148104). All these acids, with the exception of hydrofluoric, are precipitated by silver nitrate from neutral solutions. Second Group. — Precipitated hy silver nitrate from sohctions acid with nitric acid: Hydrochloric acid (HCl), hydro bromic acid (HBr), hydriodic acid (HI), hydrocyanic acid (HON or HC}), hydroferrocyanic (H4FeCy6), hydroferricyanic (HeFe^Cyia), hydrosulphuric acid (H38). Ihird Group. — Not precipitated hy any re- agents : Nitric acid (HNO3) and chloric acid (HCIO3). FIKST GEOUP. PEECIPITATED BY BAKItTM CHLORIDE FROM NEUTRAL SOLUTIONS. CHROMIC ACID (Anhydride, CrOJ. Chromic acid (H2Cr04) lias not been obtained in the free state, although many salts corre- sponding to such an acid are known, some of them beautifully crystalline. The anhydride forms beautiful scarlet crystals, deliquescent, and very soluble in water, imparting to the solution a red- disb-yellow color, whichi is visible in very dilute solutions. On ignition it yields chromic oxide (CrgOg) and oxygen. It acts as a very powerful oxidizing agent, and as a caustic on living tissues. A mixture of chromic anhydride and concentrated sulphuric acid forms an extremely active oxidizing mixture. It is used in the analysis of pig iron to oxidize the graphite to carbonic acid : 4Cr03+6H2S04+3C=2Cr,(S04)3+3CO,+6H,0 The chromates are all red or yellow, the color being visible even in very dilute solutions. Those of the first and second group metals, with the ex- 188 CHROMIC ACID. 189 ceptiou of barium, are soluble in water; all the others insoluble, or nearly so. Heduction of OTiwmiG Acid or Ch-omates to Salts of Chromic Oxide (^Gr^O^. — Some of the most characteristic reactions of chromic acid de- pend on its active oxidizing properties, the chromic acid or chromate being at the same time reduced to chromic oxide. The reduction is clearly indicated by the change in color from the red or yellow of the chromate to the bluish-green of the chromic salt. If HgS be added to an acid solution of a chromate a precipitate of sulphur is formed, and the solution changes in color from red or yellow to bluish green : K2Cr20r + 8HCl + 3HaS = CraCle + 2KC1 + 3S + 7H,0. This reaction is of special importance, for it is evident that although chromium may be originally present as the acid, the precipitation of tftie metals with HaS will always reduce it to chromic salt, and it will consequently be found among the bases. The change in color will, however, surely denote its presence. A means of distinguishing between it and the basic chromium will be given later. 190 CHKOMIC ACID. Ammonium sulphide causes the same reduction and then precipitates the chromic salt formed, yielding a precipitate of hydroxide. Many other substances reduce chromic acid or chromates to chromic salt, such as sulphurous acid (a), concentrated hydrochloric acid (b), di- lute; hydrochloric acid and alcohol (a), stannouis chloride or zinc in acid solution (not nitric) (d and e), tartaric and oxalic acids, etc. {a) K,Crs,0, + SSOg + H,S04 (5)K,CrA + 14HCl = Cr^Cle + 2KC1 + SCTg + 7H2O ; (c) K^CrgO, + 9HC1 + 4C2H5OH = CraCle + C2H5CI + SC^H^O + 8H2O + 2KC1 ; {d) K,Cr,0, + 14HC1 + SSnCl^ = Cr.Clfi + SSnCl^ + 2KC1 + TH^O ; {e) KaCr^O, + 3Zn + 14HC1 = Cr^Cle + 2KC1 + 3ZnCl, + lYi^O. All these reductions are made evident by the change in color from red or yellow to bluish green. CHROMIC ACID. 191 Reactions of CJiromates %oith the Bases. Barium Chromate {BaCrO^. — Barium chlo- ride added to solutions of the chromates gi^ea a yellow precipitate of barium chromate (BaCr04)- The precipitate is soluble in hydrochloric aud nitric acids, and decomposed by sulphuric acid. It is nearly insoluble in acetic acid. K,Cr04 + BaCl, = BaCr04 + 2KC1. Lead Chromate (PhOrO/^. — Soluble chromates added to lead salts give a yellow precipitate of lead chromate, slightly soluble in dilute nitric acid, insoluble in acetic acid and in ammonia. Soluble in sodium or potassium hydrates, from which solu- tions it is reprecipitated on acidifying with acid (best to use acetic) : K3Cr04 + Pb(N03)3 = PbCr04 + 2KNO3 ; PbCr04 + 4K0H = K^QvO, + K^PbO^ + 2HaO. Silver Chromate. — Normal chromates in neutral solutions of silver salts give a dark-red precipitate of silver chromate (Ag.jCr04). In slightly acid solu- tions the dichromate is precipitated (AggCrgO^), 192 SULPHtTRIC ACID. Both pi'ecipitates are soluble in nitric acid and in ammonia ; K,Cr04 + 2 AgNOa = Ag.CrO, + 2KNO3. Perchromic Acid (^H^Cr^O^f). — A very beau- tiful and delicate test for chromic acid is made by taking a little of a very dilute and acid solution of hydrogen peroxide, adding a small quantity of ether, and then the solution containing the chromic acid or chromate. The liquid in the tube changes to a beautiful blue coloi", and if the tube be closed and inverted several times the solution becomes (tolorless, and the color is now concentrated in the layei- of ether. This blue color is probably due to perchromic acid, but the compound is so unstable that its composition has not been satisfactorily determined. The following equation may, perhaps, represent the action that takes place : 2R.,CvO, + H^Og = H,Cr,08 + 2H,0. SULPHURIC ACID (H,SO,). Sulphur when burned forms the oxide SOg called sulphurous anhydride, which with water forms sulphurous acid (ElaSOj) — at least the salts formed correspond to an acid of this composition, SULPHUEIC ACID. 193 altliougli the acid itself cannot be isolated, as it at once splits up into sulphurous anhydride (SOj) and water. This oxide can take up one more atom of oxygen, forming sulphuric anhydride (SO3), which combines readily with water, yield- ing sulphuric acid (H2SO4). Sulphuric anhydride when in the solid state is usually in the form of white, silky ciystals. These on exposure to the air quickly absorb moisture and become liquid. Concentrated sulphuric acid is a heavy, oily liquid, colorless and transparent. Both the anhydride and the acid dissolve in water in all proportions, and the solution is ac- companied by the evolution of much heat. Owing to their great aflSnity for water, the concentrated acid and the anhydride char many kinds of or- ganic matter. The boiling-point of the acid in 338° C. The sulphates with few exceptions are colorless. All of the normal sulphates, with the exception of barium, strontium, calcium, and lead sulphates, are soluble in water. In alcohol only the ferric sulphate is soluble. Basic sulphates, which are insoluble in water, are soluble in hydrochloric or nitric acids. 194 SULPHTTEIC ACID. Barium Sulphate. — Barium salts addea to a solution containing sulphuric acid or a sulphate give a very finely divided white precipitate of barium sulphate, even from exceedingly dilute solutions. The precipitate is insoluble in water and in dilute acids. When testing the solubility of this precipitate it should be remembered that concentrated hydro- chloric or nitric acids precipitate barium salts. This precipitate, however, is easily soluble in water; so that if the solution be diluted with water there will be no danger of mistaking one precipitate for the other, since the barium sulphate would remain insoluble, whilst the chloride or nitrate would dissolve: H3SO4 + BaClg = BaSO^ + 2HC1. Lead Sulphate (PbSOi). — Sulphuric acid or soluble sulphates give with lead salts a white precipitate of lead sulphate, the properties of which have already been given under lead. Insoluble Sulphates. — Sulphates insoluble in water and acids may be changed to soluble com- pounds by fusion with sodium potassium car- bonate. SDLPHUEIC ACID. 195 In the fusion the sulphuric acid of the insoluble compound combines with the sodium or potassium carbonate and forms a sulphate soluble in water. The base with which it was combined is changed to carbonate or oxide, insoluble in water : BaSO, + Na^COs = BaCOg + I^a^S04 ; . PbSO^ + Na^COs = PbO + Na,S04 + COa. If the fused mass be treated with water the sulphate of soda dissolves and can be tested for in the solution, after acidifying with hydrochloric acid, by the addition of barium chloride. Instead of fusing with the alkaline carbonate, the same i-esult may be obtained, although not so readily, especially with barium sulphate, by boil- ing with a strong solution of the carbonate. The bases will be changed to carbonates, and the sul- phuric acid goes into solution as sodium sulphate. Blowpipe Test. — If a sulphate be mixed with dry sodium carbonate and fused on charcoal in the reducing flame, it is reduced to sulphide. By placing the fused mass on a piece of bright silver and moistening with water a dark stain of silver sulphide is formed. It must be remembered, however, that any 196 PHOSPHORIC ACID. suljstance containing sulphur will give the same result, so the test is only conclusive when it is known that all sulphur compounds except sul- phate are absent. Detection of Free Sulpliuric Acid. — Free sul- phuric acid can be easily detected by adding a very little cane-sugar to the solution, and evapo- rating to dryness on the water-bath. If any of the acid is present in the free sta,te, a black car- bonaceous residue remains, or if the quantity of acid is very small, simply a brown color. PHOSPHORIC ACID (H,PO.). Common phosphorus is a colorless, transparent solid, insoluble in water, but easily soluble in carbon disulphide. It melts at 44.3° C, and boils at 290° C. It oxidizes when exposed to the air at ordinary temperatures, giving off white fumes which have a characteristic odor and are luminous in the dark. When heated in presence of air it ignites very readily, evolving heavy white vapors of the pentoxide, PgOg. Colorless phosphorus is very poisonous, and for this reason, as well as on account of its great PHOSPHOKIC ACID. 197 inflammability, it must be handled with extreme care. It should always be kept under water, and any pieces required for experiment ought to be cut off while it is still protected in this way. If phosphorus be heated to 250° C. in a veseel to which air has not free access, it is changed to a red modification, called amorphous phosphorus. This differs from the colorless variety in being less inflammable, is not luminous, is insoluble in carbon disulphide, and is not poisonous. Phosphoric anhydride or pentoxide (PgOg) is a snow-white and very deliquescent solid. With water it forms three different acids. If dissolved in cold water it yields metaphosphoric acid, HPO3 , which on standing gradually changes to pyrophosphoric acid, H4P2O7 , and finally to ortho- phosphoric acid, H3PO4. The latter acid is also formed by dissolving the pentoxide in boiling water or by boiling a solution containing either the meta- or pyrophosphoric acids : PA + H,0 =2HP03; PA + 2H,0=H4PA; PA + 3H,0 = 2H3P04. 198 PHOSPHORIC ACID. Salts of the first two acids are also formed by subjecting ortbophospliates containing hydrogen or a volatile base to a high temperature ; as, for example, Na(NH,)HP04 - (NH3 + H^O) = NaPOg , Sodium metaphosphate ; 2MgNH,P04 - (2NH3 + H^O) = MgaPjO^ , Magnesium pyrophosphate. Compounds of pyro- and metaphosphoric acid are seldom met with, almost all the phosphates in nature and in analysis being salts of the ortho- acid (H3PO4). Magnesium pyrophosphate, how- ever, is of importance, since it is the form in which magnesium is generally weighed in analy- sis ; and "phosphorus salt" (]Sra(NH4)HP04), which on fusing into a bead becomes metaphos- phate, is a useful blowpipe reagent. The salts of orthophosphoric acid are all insolu- ble in water, with the exception of those of the alkalies. In acids their solubility varies : those of the alkaline earths are soluble in hydrochloric, nitric, and acetic acids ; ferric and aluminium phos- phates are insoluble in acetic but soluble in the mineral acids. PHOSPHORIC ACID. 199 Heactwns of Orthophosphoric Acid. Barium Phosphate. — In neutral or alkaline solutions of phosphates barium chloride produces a white pre- cipitate of barium phosphate, easily soluble in acetic, nitric, and hydrochloric acid : Na,HP04 + BaClg = BaHP04 + 2NaCl. Magnesiurti Phosphate. — The precipitation of phosphates from neutral or alkaline solutions by magnesium salts is very important. It will be remembered that this is the reaction made use of for determining magnesium, and when we have an alkaline solution it affoi'ds an equally good test for phosphoric acid. Its use for the detection of phosphates, however, is very limited, owing to the fact that the test has to be made in the presence of ammonia; and since all the phosphates, with the exception of those of the first-group metals, require acid to hold them in solution, it can only be used when we have a phosphate of an alkali metal. Arsenic acid also must be absent, as it too precipitates magnesium salts from alkaline solution. In neutral solutions magnesium salts give a white precipitate of magnesium hydrogen phos- 200 PHOSPHOEIC ACID. ptate (MgHP04) (a), or magnesium phosphate (Mg3(P04)2), according to the conditions; but in presence of ammonium chloride and free ammonia a white crystalline precipitate of magnesium ammo- nium phosphate is formed, even in very dilute solu- tions(5). All of these precipitates are soluble in acetic and mineral acids. (a) Na,HP04 + MgCl3.= MgHP04+ 2NaCI. (h) Na^HPO, + MgCl, , NH4CI + NH4OH = MgNH4P04 + 2NaCl + NH^Cl + H^O. Silver Phosphate (Ag^PO^). — Nitrate of silver gives with neutral solutions of phosphates a yel- low precipitate of silver phosphate soluble in nitric acid and ammonia. Na,HP04+3AgN03=Ag3P04+2NaN03+HN03. Perric Phosphate. — If to an acid solution con- taining phosphate an excess of sodium acetate be added and then ferric chloride drop by drop, a yellowish- white precipitate of ferric phosphate will be formed. If the addition of the ferric chloride is continued until the solution assumes a reddish tinge, and the solution boiled, all the phosphoric acid will be precipitated in combination with the PHOSPHORIC ACID. 201 iron, and the excess of iron present will be thrown down as basic acetate. This reaction is of great importance, as it enables us to separate phosphoric acid from its combiiiations with the alkaline earths. 2Na3HP04 + Fe^Cle + 21^3.0^11,0, = 2FeP04+ 6NaCl + 2HC2H3O2 , or 2MgHP04 + Fe^Cle + ^'NnO.'R.O, = 2FeP04 + 2MgCl2 + 2HC2H3O3 + 2NaCl. Ammonium, PhospTiom,olybdate. — Of all the tests for phosphoric acid this is the most valuable, as it enables us to precipitate phosphoric acid not only from acid solutions, but from solutions con- taining almost all the metals. If to a solution of ammonium molybdate be added a few drops of a neutral or acid solution containing a phosphate, a yellow precipitate of am- monium phosphomolybdate((NH4)3P04,12 M0O3) is formed. The formation of the precipitate is hastened by heating the solution gently, not above 70° C. Care must be taken in making this test always to have the ammonium molybdate present in considerable excess, otherwise the precipitate may not be thrown down. Arsenic and silicic 203 BORIC OE BOKACIC ACID. acids, if present in the substance to be tested, must pt'eviously be removed, as they also are precipi- tated by ammonium molybdate, although not so readily as the phosphate. The solution best suited for the reaction is one slightly acid with nitric acid. The precipitate of phosphomolybdate after being washed with a solution of ammonium nitrate may be dissolved in ammonium hydrate and the phosphoric acid precipitated with " magne- sia mixture" (MgCl, , NH4CI + NH4OH). 3ORIC OR BORACIC ACID (H^BO^). Boron is infusible and non-volatile, and when neated in the air burns to boric anhydride (B2O3). This is a colorless glass fusible at a red heat, and soluble in water and alcohol. On evap- orating these solutions boric acid volatilizes with the water or alcohol. For this reason acid solu- tions of borate must always be neutralized before cojicentration. Boric anhydride forms with water metaboric acid (HBOg) and orthoboric acid (H3BO3), also sev- eral other hydrates of more complex composition. The borates of the alkalies are soluble in water, the others insoluble. They are not decomposed BORIC OR BORACIO ACID. 203 on ignition. Although boric acid is a very weak one when in solution, yet iu the dry state it acts very powerfully, driving out the strongest acids from their combinations, by simply fusing them with the acid or its anhydride. Although boric acid or borates form insoluble precipitates with many metals, yet none of them are of much analytical importance. With barium salts borates give a white precipitate soluble in acids and in ammonium salts. Silver nitrate in concentrated solutions gives a white precipitate of silver borate, in dilute solutions a precipitate of silver oxide. Both are soluble in nitric acid and ammonia. Turmeric Test. — A very delicate test for boric acid is made by dipping a piece of turmeric paper into a solution of a borate made slightly acid with hydrochloric acid ; the paper is then dried on the water-bath, when it assumes a characteristic red tint. If the paper colored in this way be moistened with alkali the color changes to black or dark green, but is restored again by addition of a little dilute hydrochloric acid. In making this test much free hydrochloric acid must not be present, as it gives a brownish- black color with turmeric paper. Ferric chloride 204 BOEIC OR BOEACIC ACIC. also gives a brownish-red color, and for this reason should be removed before making the test. Flame Tests. — If alcohol be poured over dry boric acid or a borate, in the latter case concen- trated sulphuric acid being added to liberate the boric acid, and the alcohol lighted, the Hame vrill be colored green by the boric acid. Frequently, the color is only seen on the edges of the flame. The test is made more delicate by blowing out the flame after allowing it to burn for a few moments, and then relighting it, when the edges of the flame will be tinged green — only momen- tarily if the quantity of boracic acid is small. Chlorides and copper salts interfere with this test, as they give the same color to the flame. Copper may be i-emoved by precipitating with HjS and the chlorides by silver sulphate. The same coloration may be obtained by making a borate into a paste with a drop or two of concentrated sulphuric acid, heating the mix- ture gently in the Bunsen flame on the loop of a platinum wire, so as to expel excess of sulphuric acid, and then dipping the bead into glycerine. It is then lighted by holding in the flame for a moment, and on removing from the same it con- tinues to burn with a green flame. FLUORINE AND HYDROFLTJOEIC ACID. 205 FLUORINE (F) AND HYDROFLUORIC ACID (HF). Fluorine is exceedingly difficult to prepare in the free state, owing to the ease Math which it combines with almost all substances to form fluorides. It decomposes water instantly, forming hydro- fluoric acid and ozone. It also attacks glass or other silicates readily, forming silicon fluoride (SiF,). Fluorine is always found in nature as fluoride, chiefly as fluor-spar (CaFg) ; it is also contained in cryolite (NaAlF4), and in small amount in apatite (Ca3(P04)2). Hydrofluoric Acid. — In the anhydrous state this is a colorless, fuming liquid, boiling at 20° C. and readily soluble in water. It differs from all other acids in its property of dissolving silica, and for this reason is especially valuable in the analysis of silicates.- All metals with the excep- tion of gold and platinum are soluble in it with evolution of hydrogen. The fluorides of the alkalies, and a few others are soluble in water. Strong sulphuric acid liber- ates hydrofluoric acid from fluorides. 206 FLUORINE AND HYDEOFLTJOKIO ACID. Barium Fluoride. — Barium chloride gives with aqueous solutions of hydrofluoric acid or of fluo- rides a white precipitate of barium fluoride, very- insoluble in water, and difficultly soluble in cold hydrochloric or nitric acid. From these acid solutions it is not reprecipi- tated on the addition of ammonia, owing to the solubility of barium fluoride in ammonia salts : 2NaF + BaClg = BaF^ + 2NaCl. Calcium Fluoride. — Calcium chloride added to aqueoxis solutions of hydrofluoric acid or fluorides gives a gelatinous precipitate of calcium fluoride of much the same solubility as barium fluoride. Addition of ammonia helps the precipitation. Ammonium salts prevent it. Etching Test. — The most characteristic test for hydrofluoric acid depends on its property of com- bining with the silica of silicates, with the forma- tion of volatile silicon fluoride. If a finely pul- verized fluoride is placed in a lead or platinum dish, and made into a thin paste with concentrated sulphuric acid, and then gently wai'med, hydro- fluoric acid is evolved. A watch-glass covered with a thin coating of wax or paraffin, in which some letter or figure has been traced with a hard FLUORINE AND HYDEOFLUOEIC ACID. 207 wooden point, is placed closely over tlie dish, and the latter gently warmed for half an hour. (Care must be taken not to let the heat get high enough to melt the wax, for if this happens, of course, the etching is prevented.) At the end of this time the glass is taken off and the wax removed by warming, and wiping with a cloth ; the figure made in the wax will be found etched into the glass. CaF, + H2SO4 = CaSO, + 2HF ; SiO^ (of the glass) + 4HF = SiF^ + SH^O The presence of silica or a silicate in the flu- oride to be tested prevents or impairs this reac- tion, since the hydrofluoric acid libei'ated com- bines with that in preference to the silica of the glass. In this case the test is made as follows : Place the finely pulverized fluoride containing silica or silicate in a dry test-tube. Add concen- trated sulphuric acid in sufiicient quantity to cover the fluoride and heat gently. If now a drop of water in a loop of platinum be held in the mouth of the tube the drop will become opaque and solid, due to the silicon fluo- ride being decomposed by the water into hydror fluosilicic acid (HgSiFg) and silica (SiOg). 208 FLUORINE AND HTDKOFLTJORIC ACID. Instead of using the droj) of water the evolved gas may be passed over into water by fitting the test-tube with a perforated cork and bent-glass tube. When making the test in this way care must be taken that the end of the tube that dips into the water does not become choked by the separated silica, as this might cause an explosion ; also, do not allow the water to dra\v back into the tube containing the sulphuric acid for the same reason. The water into which the gas has been passed will be found to have an acid reaction, due to the hydrofluosilicic acid formed. The first re- actions are the same as in the previous test, but the silicon fluoride formed in this last one is de- composed by the water as follows : 3SiF4+ 4H2O = 2H3SiFe + H4SiO,. Silicates not decomposable by sulphuric acid should first be fused with sodium potassium car- bonate, the fusion dissolved in water and filtered, the filtrate made very sligJitJy acid with hydro- chloric acid, and boiled to expel carbonic acid. After cooling it is made alkaline with ammonia and calcium chloride added, and any precipitate that forms is dried and tested by one of the methods just given. SILICIC ACID, 209 SILICIC ACID (H,SiO„H,SiO,). Silicon forms but one oxide, SiO^, called silica or silicic anhydride. This oxide is found very abundantly in the mineral kingdom, as quartz, rock crystal, etc., and in the form of silicate it forms the greater part of the earth's crust. It is insoluble in water and acids, with the exception of hydrofluoric. Under certain conditions it unites with water to form hydrates or acids as H4Si04,H2Si03. Amorphous silica and the hydrated acids dis- solve in hot sodium or potassium hydrates, and in their carbonates, forming silicates. Also some- what soluble in ammonia, but not in ammonia carbonate. Fusion with alkali or alkaline carbonate yields a silicate soluble in water. Silica is not affected by heat even at extremely high temperatures. Of the silicates only those of the alkalies are soluble in water. Many of those insoluble in water are decomposed by acids, while others re- sist their action entirely, and can only be gotten into solution by previous fusion with alkaline carbonate, whereby the silicate is decomposed and 210 SILICIC ACID. the silica unites with the alkali to form a soluble silicate. Silicic Acid. — Solutions of alkaline silicates are easily decomposed by all acids. If excess of hydrochloric acid be added to a solution of sodium or potassium silicate, no pre- cipitate forms at first, but after standing for some time the whole mass solidifies as a transparent Jelly. If, however, the acid be added drop by drop,' silicic acid is precipitated in a gelatinous form. By adding hydrochloric or nitric acid in excess to a solution of a silicate and evaporating the solu- tion to complete dryness (110°-115° C), the silicic acid separates as anhydride (SiOg). By treating the dry residue with hydrochloric acid and water the insoluble silica is left as a residue. The last reaction is ver}'' generally made use of for separat- ing silica from its combinations. With silicates decomposed by acid, all that is necessary is to first boil the very finely pulverized mineral with an excess of strong hydrochloric acid, during which operation the geater part of the silicic acid sepa- rates, generally as a gelatinous precipitate, some- times as a pulverulent one. As soon as the mineral is completely decom- SILICIC ACIU. 211 posed the mixture is evaporated to dryness and heated in the air-bath, with occasional stirring, until no odor of acid is perceptible. The dry residue is then moistened with a little concen- trated hydrochloric acid, a small quantity of water added, and the mixture heated until everything except the silica has gone into solution. The silica is then filtered off, and on drying appears as a white gritty powder. In the case of silicates not decomposable by acid, the mineral, as already stated, must first be fused in a platinum crucible with a mixture of four or five parts of sodium potassium carbon- ate. (The fusion should be continued until the mass is perfectly liquid, and no more gas is given off.) The fused mass is then treated with water, and, without previous filtration, sufficient hydro- chloric or nitric acid is added to make the solu- tion acid ; the mixture is then evaporated to dry- ness and the silica separated as in the previous method. When it is required to test for the alkalies in a silicate that is not decomposed by acid, the fusion method just given is, of course, not admissible. In such cases the mineral is fused with barium or calcium oxides instead of sodium or potassium 212 SILICIC ACID. carbonate, and the silica separated in the same way as if the carbonates had been used. Silicon Fluoride. — Hydrofluoric acid as gas, or in strong aqueous solution, converts silica into sil- icon fluoride, SiFi (see hydrofluoric acid); in dilute solution it dissolves silica with formation of hydrofluosilicic acid : SiOa + 6HF = H^SiFe + 2H,0. Silicates when treated with hydrofluoric acid yield silicofluorides (a), but if sulphuric acid be present the silicofluoride is decomposed into sul- phate, silicon fluoride, and hydrofluoric acid. {a) CaSiOa + 6HF = CaSiF6+ SH^O ; (5) CaSiFe + H^SO, = CaSO^ + 81^4 + 2HF- Ammonium or calcium fluoride (of course free from silica) may be used in place of hydi'ofluoric acid. To a mixture of three or four parts of the fluoride and one of the silicate, in a platinum ci'u- cible or dish, concentrated sulphuric acid is added and the mixture warmed. If a drop of water in the loop of a platinum wire be now held over the dish, the water will become cloudy or opaque, due to the separated silicic acid (see hydrofluoric acid). CARBONIC ACID. 213 This metliod of decomposing a silicate may be used to advantage when it is required to test for the alkalies in silicates undecomposable by other acids. Blowpipe Test.— It silica or a silicate be fused in a salt of phosphorus bead (NaPOj), the bases with which the silica was combined dissolve, while the silica floats about in the bead in the same form as the piece of mineral added, yielding the so-called silica skeleton. This test is not alto- gether reliable, as certain silicates such as the zeo- lites dissolve completely, while- some other sub- stances containing no silica, as apatite, are not acted on. CARBONIC ACID (Anhydride, CO,). Carbon is found in many different forms — such as the diamond (crystallized carbon), which is ex- ceedingly hard and transparent ; as graphite, which is soft, opaque, and black ; all the different kinds of coal, charcoal, etc. It is insoluble in acids, and infusible. Some varieties are easily combustible ; others, as the diamond and graphite, are burnt with difficulty. When burned with an excess of oxygen or air it yields carbon dioxide or anhydride (COg) 214 CAEBONIC ACID. (commonly called carbonic acid) ; if the, carbon is in excess the product is carbon monoxide (CO). Garhonic Anhydride (^00^. — Carbon dioxide or anhydride is at ordinary temperatures and pressures a colorless and almost odorless gas of high specific gravity (1.52). At 0° C. and a pres- sure of 36 atmospheres it condenses to a liquid, which if allowed to evaporate rapidly causes a fall in temperature to — 80° C, and a portion of the liquid solidifies as a snow-white mass. The gaseous dioxide is somewhat soluble in water, one volume of cold water dissolving about one volume of the gas. Air containing carbon dioxide even in comparatively small amount is unfit for respiration, and many a fatal accident has been caused by inhaling it. The dreaded " after-damp" of coal-mines consists largely of this gas. It will not support combustion, a lighted candle being instantly extinguished if plunged into a vessel filled with it. This may be made use of as a rough test to detennine whether the air in any confined space, as a shaft or well, con- tains much carbon dioxide. If on lowering a lighted candle into a shaft the light is extinguished, it is proof that the air. contains the gas in considerable quantity. A CARBONIC ACID. 215 mixture of air and gas containing only 25^ of COj will extinguish a burning candle. Carbon dioxide is a very weak acid. It turns blue litmus red, but on drying the blue color reappears. It forms carbonates with many of the metals, and of these only the carbonates of the alkalies are soluble in water. All carbonates, however, are soluble in acids, even in very weak ones, except in hydrosulphuric and hydrocyanic. The carbonates are very widely distributed in nature, as limestones, marbles, chalk, dolomite, etc., also in many mineral springs, and in almost all waters to a greater or less extent. The at- mosphere also contains a certain amount of carbon dioxide, which although small in amount (0.05^ to 0.1^) is a very important constituent, since it supplies plants with their carbon, their leaves with the aid of sunlight being able to decompose the carbon dioxide, assimilating the carbon and liberating oxygen. As already stated, carbonates are soluble in almost all acids, most of them dissolving even in the cold ; a few, however, require heat (magne- site, siderite, etc.). The decomposition is accom- panied by effervescence, caused by the escaping carbon dioxide (a). The acid used to decompose 216 CARBONIC ACID. the carbonate stould always be added in excess so as to prevent the formation of any acid carbon- ate (^), since carbonic acid forms acid carbonates with a few of the metals, particularly those of the alkalies (KRCO^'NailCO,, etc.). To prove that the gas evolved is really carbon dioxide it is passed into a solution of Kme or baryta water, or a glass rod that has been dipped in baryta-water is held in the test-tube in contact with the gas. If carbon dioxide is present the lime or baryta water becomes turbid, ovring to the formation of a precipitate of carbonate (c). If an excess of the gas is passed into the solu- tion the precipitate will dissolve, but on boiling it again precipitates, the excess of the dioxide be- ing driven off by the heat. (a) NagCOs + 2HC1 ^Cb,-\- 2NaCl + H^O ; (b) Na^COg + HCl = NaHCOs + NaCl ; (c) Ca(0H)2 + CO, = CaCOa + H^O. Barium or Calcium Carbonate. — With solu- tions of the normal carbonates barium and calcium chlorides give white precipitates of carbonate. CAEBONIC ACID. 217 ' In dilute solutions of acid carbonates tte pre- cipitate only forms on boiling. With solutions of carbonic acid no precipitate is formed unless the solution be neutralized : Na^COg + BaClg = BaCOs + 2NaCl. Lead Ca/rhonate. — An ammoniacal solution of lead acetate or an aqueoils solution of lead chloride gives in solutions of normal or acid carbonates a white precipitate of carbonate. All neutral solu- tions of lead salts are precipitated by the normal carbonates. ACIDS OF THE SECOND GEOUR The characteristics of this group are the precip- itation by silver nitrate from solutions acid with nitric acid, and the fact that barium chloride does not precipitate them. The acids are hydrochloric, hydrohromic, hydri- odic, hydrocyanic, hyd/ivferrocyanic, hydroferri- cycmic, and hydrosulphurie. HYDROCHLORIC ACID (HCI). Chlorine is a neavy, yellowish-green gas, pos- sessing a characteristic odor, and is extremely irritating to the respiratory organs. It is soluble- in water, forming a solution of a faint yellowish- green color. This solution on exposure to light decomposes into hydrochloric acid and oxy- gen (a). Chlorine, both as a gas and in solution, acts as a strong bleaching agent, destroying many vege- table colors, and in solution is a very powerful 818 HYDROCHLORIC ACID. 219 oxidizing agent. It combines freely with the alkaline hydrates, forming hypochlorite (b) in cold solution and chlorate in hot (c). (a) 2H,0 + 2C1, = 4HC1 + O, ; (b) 2K0H + CI, = KCIO + KCl + HgO ; (c) 6KOH + 3Cl3=KC103 + 5KCl + 3HA Hydrochloric acid at ordinary temperatures is a colorless, suffocating gas, fuming in moist air, and very soluble in water. Concentrated solutions lose a considerable amount of the gas on heating. The acid is easily formed from chlorides by heat- ing them with sulphuric acid. It is made com- mercially from common salt in this way : 2NaCl + H3SO4 = NagSO^ + 2HC]. The normal chloiides are all easily soluble in water, with the exception of silver, lead, mercurous and cuprous chlorides. Bismuth and antimony chlorides require the presence of some free acid to hold them in solution. Chlorine occurs in nature principally as sodium chloride (common salt), either in the form of rock salt or, dissolved in water, as sea- water. It is also found to some extent as potassium and magnesium chlorides (Stassfurth deposits). 220 HYDROCHLORIC ACID. Silver Cliloride. — Hydrochloric acid and almost all chlorides, even in exceedingly weak solutions, give with silver nitrate a white precipitate of silver chloride, which changes to a violet or black color on exposure to light. For solubilities, etc., of precipitate, see silver chloride, under head of Silver. Mercurous and Lead CMorides. — Solutions of either mercurous or lead salts are precipitated by hydrochloric acid or soluble chlorides, as already given under those metals. Evolution of Chlorine. — If hydrochloric acid be heated with manganese dioxide or lead dioxide, chlorine is evolved (a), which is easily recognized by its yellowish-green color, its odor, and bleach- ing action on vegetable colors. If a piece of moist litmus paper is held in the fumes the color will be destroyed. In the case of chlorides some strong sulphuric acid must be added as well as the dioxide (b). Chlorides heated with manganese or lead diox- ides and acetic acid do not evolve chlorine. {a) MnOa + 4HC1 = 018 + MuCl^ + 2H2O • {V) 2NaCl + Mn08 + 3H2S04 = CI2 + 2NaHS04 + MnSO^ + 2H3O. HTDEOCHLOEIC ACID. 221 Free hydrochloric acid alone or in presence of chlorides is easily detected by simply heating the solution with manganese dioxide, when chlorine is evolved as in reaction (a). Chromio Oxy chloride {CrO^Gl^. — If a dry chloride be mixed with potassium dichromate and placed in a small flask or test-tube, and some con- centrated sulphuric acid added and the mixture gently heated, chromic oxy chloride or chromyl chloride (also called chloro-chromic acid) will be liberated as a reddish-brown gas that condenses on cooling to a red liquid (a). If the flask or test- tube be fitted with a delivery-tube, and the gas passed into ammonia, a yellow solution of ammo- nium chromate is obtained (h). On acidifying this solution with acetic acid, and adding a drop or two of lead acetate, a yellow precipitate of lead chromate is formed. (a) KA207 + 4NaCl + 6HaS04 = 2CrOici3+ 2KHSO4 + 4]S'aHS04 + SH^O ; (5) Cr02Cl8 + 4NH40H = (NH,),Cr04 + 2NH4CI + 2H2O. This test is of particular value for determining chlorides in the presence of bromides, the latter 222 HTDBOBKOMIO ACID. forming no corresponding compound, but liberat- ing bromine instead : 6KBr + KgCr^O^ + 1 iH^SOi = sS-s + 8KHSO4 + Cr2(S04)3 + 7H2O. The bromine when passed into ammonia gives a colorless solution. Iodides with dichromate and sulphuric acid act in the same way as bromides. Insoluble Chlorides. — For the detection of chlorine in the chlorides insoluble in water and nitric acid, they should be fused with sodium potassium carbonate, the fusion extracted with water, and the chloride tested for in the filtrate. Blowpipe Test. — If a salt of phosphorus bead be saturated with cupric oxide, and then a small quantity of any substance containing a chloride be added, and the substance heated in the redu- cing flame, the flame will be colored blue. HYDROBROMIC ACID (HBr). Bromine is a heavy, dark-red liquid which boils at 63° C. Even at ordinary temperatures it vola- tilizes freely, giving off heavy red fumes that are exceedingly irritating. It is more soluble in water than chlorine (a sat- urated solution contains about 3^ of bromine), HYDROBROMIO ACID. 223 and easily soluble in ether, chloroform, or carbon disulphide. If an aqueous solution of bromine be shaken with one of the substances just mentioned they will extract the bromine from the water solu- tion, the latter becoming colorless, and the ether or disulphide appears of a red or yellow color. Bromine, like chlorine, is a strong bleaching and oxidizing agent, and the reactions are similar. Hydrobromic acid at ordinary temperatures is a colorless gas, very soluble in water. The acid itself or -its aqueous solution decom- poses slowly in presence of air, bromine being lib- erated, which colors the solution red or yellow. Bromides are only found in small quantities. They are contained in sea- water, and in many salt springs. In chemical properties the bromides resemble the corresponding chlorides very closely. Silver Bromide. — As already stated under the reactions for silver, bromides or hydrobromic acid precipitate silver salts as bromide, insoluble in nitric acid. For further properties of the precipitate, see silver bromide, under head of "Silver." Liberation of Free Bromine. — Hydrobromic acid and all bromides except silver bromide are 224 HYDROBEOMIC ACID. decomposed by heating with nitric acid, bromine being set free ; 6KBr + 8HNO3 = 3Br2+ 6KNO3 + 2N0 + m,0. Chlorine either as gas or in solution decomposes bromides, liberating bromine. Alkaline hypochlorites as sodium hypochlorite (NaClO) effect the same reaction if added to solu- tions slightly acid with hydrochloric acid. The addition of either reagent must be made with care, as an excess of it converts the bromine into color- less bromine chloride. If before liberating the bromine a few drops of carbon disulphide or chloroform be added to the solution, and then the chlorine water or hypochlorite and the mixture be gently shaken, the bromine is all taken up' by the disulphide or chloroform, which now shows a red or yellow color. It is always advisable to test the reagents by making a blank test (i. e., making a test in the same way, except that pure water is substituted for the solution to be tested for bromine), .to see if they give any color with the carbon disulphide. Should any color appear, either new reagents must be used or the color obtained in the blank test should be compared with that given by the substance under examination. HYDROBROMIC ACID. 225 Bromides with manganese or lead dioxides and sulphuric acid give the same reactions as chlorides, the only difference being that a more dilute solu- tion of sulphuric acid can be used than with the latter. With manganese dioxide and acetic acid bro- mides are not decomposed, but with lead dioxide and acetic acid bromine is liberated (a distinction from chlorides). Action on Starch. — If a bromide be mixed with manganese dioxide and sulphuric acid in a small beaker, and the latter covered with a watch-glass, having a strip of paper attached to the under sidp which has been moistened with starch-paste and sprinkled with starch powder, on heating the mix- ture gently the liberated bromine colors the starch- paste yellow. Insoluble Bromides. — Bromides insoluble in nitric acid are fused with sodium potassium car- bonate, and the fused mass treated in the same way as for chlorides. Blowpipe Test. — Bromides when heated in re- ducing flame in a salt of phosphorus bead which has been saturated Avith cupric give a blue flame inclining to green. 226 HYDKIODIC ACID. HYDRIODIC ACID (HI). Iodine generally occurs in the form of black crystalline scales having a decided lustre. It fuses at 115° C, and boils at 200° C. (Stas), with evolution of violet fumes vrhich con- dense in a crystalline form. It is only very slightly soluble in water, but dissolves easily in alcohol, ether, chloroform, carbon disul- phide, and in aqueous solution of iodides. The solutions in chloroform and carbon disulphide are violet, the others brown or reddish-brown. , Certain reducing agents as thiosulphates (a) and sulphites (b), dissolve free iodine, yielding colorless solutions. 21 + 2Na2S203 = 2NaI + Na2S406 (sodium tetra- thionate); 21 + Na^SOs + HgO = 2HI + Na^SO^. Starch-paste, even with minute traces of iodine, gives a beautiful blue color. Iodine has only very weak bleaching properties ; the action on litmus and other vegetable colors is very slow. It stains the skin brown. Hydriodic acid is a colorless gas at ordinaiy temperatures, and is very soluble in water, giving HYDRIOBIC ACID, 227 a colorless solution. This solution soon turns brown, owing to the gradual decomposition of the hydriodic acid, iodine being liberated and dis- solving in the remaining acid. Iodides. — Iodides in many ways correspond to the chlorides and bromides, but more of them are insoluble in water, and generally these insoluble compounds are soluble in excess of the iodide, forming double iodides. Silver Iodide. — Soluble iodides give with solu- tion of silver nitrate a yellow precipitate of silver iodide. The precipitate is soluble in excess of the iodide, forming a double iodide (KI.AgI). It is insoluble in dilute nitric acid, and almost insoluble in dilute ammonia. Soluble in thiosulphate and potassium cyanide. Mercurous cmd Lead Iodides. — These have already been described under their respective metals. Oaprous Iodide. — Sulphate of copper precipi- tates soluble iodides as white cuprous iodide, free iodine being liberated at the same time (a). In presence of reducing agents, as ferrous sulphate {b) or sulphurous acid (c), no free iodine is formed. Addition of ammonia assists the precipi- tation. 228 HYDRIODIC ACID. (a) 4KI + 2CUSO4 == CuJa + 12 + 2K2SO, ; (b) 2KI+2CuS04 + 2FeS04 = CuJ, + Fe,(S0,)3 + K,S0,; (c) 4KI + 2CUSO4 + SO, + 2H3O = CuJa + 2K,S0, + H2SO4 + 2HI. Chlorides and bromides are not precipitated by these reagents. Liberation of Free Iodine. — The most delicate and characteristic tests for iodides depend on the liberation of free iodine, the presence of the iodine being then made evident by suitable means. Iodides are decomposed by many different re- agents, iodine being set free in the reaction, such as concentrated sulphuric or nitric acids, nitrous acid even in dilute solution, chlorine, bromine, etc. If a solution containing an iodide, even in ex- tremely minute quantity, be acidified with a few drops of dilute sulphuric acid, then some starch- paste* added, and finally a couple of drops of * Starch-paste is made by mixing one part of starch with about 100 parts of cold water and then heating until a solution is ob- tained that is almost clear. This solution should be freshly pre- pared for each experiment. HYDRIODIC ACID. 229 potassium nitrite, the solution will become blue, or if only traces of iodine are present, a reddish- violet. The test should always be made in a cold solution, as heat destroys the color. . Chlorine may also be used for this test; but is not so well suited for it, as an excess bleaches the color. Another very delicate test for iodine is made by adding to the solution to be tested a few di'ops of chloroform or carbon disulphfde, and then very carefully, drop at a time, a solution of chlorine water or a hypochlorite ; on shaking the mixture the liberated iodine dissolves in the chlo- roform or earbon disulphide with a beautiful violet or purple color. Other reagents that liberate iodine may be use'd for this test, such as nitrite and sulphuric acid. When chlorine or hypochlorite is used the reagent must be added carefully, since an excess of the reagent forms colorless iodine chloride. If iodides are heated with ferric sulphate or chloride, iodine is set free : 2KI + Fe^CSO^)^ = I^+ 2FeS04 + K^SO, . This maybe used as a separation from chlo- rides and bromides. With manganese or lead dioxides and dilute 230 HYDROCYANIC ACID. sulphuric or acetic acids, and with potassium di- chromate and dilute sulphuric acids, iodine is set free, and, unless in very small quantity, may be recognized by its violet color. Insoluble Iodides. — Iodides insoluble in acid are fused with sodium potassium carbonate and the fusion treated as already given for bromides or chlorides. Blowpipe Test. — Iodides when fused in the reducing flame in a salt of phosphorus bead satu- rated with cupric oxide give a green flame. HYDROCYANIC ACID (HCN or HCy). Cyanogen (C2N3) is a colorless gas, having a characteristic odor, and when lighted burns with a violet flame. It may be prepared by the igni- tion of silver or mercuric cyanide. Hydrocyanic acid is a colorless, volatile liquid, boiling at 27° C, and having a very characteristic- odor, resembling that of bitter almonds. It is soluble in water and alcohol in all propor- tions, but these solutions soon undergo decom- position if concentrated ; in dilute solution it is more stable. Hydrocyanic or prussic acid is one of the most poisonous substances known. The soluble cya- HYDEOCTANIO ACID. 231 nides are also exceedingly poisonous, and great care must be taken in handling them. The cya- nides of the alkalies and alkaline earths are solu- ble in water, and the solutions obtained are decomposed by acids, even by carbonic acid. AH cyanides are decomposed by evaporating with concentrated sulphuric acid. Potassium and sodium cyanides are valuable reducing agents. They can be fused without de- composition if air be excluded ; but when fused with substances capable of yielding oxygen, such as metallic oxides, they reduce these and are con- verted into cyanate (a). Sulphides are reduced in the same way, yielding thiocyanate (5) : SnOa + 2KCN = Sn + 2KCN0 ; PbS + KCN = Pb + KCNS. The cyanides of the heavy metals are all decom- posed on ignition, yielding either the metal or a carbide of the metal. Cyanides combine with the bases to form dif- ferent classes of compounds — as the simple cya- nides (as KCN), double cyanides (as AgCN,KCN), and compound cyanides (as K4FeCy6 or KeCo^Cyia). The simple cyanides (a) have a great tendency to combine and form soluble double cyanides (h), 232 HTDEOCTANIC ACID. from which solutions they are reprecipitated on acidifying with acid (c) : {a) Co(N03)2 + 2KCy = CoCy^ + 2KNO3 ; (b) CoCy, + 4KCy = CoC~^4KCy ; ^ (c) CoCy5,,4KCy + 4HC1 = CoCy^ + 4KC] + 4HCy. In a compound cyanide the cyanogen is com- bined with the metal in such a way that the two act as if they were a simple element (such a com- bination is called a compound radical), and the metal no longer acts as it does in simple solutions ; for example, the iron in potassium ferrocyanide no longer gives the ordinary reactions of that metal, and, indeed, we use this very solution as a reagent to test for iron. These compound cyanides on addition of dilute acids give no precipitate of a cyanide, as the double cyanides do. On evaporation with strong sul- phuric acid, however, they are decomposed into hydrocyanic acid and metallic sulphatp, which re- mains as a residue. Simple Cyanides. — Silver nitrate added to a solution of hydrocyanic acid or a soluble cyanide gives a white precipitate of silver cyanide (a) which is easily soluble in potassium cyanide, form- HYDROCYANIC ACID. 233 ing double cyanide (b) ; it is also soluble in am- monia; insoluble in dilute nitric acid. On boil- ing with strong hydrochloric acid it is decomposed into silver chloride and hydrocyanic acid : (a) KCN + AgNOg-AgCN + KNOa; (h) AgCN + KCN = aJcN,KCN. Silver cyanide on ignition is decomposed, yield- ing metallic silver, cyanogen, and some silver para- cyanide. Formation of Ferrocyanide. — If to a solution containing hydrocyanic acid or a cyanide a mix- ture of ferrous and ferric (a drop or two only of the ferric is required) salt be added, then suffi- cient sodium or potassium hydrate to make the solution alkaline, and the mixture gently heated, the cyanide is changed to ferrocyanide, and on acidifying the solution a blue precipitate of Prus- sian blue is obtained ; 6KCN + FeSO, = K4Fe(CN)6 + K^SO^. ' Formation of Thiocyanogen {Stibphocyanogen). — Another very delicate test for hydrocyanic acid or cyanides is made by adding to the solution a few drops of ammonia and a little strong ammo- nium sulphide (polysulphide) and evaporating to 234 HYDROCYANIC ACID. dryness at a gentle heat (best on water-batL). The residue, which now contains thiocyanate, is made acid with hydrochloric acid and a drop of ferric chloride added, when the well-known blood- red color of ferric thiocyanate appears : KCN + (NIl,)S. = KCNS + (NH,),S,_j. The evaporation to dryness decomposes the ex- cess of ammonium sulphide, so that the residue is free from that compound. When making this test it is well to add the hydrochloric acid in moderate excess, otherwise acetates, phosphates, etc., if pres- ent, will interfere with the reaction. Hydrof&i'rocyanic Acid {H^Fe{CN')^. — This acid is a colorless, crystalline solid, easily soluble in water. The ferrocyanides of the alkalies and alkaline earths are soluble in water, most of the others insoluble. The insoluble compounds are easily brought into solutioji by boiling with sodium or potassium hydrates, the metal with which they were com- bined, if its hydi'ate is insoluble in the alkali, being left as a residue. Fe4(Fe(GN)6)3+12KOH = 2Fe,(OH)o + 3K4Fe(CN)6. HYDROCYANIC ACID. 235 If the metallic hydroxide is soluble in alkali (as Zn(0H).3,Pb(0H)a, etc.) it also goes into solu- tion, and must be precipitated from the alkaline liquid b}'^ HgS. All ferrocyanides, as already stated, are decomposed by evaporation with strong sulphuric acid. Fusion with sodium or potassium nitrate effects the same result. Silver Ferrocyatiide. — ^Nitrate of silver gives with ferrocyanides a white precipitate of silver ferrocyanide (Ag4Fe(CN)6), insoluble in nitric acid and ammonia, soluble in potassium cyanide. The most characteristic test for ferrocyanides is made by adding to the slightly acid solution a few drops of a ferric salt, when, if ferrocyanides are present, a deep-blue precipitate (Prussian blue) will be formed. Cupric salts also give a characteristic red pre- cipitate. Hychoferricyanic Acid (^H^Fe.2,{CN)'i^. — Many of the ferricyanides are soluble in water, and all are decomposed in the same manner as the ferrocyanides. Nitrate of silver gives an orange-colored pre- cipitate of silver ferricyauide (Ag6Fea(CN),2), insoluble in nitric acid, soluble in ammonia, and potassium cyanide. 236 HYDKOSULPHDIUO ACID, Fei'ric salts give no precipitate, but ferrous give a deep-blue precipitate of ferrous ferricyan- ide (Fe3Fea(CN),2). This test is characteiistic and exceedingly delicate. Thwcyanic Acid (^Hydrostdphocyanic). — Thio- cyanic acid or ttiocyanates (sulphocyanides) are readily detected by the blood-red color given by ferric salts in hydrochloric-acid solution. HYDROSULPHURIC ACID (H,S). Sulphur at common temperatures is a solid body, generally of a yellow color, but precipitated sulphur is white. It is insoluble in water and alcohol, soluble in benzol, petroleum, and car- bon disulphide, the latter substance being its best solvent. Amorphous sulphur is generally insoluble in the disulphide. There is also a col- loidal modification that is soluble in water. Sul- phur fuses at 115° C. and boils at 444° C. When heated in the air it burns with a bluish flame to sulphur dioxide, the latter being recog- nized by its well-known and characteristic odor. Strong oxidizing agents, as concentrated nitric acid, aqua regia, hydrochloric acid, and potassium chlorate, etc., with the aid of heat, dissolve sulphur gradually, forming sulphuric acid. Potassium and HYDEOSULPHUBIC ACID. 237 sodium hydrates dissolve it, on beating, to a yellow liquid containing sulphide and thiosulphate. Sulphur is often found in the free state in connec- tion with deposits of gypsum and rock salt, and in the vicinity of volcanic craters. Sicily furnishes a large proportion of the world's supply. Free sulphur in a substance can generally be easily detected by igniting it in the air, when it burns with a blue flame and the characteristic odor of sulphur dioxide. The test is best made by heating the mixture in a glass tube open at both ends, the tube being held in a slanting posi- tion and the heat applied to the part where the substance rests. If heated in a tube closed at one end the sul- phur volatilizes as yellowish-brown vapors that condense in the upper, cold portion of the tube to brown drops, turning yellow on cooling. HydrosulpTiuric Acid or Hydrogen Sulphide at common temperatures and pressure is a color- less, poisonous gas, having the odor of rotten eggs. It is easily inflammable, burning to sulphur diox- ide and water. It is soluble in water, but the solution is not stable. For this reason, when used as a precipi- tating agent it should be freshly made. It is 238 HYDEOSULPHURIC ACID. generally obtained by the action of acids ou metallic sulphides (usually ferrous sulphide). FeS + 2HC1 = FeCIa + hJ. Sulpliides. — The sulphides of the heavy metals are found very widely distributed in nature, aud form many of the most valuable ores. The sul- phides of the alkalies and alkaline earths are the only ones soluble in vpater. These and the sul- phides of iron, zinc, aud manganese are decom- posed by cold dilute minei'al acids, with evolution of hydrosulphuric acid. Of the other sulphides, some are soluble in hot, strong hydrochloric acid ; others require nitric acid or aqua regia (i.e., chlorine). When nitric acid or mixtures yielding chlorine are used, hydrosulphuric ^.cid is not liberated, but sulphur is set free {a and V), and frequently the action goes further, the sulphur being oxidized to sulphuric acid (c and d) : {a) 8Ag,S + 8HN03 = eAgNOg + 3S -\-ARf> + 2N0 ; {h) HgS + CL = HgCl, + S_; (c) 3PbS -i 8liNO,(hot and cone.) = SPbSO^ + 8N0 + 4H,0 ; HYDROSXJLPHURIC ACID. 239 {d) S + 3C1, + 4H3O = H2SO4 + 6HCL Detection of HydrosulpTiuriG Acid. — When in the gaseous state or in solution this acid is easily recognized by its odor, or by bringing the acid in contact with lead or silver salts, when black pre- cipitates of sulphide are formed. A convenient way of making the test is to hold in the gas a piece of filter-paper that has been moistened with lead acetate and then made alkaline by the addition of a drop or two of ammonia; if HjS is present, even in minute quantity, the paper shows a brown or black lustrous film of lead sulphide. Sulphides decomposed by hydro- chloric acid are tested in the same manner, by placing them in a test-tube with the acid and heating, a piece of paper moistened with lead acetate and ammonia being held at the mouth of ' the tube. Sulphides decomposed with difiaculty by hydro- chloric acid, if treated in a small flask or test-tube with hydrochloric acid and finely divided iron free from sulphur (ferrum alcoholisatum), evolve HaS along with the hydrogen liberated by the action of the acid on the iron. If the flask or tube has been loosely corked, and a piece of acetate of lead paper fastened to the under side of the cork, 240 HYDKOSULPHUKIC ACID. the paper will be turned black or brown by the sulphide. Realgar, orpiment, and molybdenite are not decomposed. To detect sulphide in presence of free alkali or alkaline carbonate, a solution of lead hydroxide in sodium or potassium hydrate (made by adding the alkaline hydrate in excess to a soluble lead salt) is added, when even a ti-ace of sulphide gives a black or brown precipitate ; or if the amount is very minute, simply a brown color (1 part of lead in 1,000,000 parts of water may be detected in this way). Sodium nitroprusside (Na,Fe,(CN),o(NO)„2H,0) added to alkaline solutions is also a very delicate test fo)' sulphides, giving a reddish-violet tint to the solution. Sulphides insoluble in acids should be fused with sodium or potassium hydrate in a silver or nickel crucible. A soluble sulphide of the alkali is formed that can be tested by any of the methods just given, or, what is simpler, the fused mass is placed on a piece of bright silver and moistened \vith water, when the silver is turned black by the sulphide pi'esent. HYDKOSULPHURIC ACID. 241 Oxidation to Sulphate. If a sulphide be fused with sodium potassium carbonate and an oxidiz- ing agent, as sodium nitrate or potassium chlorate, the sulphide is oxidized to sulphate. The fused mass is then digested with water and filtered, and the filtrate tested for sulphuric acid. This test is, of course, only conclusive in the absence of sulphates and all other substances containing sulphur (except sulphides), for all on fusion in this way yield sulphates. Blowpipe Test. — If a sulphide be heated in a glass tube open at both ends, in the same way as given for free sulphur, sulphur dioxide will be given off, and may be recognized by its odor. Also, if heated in oxidizing flame of the blow- pipe, sulphides yield sulphur dioxide. THIRD GEOUP OF ACIDS. Acids which are not precipitated by either barium chloride or silver nitrate — viz., nitric acid and chloric acid. The chief characteristic of this group is that all the nitrates and chlorates are soluble in water, with the exception of a few basic nitrates. NITRIC ACID (HNO3). Nitrogen combines in five different proportions with oxygen— viz., N,0,— NjO, or NO,— N.Og,— N,04 or NO,. The N20,Na03 and NA unite with water to form the acids — hyponitrous (HNO), nitrous (HNO,), and nitric acid (HNO3). Nitric acid is the only one of common occur- rence, so we will only consider its reactions. The pentoxide (N^Og) is a colorless substance, crystallizing in six-sided prisms, which melts at 343 NITRIC ACID. 243 30° C. It dissolves in water, with evolution of much heat, forming nitric acid : NA + HgO^SHNOg. Nitric acid is a colorless (when free from lower oxides of nitrogen), fuming liquid, very corrosive, and acting powerfully on organic tissues. It boils at 86° C, and its specific gravity is 1.52. It is a strong oxidizing agent, and is a solvent for many metals, in the latter case generally giving off nitric oxide (NO). Nascent hydrogen in alkaline solution reduces nitric acid to ammonia: KNO3 + 8H = KOH + NH4OH + H2O. On ignition all nitrates are decomposed, leaving a residue of oxide of the metal. Commercially, nitric acid is made by heating sodium nitrate (Chili saltpetre) with strong sul- phuric acid : 2NaN03 + H2SO4 = Na^SO^ + 2HNO3. Since nitric acid forms no insoluble compounds with any of the metals, we are obliged to make use of other reactions for its detection. Its oxi- dation reactions afford an easy means of recogniz- ing it. 244 WITKIC ACID. Action on Copper. — If a nitrate be heated with metallic copper and concentrated sulphuric acid (the latter is added to liberate the nitiic acid) in a test-tube, the tube becomes filled with reddish- brown fumes of nitrogen peroxide, formed by the oxidation of the nitric oxide : 3Cu + 8HNO3 = 3Cu(N03)2 + 2N0 -f 4H2O ; 2N0 + O2 = 2NO2 (reddish-brown fumes). Action on Ferrous Salts. — If a solution of nitrate is mixed with an equal volume of concen- trated sulphuric acid, and after the mixture has been cooled a little feri'ous sulphate be poured very carefully down the side of the tube, so that the two solutions do not mix, a brown ring will foim at the junction of the two liquids. In mak- ing this test it is advisable to have the solutions concentrated. On heating, the brown color dis- appears : 2KNO3 + 4HaS04 + lOFeSO, =2(FeS04),NO + 3Fe,(S04)3 + K^SO^ + Alif>. (Brown ring ) It will be peen fioni the equation that a part of the ferrous salt is oxidized, and that another por- tion combines with the nitric oxide liberated NITRIC ACID. 245 by the reaction to form the brown compound (FeS04)2NO. It may perhaps make the action plainer to write the equation in two stages, thus : 2KNO3 + 6FeS04 + 4H,S04 = 3Fe,(S04)3 + 2N0 + K2SO4 + 4H,0 ; 4FeS04 + 2N0 = 2(FeS04)8NO. Action on Indigo. — A little hydrochloiic acid is boiled in a test-tube for a few moments, then a drop or two of indigo in sulphuric acid is added, and the solution again boiled ; the solution will re- main blue if the hydrochloric acid was free from chlorine. If a solution containing nitric acid or a nitrate is now added, and the solution again boiled, the blue color disappears. The action is due to the liberation of chlorine, which reacts on the indigo. For this reason, any substance yielding chlorine (as chlorate) will give the same reaction. In a negative way the test is very useful, for although the bleaching of the solution is not conclusive for nitrates, still if the color is not destroyed we know that nitrates (also chlorates) are not present. Action on PTienol. — A test that is well suited to the detection of nitrates when present only in minute quantities (as in well-water, etc.) is made 246 NITRIC ACID. by adding two or three drops of a mixture of one part phenol (carbolic acid), four parts of concen- trated sulphuric acid, and two parts of water to a solid nitrate. A reddish-brown color is produced, which changes to yellow on the addition of a few drops of ammonia. The color is due to the formation of a nitro compound with the phenol. (CeH2(N02)30H, picric acid.) When it is desired to test a solution for nitrate in this way the solution, which must be neutral or alkaline, is evaporated to dryness in a porcelain dish and the mixture of phenol, etc., added to the residue. Chlorates, if present, must be removed by ignition before making this test, as they give mucli the same color. Nitrates cause deflagration if heated to a dull red with charcoal or organic matter, such as paper. The deflagration is caused by the oxidation of the carbon by the oxygen of the nitrate. M'ee nitric acid may be detected by adding to the solution a few quill-cuttings and evaporating to dryness in a porcelain dish or beaker on the water-bath. If any free nitric acid is present the quills will be turned yellow. CHLORIC ACID. 247 CHLORIC ACID (HCIO,). Chlorine forms four acids with oxygen — viz., hypochlorous acid (HCIO), chlorous acid (HCIO2), chloric acid (HCIO3), and perchloric acid (HCIO4). Hypochlorous and chloric acids are frequently met with, the first in the form of bleaching-powder, also as Javelle water (NaClO), the latter as potassium chlorate. Hypochlorites act as strong bleaching agents, especially in acid solutions, due to the liberation of chlorine. Solutions of hypochlorites are decomposed on boiling into chloride and chlorate. Hypochlorites may be detected by their bleach- ing action on litmus paper, even in alkaline solutions, and by shaking their solutions with mercury the latter is changed to the yellow or red oxide. Chlorine under the same conditions gives mercurous chloride, white, and chlorates in neutral or alkaline solution do not attack mer- cury. With manganese salts hypochlorites give a brownish-black precipitate of hydrated manga- nese dioxide, and with silver nitrate they give' a pi'ecipitate of silver chloride. The chlorates are all soluble in water. On 248 CHLORIC ACID. ignition they yield oxygen, or, in some cases, oxygen and chlorine; in the first case the residue left is a chloride, in the latter an oxide. This action is employed in the production of oxygen from potassium chlorate : 2KC103 = 2KCl+302. Heated with organic matter, the chlorates cause deflagration acting much more strongly than the nitrates. A mixture of potassium chlorate and nitric acid forms a very powerful oxidizing agent. Hyd/rocMoric ticid decomposes chlorates, yield- ing chlorine and oxides of chlorine, the action varying with the conditions. If the chlorate present is not in too small quantity, the test-tube in which the experiment is made is filled with a yellowish-green gas of a very peculiar and dis- agreeable odor, and the solution as well becomes yellow. Solution of indigo is quickly bleached by it. Sulphuric acid also decomposes chlorates into perchloric acid and chlorine tetroxide : 6KCIO3 + 3H2SO4 = 2C1A + 2HC10, + 3K3SO, + 2HA If a few drops of concentrated sulphuric acid are placed in a small e\'aporating-dish, and a very CHLORIC ACID. 249 S7nall quantity of chlorate added to it, the sul- phuric acid assumes an intense yellow coloi", due to the chlorine tetroxide. No heat should be used in this test, and only a small amount of the chlorate, otherwise a danger- ous explosion Tusij result. If to the yellow solution Just mentioned a drop of an aqueous solution of aniline sulphate be added, and then a few drops of w^ter, a deep- blue color is produced (distinctive from nitric acid). Indigo in Presevce of Sulphite. — If a solution containing chlorate be colored light-blue with a drop or two of indigo in sulphuric acid, the solu- tion made slightly acid with sulphuric acid, and sodium sulphite added di'op at a time, the blue color of the indigo is at once destroyed, due to the sulphurous acid reducing the chloric acid and liberating chlorine or oxides of chlorine. 250 CHLOKIC ACID. SOLUBILITIES. Name of Salt. Soluble in Water. Soluble or Decom- posed by Acids. Insoluble in Water and Acids. Carbonates Sulpbates Sulphites Sulphides Nitrates Nitrites Chlorides Chlorates Hypochlorites Iodides Fluorides Phosphates I (Ortho) 1 Silicates Acetates Oxalates Tartrates Alkaline carbonates All sulphates except those in 4 Alkaline sulphites and acid sulphites of alkaline earths t Sulphides of the al - < kalies and alkaline ( earths All except some basic nitrates All those All All except those given in 4 All sulphites are deobmposed by acids, yielding (SOj) All: some by HCl, HjS evol.; others by HNO3 or aq. reg. S. separating All BaSOi , SrSO, , CaSO, , PbSO, 1 All except ( given in 4 All All f All except Pblj j Hglj, Bil3, Cujis ) and those given in I. *• ( Those of the alkalies < and a few others : I AgF, HgFj AH except those given in 4 All except those in 4 AgOl.PbClj.HgjClj (HgoCIo sol. in HNO3 or CI) Agl, Hgjij, decom- posed by CI Citrates Arsenites and Arsenates Chromates Those of the alkalies [ Potassic and sodic silicates (not na- I tive) All normal Those of the alkalies and Cr and Sn" (Mg and Fe" spar- ingly) I Normal alkaline tar I trates Alkaline citrates All decomposed by strong H2SO4 with few excep- tions All Many are decom posed by acids separati ng HjSiOi All Many native sili- cates (decom posed by HF) All All All f Those of the alkalies All I Those of the alkalies and Sr, Ca, Mg, Zn, I Fe", Cu All, with a few ex- ceptions DETECTION OF THE ACIDS. We have now become acquainted with the most important of the individual reactions of the acids, and the next thing in order is to determine some plan for recognizing them either when alone or mixed with others. The reactions occurring on the addition of hy- drochloric and hydrosulphuric acids in the separa- tion of the bases should always be carefully noted, as the presence of ceitain acids will be revealed on the addition of these reagents. If the addition of hydrochloric acid causes an evolution of gas, one or all of the following acids may be present — carbonic, sulphurous, or hydio- sulphuric acids. The first is indicated by being odorless and giving a white precipitate with lime or baryta water ; the others, by their odor and the action of hydrosulphuric acid on acetate-of-lead paper. Hydrocyanic acid, if present, would also be liberated at this point, and rdight be recognized by its odor and by confirmatory tests. If the 251 252 DETECTION OF THE ACIDS. original substance is a solid the same remarks ap- ply to the operation of dissolving it in acids. The color of the original solution and the ac- tion of hydrosulphuric acid show whether chromic acid is present or not. If the solution is red or yellow at first, and after addition of the hydro- sulphuric acid changes to a bluish green, chromic acid is present. It is not possible to make any scheme for the separation of the acids, similar to those used for the bases, but in the one given on the following pages I have endeavoi'ed to systematize the sepa- rations as much as possible. DETECTION OF THE ACIDS. SCHEME.— ACIDS. 253 Boil the solution containing the acids, with excess of NaaCOg ; filter hot. (Note 1.) Precipitate A. Contains bases and perhaps silicates, phosphates, and fluorides. Divide into two parts. Part 1°. Acidify with HNO3, and evaporate to dryness. Take up the residue with HNOj and HjO ; filter and wash. Ebsidhe 1°. A white, Rritty powder = SiOa. Test with meta- phosphate bead. Filtrate 1°. Add (NH4)eMo,Oa, a yellow crys- talline ppt., = H3PO4. Part 2°. Acidify with HCaHjOj, and filter. Residue 2°. Filtrate 2°. Reject. If no SiOj has been found in resi- ^w^HaH^H^^;^ due 1", treat residue 2°uithH^y04 (cone.) in a Pb or Pt dish, and see if fumes will etch Rlass; if they do=HF. Should SiPa have been found in resi- due 1°, treat 2° according to method given for fluorides in presence of silica. Filtrate A. Contains the acids. Divide into two unequal parts. Part 1° (M of A). Add HNO3 to faintly acid reaction; then NH.OH until slightly alkaline, and boil until all free NHjOH is driven oft. Divide into two parts (Note 2.) Part 1°. ■ Take a small portion of this neutral solution, and add BaClj. Awhite ppt. = acids of Group 1 . To thi'i ppt. add HCl: aIldissolveexceptBaS04; proves HjSOi. Part 2». Take a small portion of this nsntral solution, add AgNOs, and then acidify with HNOj. Ppf = acids of the ad Group yote the color of the ppt before and after adding HNO3. (Note 3.) Part 2° (14 of A.) Test a portion for HCIO3' by evaporation to dryness and adding a few drops of, concentrated sulphuricacid . The acid becomes yellow if chloric acid is present. Evaporate another por- tion to small bulk, cool and acidify with cone. HaSOd. and test for HNO3 with FeSO,. Brown ring = HNO3. HCIO3, HBr, HI, and H3Cr04 impede this reac tion, and should be removed. (Note i.) 254 DETECTION OF THE AOIDS. SCHEME.— Concluded. If Group 1° has been found by test on part la, take the remainder of this solution and divide into four parts: 1°. Acidify with HCI. and test for H3BO3 with turnneric paper. A red" colter := H3BO3. (Note 5.) Evapo- rate the same so- lution to dryness, and take up resi- due with a little HCI and water; a white iiisol. resi- due = SiOa. (Note 6.) Add CaClj and a httle NHiOH; filter ppt. if any (Note 7); dry, and if SiOa has not been found in 2^ test tor HF by etching test. If SiOg has been found, then te^ as gjiven for hydro- fluoric acid in pres- ence of silica. 3°. Acidify withHNO, add (NH4)eMo,Oai; a yellow ppt. = H3PO4. If ASjOe has been found in the test for bases, it should be removed by acidi fying with HCI. passing in HgS gas, and filtering before testing for H3P04. (Note 8.) I f HaCrO, i s present, the solu- tion will be yellow. Confirm by acidify- ing with HCjHjO,, and adding Pb(CQH302), , a yellow ppt. = PbCrO,. (Note 9.)' H3CO3. Test the original solution with HCI; if effervescence, pass the gas into lime-water. A white ppt., soluble in HCgHaOa, with effervescence proves HjCOs. If 2d Group has been found to be present by test of Part S», test remainder of solution for acids of this Group as follows : Divide into four parts. 1. Add a little CRa. and then a few drops of HCI, then NaClO verv care- fully; the I is set free, and colors CS2 purple; add more NaClO very carefully, and ti nally the I color disappears, and the Br is theU lib- erated and colors CSj brown. (Note 10.) 2. Acidify with HCI. and test with FeSO,, a blue ppt. =He(FeCy,),. The FeS04 solution must be recently made. Acidify with HCI and add FeoClj ; blue ppt. = HiFeCye. Acidify slightly with HjSOj.andif HCy is present it is known by its pe- culiar odor. Care should betaken not to inhale this gas. If HjFeCy,. or He(FeCy,),, has been f oundin 3 and 3, test for HCy as given in note 11. HCI. When HI,HBr,H4FeCy6,HeFe2Cy,3 and HCy have not been found, a white curdy ppt. on addition of AgNOj, in- soluble in HNO3 , and soluble in NH4 OH = HCI. HCI mny also be tested fnr by means of H2a04 and MnOj. If HI, HBr, HCy, HjFeCye. and H.Fe,- Cy]3 are present, see Note 12 for de- tection of HCI. H,S. If present, will be found on acidify- ing the sodium carbonate solution with acid, when HjS will be given off, and will turn paper moistened with Pb(CjH30,), black. H.S is usually found in testing for the bases, as it is liberated when dissolving the original substance. (Note 13.) NOTES TO SCHEME. 255 Note 1. — Since the presence of the heavy metals interferes with many of the tests for acids, it is always best to remove them before beginning the analysis. If the substance under examination is a solution, or a solid that is easily soluble in water or acids, they are precipitated by adding to the solution or solid substance a slight excess of sodium carbonate, heating to boiling, and filter, ing. The acids, with the exception of those pre- cipitated from acid solutions on neutralization (phosphates, etc.), will be in the filtrate as soda salts. This and the residue are tested as given in the scheme. Another way of obtaining the solution for analysis is, if the substance is a solid, to extract it thoroughly with boiling water and filter. The filtrate is boiled with a slight excess of sodium carbonate and filtered, and the filtrate tested for acids according to the scheme. The precipitate contains only the bases, and may be rejected. The residue left after extracting the original substance with water is dried and fused in a platinum crucible, if no reducible metals are present, with about four parts of sodium potassium carbonate ; the fusion is boiled with water until completely disintegrated, and filtered. 256 DETECTION OF THE ACIDS. The filtrate is tested in the same way as that obtained by boiling the original substance with the carbonate. If phosphoric acid is not found in the filtrate the residue, or a portion of it, should be dissolved in nitric acid and tested for that acid, as phosphates of the alkaline earths and alumina ai-e not completely decomposed by fusion in this way. Note 2. — Care must be taken to add only euough nitric acid to make the solution slightly acid ; it should then be boiled to expel carbonic acid before adding the ammouia. A precipitate sometimes forms at this point (bases soluble in excess of hot sodium carbonate) which should be filtered off. If much nitric acid is used the subsequent ad- dition of ammonia will form a considerable quan- tity of ammonium nitrate, which interferes with the test for borates, fluorides, etc. Note 3. — The color of the silver precipitate often indicates the presence of certain acids. For example, if the precipitate is yellow in the neutral solution and changes to white on the addition of nitric acid, this would show the probable presence of phosphoiTc acid (or arsenious), as well as the presence of acids of the second group. Acids of DETECTION OF THE ACIDS. 257 the latter group giving strongly colored precipi- tates could not be present (iodides, sulphides, and ferricyanides) if the insoluble portion of the pre- cipitate was white. Note 4. — The test for nitric acid should be de- ferred until it has been determined if any of these acids are present. Chloric acid, if present, is removed by evaporat- ing a portion of the alkaline solution to dryness and then igniting it. The residue is dissolved in very little water and tested for nitrate with sul- phuric acid and ferrous sulphate. Chromates (in- dicated by their yellow color) are removed by making the solution slightly acid with sulphuric acid and adding sodium sulphite ; the sulphurous acid liberated reduces the chromate to chromium sulphate, which is then precipitated by ammonia, and the filtrate after concentration tested for nitric acid. Hydriodic and hydrobromic acids are re- moved by precipitation, with silver sulphate. (This is readily made by dissolving a silver coin in a little hot concentrated sulphuric acid and after cooling diluting with water.)' Note 5. — Remember, in making this test, that the solution must be acid with hydrochloric, but not strongly so. 2£8 DETECTION OP THE ACIDS. ]S|"oTE 6. — Silica is not generally met with ex- cept in minerals, and if the mineral is one decom- posable by acid (except hydrofluoric) it is readily found by evaporating the acid mixture to dryness, with frequent stirring, at a temperature not ex- . ceeding 110°-115° C. The thoroughly dried mass (which snould give no odor of acid) is moistened with a little concen- trated hydrochloric acid (if silver or lead are present use nitric acid), some water added, and the mixture boiled for some time. On filtering, the silica will be left as a white, gritty powder. Sili- cates not decomposed by acids are fused with sodium potassium carbonate, the fusion digested with water, then, without filtering, acidified with hydrochloric or nitric acid and evaporated to dry- ness, and treated as just given for "decomposable silicates. Note 7. — Of course, other acids besides hydro- fluoric will give a precipitate at this point (phos- phoric, sulphuric, etc.), so the mere fact of a precipitate forming is no proof of hydrofluoric acid. It is well to try its solubility in acetic acid ; if soluble, no fluoride is present, and no further test is necessary. In the case of minerals decomposable by sulphuric acid, the test may be DETECTION OF THE ACIDS. 259 made to advantage directly on the finely pulver- ized substance. , Note 8. — If arsenic acid is present care must be taken to remove it completely by passing H^S into the hot solution (70° C.) for at least half an hour. The solution should only be slightly acid with nitric acid, since the latter in strong solution decomposes the HgS. Boil out the HgS thoroughly before adding the molyb- date. Phosphates, unless present in very minute quantity, may be tested for with " magnesia mix- ture" (MgCl, + NH4CI + NH4OH) instead of the molybdate. In this case the arsenic, if present, mufet be removed, as before, by acidifying (but with HCl) and precipitating with HgS. The filtrate is then made strongly alkaline with ammonia and the ma2;nesia mixture added. Note 9. — When chromates are present in a mixture, as already stated, they are reduced by H^S when precipitating the metals, and chromium will then be found among the bases as well as with the acids, although it may only be present in the acid form. In such cases, in order to deter- mine if basic chromium is present, it should be tested for in precipitate A where it will have 260 DETECTION OF THE ACIDS. been left by the treatment with sodium car- bonate. Note 10. — The hypochlorite should be added only a drop at a time at first and a weak solution used, for if the iodine is present in small quantity any excess of the hypochlorite would destroy the color. When a large amount of iodine is present, and the carbon diaulphide becomes very dark in color, instead of adding more hypochlorite to remove the color it is a better plan to filter the solution through a wet filter ; the disulphide, with its dissolved iodine, remains on the paper, and to the filtrate is added fresh carbon disulphide, and the test proceeded with as before. For traces of iodine it would be better to use the starch test. Cyanides, if present, should be removed before testing, as they interfere with the liberation of the iodine. For method of doing this, see Note 12. Note 11. — Unless the odor of the hydrocyanic acid is so decided as to be unmistakable, it is better to make the test with ferrous and ferric salt, obtaining the precipitate of Prussian blue. Or the test with alkaline polysulphide can be used, forming thiocyanate. In the presence of f erro or f erri cyanides these DETECTION OF THE ACIDS. 261 tests cannot be used. The simplest metliod of separation in such cases is to add an excess of hydrogen sodium carbonate (bicarbonate) to the solution in a flask fitted with a cork and bent delivery-tube (or better, connected with a con- denser), and boil, passing the vapors over into a solution of potassium hydrate. The heating should be continued until the liquid in the flask has nearly all distilled over. The hydrocyanic acid of the cyanide is set free in the reaction and passes over with the steam into the alkali, form- ing potassium cyanide, on which any of the usual tests for cyanide cau be made. Note 12. — The presence of cyanides and ferro- and ferricyanides interferes with the usual test for hydrochloric, hydrobromic, and hydriodic acids, so they must be removed before these acids can be tested for. This is done by pi-ecipitating them all as silver salts, filtering, drying, and ignit- ing the precipitate. The cyanogen compounds are all destroyed by this treatment, while the iodide, bromide, and chloride of silver are unchanged. The ignited residue is then fused with sodium potassium carbonate, and the fusion extracted with water and filtered. The filtrate contains the acids as sodium iodide, bromide, and chloride. 262 DETECTION OP THE ACIDS. Another simple way of decomposing the fused silver salts is to cover them M^ith a little water, add a few drops of sulphuric acid and a piece of zinc, and allow to stand for some time. The solu- tion will contain the acids as zinc iodide, bromide, and chloride, and the residue metallic silver. 2 AgCl + Zn = 2 Ag + ZnCl^. Chlorides cannot be tested for in the pres- ence of iodides and bromides by the ordi- nary methods, and some special method must be adopted. There are several different ways of accomplishing it, one of the simplest of which is the following : Make the solution slightly acid with sulphuric acid, add a concentrated solution of ferric sulphate (ammonium ferric sulphate is a convenient form in which to use it), and boil ; the iodine is liberated and carried off with the steam. As soon as no more iodine is set free (indicated by absence of violet fumes, or by test with paper moistened with starch'paste giving no blue color) solution of potassium permanganate is added and the solution again boiled. If the color of the per- manganate disappears, more is to be added, and in sufficient quantity to give a violet or purple color DETECTION OF THE ACIDS. 263 to the solution that does not disappear quickly, even on boiling. The permanganate decomposes the bromide, which is evolved as free bromine, and brown ox-' ide of manganese precipitates. The solution is filtered, and the remaining permanganate decom- posed by adding a few drops of alcohol, warming, and filtering. Silver nitrate is added to the filtrate, giving, if chlorides are present, a white, curdy precipitate of silver chloride. Another method of detecting chlorides in the presence of iodides and bromides is to add to the neutral solution lead dioxide and acetic acid and boil until no more iodine or bromine is evolved and the solution becomes colorless. Filter and wash the precipitate with hot water, and test in the filtrate for chlorides with silver nitrate. Note 13. — When there is any doubt as to the presence of sulphide the surest plan is to fuse the finely divided substance with sodium or potassium hydrate in a silver or nickel crucible and test the fused mass on bright silver. (See hydrosulphuric acid.) Section III. COMPLETE ANALYSIS, INCLUDING THE METALS AND ACIDS OF ALL THE GROUPS. Having completed our study of the individual reactions of the different bases and acids as well as the separation and detection of the members of each group when present together in solution, all that now remains to be accomplished is to connect these different group - separations into one con- nected method that will be applicable in almost all cases that may arise. Before proceeding further it will be well to call attention to the fact that success as an analyst will depend on several things. Rapidity in work can only be attained by neatness and an intelligent use of time. Lack of neatness and order not only cause the loss of much valuable time, but they are also fruitful sources of error. The student should endeavor to make use of every moment, and for this reason should ac- custom himself to carry on several operations at the same time. For instance, while filtering a 264 COMPLETE ANALYSIS OF ALL THE GROUPS. 265 group-precipitate lie can test the first few drops of tlie filtrate to see if the precipitation has been complete, and, if so, he can then determine if any of the metals of the succeeding group are present. He then knows in just what manner to treat the filtrate when it is ready — whether to add the reagents for the next group, or, if his tests have shown that it is not present, he can go on to tlie precipitation of the remaining groups. Be- sides making these tests, he will very likely have had time to investigate some other precipitate that has already been filtered and washed. Another point of special importance is to know exactly under what conditions each precipitation should be made, and every one should make it a rule, which should be rigidly adhered to, never to add any reagent until lie Icnoios wJiy it is used, and what it is expected to do. No one will ever be a good and intelligent analyst who adds a reagent simply because his book tells him to do so. He must know the reason why, and if he does not, he should go no further until he does know. After a precipitate has been obtained, it is very important in most cases that it should be thoroughly washed, and the thoroughness of the washing should be determined by a confirmatory 266 COMPLETE ANALYSIS OP ALL THE GROUPS. test. For instance, test the washings with silver nitrate when washing out a solution containing chlorides. Never, because you are in a hujTy, slight this important part of your analysis. Every student should bear in mind when mak- ing an analysis that he has two objects in view. One, of course, is to determine what the constitu- ents of the substance are ; and the other, which is of equal or even greater importance at this stage of his course, is to gain an intelligent understand- ing of the methods used. Let us return now to a study of what is neces- sary in order to make a complete analysis of any mixture. In the first place, the substance to be examined, if not already in the liquid form, must be gotten into solution. Frequently this is a very simple operation, as, for example, when the ma- terial is soluble in water or acids ; but in some cases it is more difficult, as in the case of a sub- stance insoluble in acids, and requiring fusion to decompose it, such as barium sulphate and many silicates. Before attempting to dissolve any substance (other than alloys) it should always be pulverized and gotten into the finest powder possible; the COMPLETE ANALYSIS OF ALL THE GROUPS. 267 finer it is, the more rapidly will it go into solu- tion, or be decomposed by fusion if the latter proves necessary. It is well at this point, before proceeding to get the material into solution, to make a preliminary examination on charcoal and in the glass tube, as much valuable information is often gained in this way without the expenditure of much time or labor. PEELIMINAEY EXAMINATION. Test A. — Heat a portion gently with oxidiz- ing flame upon charcoal. Sb. — White pulverulent volatile coat. The compound often continues to form in dense white fumes after cessation of blast. The coat- ing disappears before reducing flame, tingeing it a pale yellow-green. As. — White very volatile crystalline coat. The coat disappears before reducing flame, tinge- ing it pale blue, and evolving a characteristic garlic odor. Test B. — Heat a portion gently with the reduc. ing flame on charcoal, and note results. Then add 268 COMPLETE AWALTSIS OF ALL THE GROUPS. some sodium carbonate and heat strongly with the reducins: flame for three or four minutes. As. — Garlic odor and a faint white volatile coat. 0(1. — Dark-brown volatile coat, sometimes shad- ing to greenish yellow, and usually surrounded by a variegated coloration resembling the colors of peacock feathers. Zn. — White not easily volatile coat, yellow when hot. 8n. — White non-volatile coat close to assay, and usually small in amount. White coats may form from Pb, Bi, or alkalis; yellow coats from Pb or Bi ; brown or red coats from Cu or Mo ; and the ash of the coal may be white or red. If any coat forms, examine it for Zn and Sn by moistening it with cobalt solution and blowing a strong blue flame on the substance. The coating turns green if Zn or Sn are present. The coatings from other elements will not prevent the cobalt coloration. Test C. — Mix a portion of the substance with more than an equal volume of bismuth flux* and * Formed by grinding together 1 pt. KI, 1 pt. KHSO,, 2 pts. S. COMPLETE ANALYSIS OF ALL THE GROUPS. 269 heat gently upon a plaster tablet with the oxidiz- ing flame. Ph. — Chrome -yellow coat, darker hot, often covers the entire tablet. 8n. — Brownish-orange coat. Afi. — Reddish-orange and yellow coat, darker hot. 8h. — Orange to peach-red coat, very dark when hot. Hg. — Gently heated, bright yellow and scarlet coat, very volatile; becomes all scarlet on stand- ing ; but if quickly heated, the coat formed is pale yellow and black. Bi. — Bright scarlet coat surrounded by choco- late-brown with sometimes a reddish frino;e. Te^ts in Glass Tube. — The glass tube used should be about 3 inches long and J inch in diam- eter, and be of hard glass. One end is closed by fusing it in the Bunsen flame. Some prefer to use a tube that has had a small bulb blown on one end. A little of the powdered substance is introduced into the tube so that it rests in the lower or closed end. It is then heated, gently at first, and finally at the highest heat possible. A. a. The substance remains unchanged; ab- 270 COMPLETE ANALYSIS OF ALL THE GEODPS. seijce of organic matter, volatile or fusible sub- stances, and those containing water of crystalliza- tion. b. The substance blackens; indicates organic matter. c. The substance fuses without blackening, and no water is given off ; add a small piece of char- coal ; deflagration indicates nitrates or chlorates. d. Water is expelled, condensing in upper part of tube ; indicates either water of crystallization (such substances generally fuse, and after expul- sion of the water solidify again), or that derived from hydroxides or salts yielding water on igni- tion (NaHCOg), or moisture. The reaction of the water should be tested ; if alkaline it indicates ammonia, if acid a volatile acid (H3SO4, HCl, etc.). B. Gases are Evolved: SO2, recognized by its odor, indicates either sulphides, sulphites, or sulphates. JV0.2, known by its red color and odor, indi- cates nitrites or nitrates. Care should be taken not to confuse it with bromine, which also forms a red gas. 01, Br, and I may all be liberated, and when alone are easily recognized by their odor and color. COMPLETE ANALYSIS OF ALL THE GROUPS. 271 NQ or HCN indicates cyanogen compounds decomposable by heat. May be recognized by the odor. H.^S indicates sulphides (not anhydrous), also thiosulphates. Recognised by odor or lead-ace- tate paper. NH^ indicates ammonium salts or the decompo- sition of cyanides containing water. Nitrogenous organic matter may also yield ammonia. C. A Subli'mate Forms : S indicates sulphur in the free state or that de- rived from sulphides, many of which liberate sul- phur on heating in this manner. Kecognized by forming reddish-brown drops that turn yellow on cooling. Hg. — Many of the mercury salts volatilize un- altered (the oxygen salts are decomposed on igni- tion), so that the sublimate in this case may be either metallic mercury or a salt of mercury. If doubtful, make another test by mixing some sodium carbonate with the substance before fus- ing, and covering this mixture, after placing it in the tube, with a layer of the carbonate, the mer- cury salts will be decomposed, yielding a subli- mate of metallic mercury. Arsenic. — Metallic arsenic gives its character- 273 COMPLETE ANALYSIS OF ALL THE GROUPS. istic dark mirror. The oxides yield a white crystal- line sublimate of arsenious oxide. The sulphides give sublimates that are reddish yellow while hot and yellow on cooling. The oxide and sul- phide are soluble in ammonia. Antimonious oxide also gives a white crystal- line sublimate, insoluble in ammonia. Besides the tests just given, it is well to heat another portion of the substance mixed with potas- sium bisulphate (made by heating potassium hydro- gen sulphate at a low heat, just sufficient to drive out the water, 2KHSO4 - H,0 = KaSO,, SO3) in a glass tube in the same way as in the first test. Many substances not afPected by heat alone are decomposed by this reagent. The gases evolved are recognized, as in first test, by their odor color, etc. Fluorine may also be detected in this way, as the gas liberated attacks and corrodes the glass, and should silica be present a drop of water held at mouth of tube will become opaque. Silica. — Test a few grains of powder in salt of phosphorus bead for silica. The preliminary examination having been com- pleted, we are ready to proceed with the opera- tion of getting the substance into solution. Should organic matter, such as oil, etc., be pres- COMPLETE ANALYSIS OF ALL THE GROUPS. 273 ent, it must be removed by digesting with proper solvents, as naphtha, etc., or it may in many cases be desti'oyed by ignition. Solution. — It is always advisable to follow some systematic plan when dissolving any mate- rial given for analysis. The following one, al- though it may seem tedious at first glance, will be found shorter in the long run than any hap- hazard attempts at solution. - It will very seldom, if evei', happen that all of the bases and acids provided for will be found in any one mixture, so that the method will be materially shortened in almost all cases. When it is desired to get an idea as to the combinations in which the different constituents of a substance exist, it is necessary to analyze the different solu- tions separately for both bases and acids. Even this will not always give the true combinations, for interactions may take place on the addition of the solvent that will wholly or partially change them. If it is only required to deter- mine the actual constituents without regard to their combinations, which is usually the case, the different solutions may be analyzed as one mixture, provided they do not precipitate each other. For example, before mixing a water or 274 COMPLETE ANALYSIS OF ALL THE GEOTTPS. nitric acid solution with one in hydrochloric acid, the former ought always to be tested with hydrochloric acid to see if any silver, lead, or mercurous salts are present, and if fcund they should be completely precipitated and filtered off, after which the solutions may be mixed together and analyzed. The student should make it a rule always to put aside a part of the original material for special tests, and in case of accident to the main analysis. The solution of a solid substance is conven- iently considered under two heads, viz. : 1°. The substance is neither a metal nor an alloy. 2°. The substance is a metal or an alloy. Substances of the 1° class are to be treated as follows : A. Boil some of the finely pulverized substance with water, repeating this treatment with fresh portions of water until thoroughly extracted. a. All dissolves. Test solution according to Scheme No. 1. h. A residue remains. Filter, and evaporate to dryness a few drops of the filtrate in a platinum capsule, and see if any COMPLETE ANALYSIS OF ALL THE QKOtJPS. 275 appreciable residue remains ; if so, test filtrate ac- cording to Scteme No. 1, or reserve to be com- bined with acid solution B or C. B. Residue insoluble in water. Treat a portion of this residue with dilute hy- drochloric acid, applying heat if necessary. If it does not dissolve completely, allow it to settle and then decant the solution, and to the residue add some concentrated hydrochloric acid, and boil.* If everything has now gone into solution the acid liquids are evaporated to small bulk, diluted with water, and pi-ecipitated with H^S in the usual way. a. In case the foregoing treatment has not effected complete solution, keep the dilute acid liquid, and the one with strong acid containing the undissolved residue for future use. C. Hydrochloric acid not liaving effected com,- plete solution, take another portion of B and boil with nitric acid, and add water. If all dissolves, evaporate to expel excess of acid, dilute with water, and add to solution A, h. Should the ad- dition of water cause a y^recipitate, it is proba- bly bismuth, which will redissolve on the addi- * Tlie presence of carbonate is always revealed at this point by the evolution of COj. Sulphides also are frequently detected here by the evolution of HaS. Also many cyanides. 276 COMPLETE ANALYSIS OF ALL THE GEO UPS. tion of a few drops of nitric acid. The solution is now tested in the usual way. a. If the boiling nitric acid does not effect complete solution, test a few drops of the liquid, after dilution with water, with dilute hydro- chloric acid for silver and lead (mercurous salts would have been oxidized to mercuric). b. Silver or lead salts are not present. — Mix the nitric acid solution and residue with the hydrochloric acid solutions (B, a) and the residue insoluble in that acid ; boil the mixture, which now forms aqua regia, for some time ; and if there is still a residue left, dilute, and filter and wash precipitate, and treat the latter according to D. The filtrate, after getting rid of the excess of acid by evaporation is combined with the water solution A h and analyzed in the usual manner. c. Silver or lead salts are present. — Dilute the strongly acid solution and filter, washing the residue thoroughly. Then treat this residue first by boiling with fresh hydrochloric acid (not that used in B, a), and if this does not dissolve it com- pletely, then add some nitric acid, forming aqua regia, and boil again. If a residue still remains, filter (after dilution) and wash and proceed with COMPLETE ANALYSIS OF ALL THE GROUPS. 277 it as in D. Reserve filtrate, which is to be mixed with the water and nitiic acid solutions later on. Silver and lead salts being present, the water and nitric acid solutions are mixed, the silver, lead, and mercurous salts (the latter from the water solution) precipitated by hydrochloric acid and the filtrate mixed with the aqua regia solution, excess of acid expelled by evaporation to small bulk, the solu- tion diluted and pi'ecipitated with H^S as usual. D. S,esldue insoluble in aqua regia. This residue may contain AgCl, PbSO^, BaSO,, SrSO^, CaSO^, SiO^ , and silicates, AlA , CrA,CBF,,C, S, SnO^. a. If the residue insoluble in aqua regia is white, take a few grains of it and add a drop or two of ammonium sulphide ; if silver or lead is present it will turn black. In case D is daik colored test a small portion on charcoal or plaster for Ag and Pb. If found, by either test, proceed according to c ; if not, according to residue 2°. b. Test another small portion in a glass tube closed at one end for S if its presence has not already been determined in the preliminary tests. c. Silver and lead salts are present. Take some of the residue, and heat with a con- centrated solution of NH^CsHgOa; filter, and re- peat treatment until lead salts are all removed. 278 COMPLETE ANALYSIS OF ALL THE GROUPS. FiLT. 1». Lead salts -i- 2 parts. 1°. 2°. Add HjSO,l Add HCI, -white ppt. filter if neces = PbSO, sary, and add a few drops of BaClj, white ppt. = BaSO, = HjSOi. Residue 1". Warm residue wilh KCy (if S is present, di- gest in the cold), filter, and repeat the treat- ment until Ag salts are all removed. FiLT. 2°. AgCy.KCy, add (NH4)2S, brown ppt. = AgaS — . Dissolve AgjS in hot HNO3. dilute, filter, and to filtrate add HCI — white ppt. — AgCl = Residue 2°. If S is present. Iieat residue in porcelain dish until the S has complete- ly volatilized; then mix the residue with NaKCOa and a little NaNOa, and fuse in a platinum ci'uci- ble; dissolve the fusion in boiling water, filter and wash. Bgsidcb 3°. BaCOa, SrCO,, CaCOa (SiOs ,_ AljOa. CrjOs, SnOj) Treat with HCjHaOj; heat and filter. FiLT. 4°. _BaA2, Sr- As, CaA^. Test in the usual way. Residue 4°. SiOj. SnOj, etc. Place in te.st-tube with Zn and strong HC31 and a few small pieces of platinum— SnOg. if present, will be reduced to Sn. Dissolve HCI, and test with HgCI,. Filtrate 3°. NajSiOa, NaF, NajCrOj. y!a^A.\,0,.N■■^,'^0,■ NajSnOa. (?) Make slightly acid «ith HC ' and evaporate to dryness; if Cr is present, add a little alcohol; (do not heat much above 100° C). Add HCI to the dry residue, boil, dilute, and filter. Ppt. 5°. SiOj, SnOj (?.) Pi'ove SiOj by salt f phosphorus bead. Sn by re- duction with Zn. Filtrate 5°. NaF. AljCle. Cr,Cis. NajRO,, SnCU (?). Pass H2B gas into the solution and filtei- f i-om any ppt. Add NH4OH to faintly alkaline reaction and filter. FiLT. 6. NaF, AljClj. CrjCle + Na^SO,. Ppt. 6. SnS„S + Pt-I-Ss ('). Add a little concentrated HCI and boil; the tin dis- solves as Sn- Cli. Nearly n en t rali ze and precipi- tate the tin as ineiastan- nic ax;id with Na.^SOj. Ppt. 7°. Alj, (OH)a, and Cr2(0H),). Test in the usual way. Filtrate 7°.— (Divide in two parts.) 1. Acidify with HCI and add BaClj, white ppt. = BaSO , = HgSOt. be to take some of residue No. 2 and test that directly, either by etching test, or if silica is present by passing SiFi into water. Add CaClj and let stand tor some time; filter, and test ppt. for Fl by etching test. A better way to test for Fl would COMPLETE ANALYSIS OP ALL THE GftOFPS. 2M Metals or Alloys. Case 2°. The substance is a metal or an alloy. Boil with HNO3 Sp. Gr. 1. 2, and evaporate nearly to dryness ; add a few drops of HNO3, dilute with water, boil and filter. Rksiduk 1° Residue is white and non-metallic* Boil with strong solution of H2C4H40a; if any residue remains filter and wash. Filtrate 1°. Test according to Scheme No. 1. FiLTBATB 2°. Residue 8°. Add a few Place in a platinum capsule with a piece of zinc and a drops of HCl. and little strong HCl. The HjSnOs will be reduced to metallic pass in HjS, an tin. Remove the zinc, dissolve the tin in HCl, and test orange ppt. = solution with HgCU; a white or gray ppt. proves Sn. SboS.. Gold and platinum, if present, will be left undissolved by the HCl, and should be tested for by dissolving in aqua regia, dividing solution into two pans, and testing tor gold with FeSO,, and for platinum with KCl. The substance now being in solution, we are ready to proceed with the analysis for the de- tection of the bases. As the separation of the metals of each group has already been explained, all that is necessary now is to show the connec- tion between the different group schemes. This is done in the following outline, in which it is seen that each group is tested for in the filtrate * If residue is metallic dissolve it in aqua regia and test solution for gold and platinum. Platinum alloyed with a large amount of silver is soluble in nitric acid. 2S0 COMPLETE ANALYSIS OF ALL THE GROUPS. obtained after precipitating the preceding group. The analysis is then, first, a separation into groups, and, after that has been accomplished, the separation of the individual members of the group from each other. Solution contains 1°, 2°, 3°, 4°, 5°, 6° groups. Add HCl. Ppt. AgC'l,PbClj,HgaCl,. FiLTBATE. 1°, 2°, 3', 4°, 5° (2° div.), and 6° groups. Add HjS. Ppt. 5° and 6° groups. Add (NH<)jSa,. Residue:. 5° group. FiLTEATK. 6° group. FiLTHATB. 1° 2° 3° 4° groups. Add NHiCl + NH4OH + (NH<)jS. Ppt. 3° and 4° groups. Filtrate. 1° and 2° groups. Add (NH4)jCOs. Ppt. 2° group (ex Mg). Filtkate. Mg and 1° group. Besides the main analysis, it is necessary to make special tests on portions of the original solution to determine the state of oxidation of certain of the metals, usually for iron, tin, and arsenic. The two oxides of mercury are gen- erally determined in the regular analysis ; but it must be remembered that if it has been necessary to boil with acids in order to get the substance COMPLETE ANALYSIS OB' ALL TIlPl GROUPS. 381 into solution, that the mercurous salt may have been changed to mercuric. In cases of this kind, if mercuric -salts only have been found in th« regular course of analysis, it is always necessary to test specially for mercurous compounds. This may be done by digesting some of residue B with cold dilute hydrochloric acid and filtering; the residue is then washed with boiling water and tested with ammonia in the usual way for mercu- rous salt. If tin has been found, the oxide present is de- termined by testing some of the hydrochloric acid solution (water solution also should be tested in same way) with mercuric chloride for stannous salts and with sodium sulphate or nitrate for stannic. The two oxides of iron are tested for in the hydrochloric acid and water solutions with potas- sium ferricyanide for ferrous and thioc}anate for ferric salts. For arsenic, the difference in color of the silver salts, the tests with copper sulphate, the mag- nesium mixture test, and the time required for precipitation by HgS will generally show which oxide is present. 282 COMPLETE ANALYSIS OF ALL THE GROTTPS. THE ACIDS. The different ways of preparing the solution for the determination of the acids have already been explained, as well as the methods employed for their detection ; so it will be unnecessary to say anything further on that point, with the ex- ception of a few words in regard to the analysis of silicates and cyanides. Silicates. — The method generally employed in the analysis of silicates is to decompose them either by boiling with acid or by fusion with sodium-potassium carbonates (in the latter case the fusion is dissolved in water and made acid with hydrochloric acid), evaporating the acid mix- ture to dryness, heating to about 110" to 115° C. until no odor of acid is perceptible, then moisten- ing the residue with a little strong acid, adding water, and boiling until only a white residue of silica is left. This is filtered out, and the filtrate analyzed in the usual way. The alkalies of course cannot be tested for in this solution in case the silicate has been decomposed by fusion. When required to determine them, a separate portion is treated with hydrofluoric and sulphuric acids, or COMPLETE ANALYSIS OE ALL THE GROUPS. 283 a part may be fused with baiium hydrate, as explained under tests for silicic acid. Cyanides and Ferro- and Ferricyanides insolu- ble in acids ai-e best decomposed by boiling with strong sodium or potassium hydrate, then adding sodium carbonate and boiling again. On filtering the residue will be free from cyanogen (except silver cyanide), and the filtrate will contain all of the cyanogen in combination with the alkalies, and besides this all of those metals whose hydrox- ides are soluble in alkali (Zn, Al, Pb, As, etc.). Before testing this solution for cyanide, ferro- and ferricyanide, etc., sodium sulphide should be added to it drop by drop as long as it causes any pre- cipitate, but avoid a large excess of the reagent. This treatment removes those metals whose sul- phides are insoluble in alkalies (Zn, Pb). Filter and add dilute sulphuric acid to the filtrate until it is slightly acid, then add HoS if the solution does not already smell of that gas. This will pre- cipitate any metals of the sixth group, and after filtering the solution may be tested for the cyano- gen acids, as already explained. REAGENTS. Hydrochloric acid, HCl (strong), sp. gr. 1.12, 24^ acid. Hydrocliloric acid HCl (dilute), sp. gr. 1.06. Nitric acid, HNO3 (strong), sp. gr. 1.2, 32^ acid. Nitric a£id, HNO3 (dilute), sp. gr. 1.1. Sulphuric acid, H2SO4 (cone), sp. gr. 1.84, 95^ acid. Sulphuric acid, HaS04 (dilute), sp. gr. 1.14. Acetic acid, HCgHgOg, sp. gr. 1.04, 30^ acid. Hydrosidphuric acid, H.^S, is used either in the gaseous form or as a solution in water. It is generally prepared from ferrous sulphide (FeS) and hydrochloric or sulphuric acid. The water solution is obtained by saturating water with the gas. It does not keep well and must be freshly prepared from time to time. 384 REAGENTS. 385 Tartaric acid, ^.f^Jlf)^ (crystallized), dis- solved in water as needed for use, 1 part acid to 3 of water. Potassium hydroxide, KOH, solution of sp. gr. 1.125, Ibfo KOH. Sodivm hydroxide may be nsed instead of the above. Potassium chromate, K2Cr04, 1 part to 10 parts water. Ainmonium hydroxide, NH4OH, solution of sp. gr. 0.96 = 10^ NHg. Sodium carho7iate, NagCO3,10H2O, 2.7 parts to 5 of water. Sodium hydrogen phosphate, ]Sra2HP04,12H;jO, 1 part to 10 parts of water. Sodium acetate, '^SiC^^O.^jd^.f), 1 part to 10 parts of water. Sodium hypochlorite, N^CIO. Some bleaching powder (chloride of lime) is dissolved in water, and sodium carbonate solution added as long as a precipitate forms. Allow the precipitate to settle and then siphon off the clear liquid for use. Ammonitim carhonate, (^'^^.f^O.,,, 1 part of salt, 4 parts of water, and 1 part of ammonium hydroxide. 286 REAGENTS. Ammonium sulphide, (NIii)S and (NH4)8Sx, Two solutions of ammonium sulphide are re- quired, one of the normal sulphide for the precip- itation of the 3° and 4° groups and one of poly- sulphide for the separation of 5° and 6° groups. The method ordinarily used for the preparation of the noi'mal sulphide does not yield this com- pound according to the researches of Bloxam (see Jour. Chem. Soc, Apr. 1895), and he gives the following: HaS gas is passed into ammonium hydroxide solution and small portions are taken from time to time and tested as follows : An excess of solution of I'ecrystallized cupric sulphate (1 cc. = about 0.1 grm. of CUSO4) is placed in a flask, water is added, and then a few drops of the sulphide solution, and the mixture well shaken. It is then filtered (the filtrate must have a blue coloi-, showing the presence of an excess of copper sulphate) and the precipitate washed with boiling water until all the copper sulphate has heen washed out. Dilute solution of potassium hydroxide is added to the filtrate a drop at a time. If the first drop or two produces a precipitate, showing that the solution is neutral, the precipitate of sulphide, that has been thor- REAGENTS. 287 ouglily washed, is shaken with airnnonium hy- droxide and filtered. A blue filtrate shows that the portion of ammonium sulphide tested con- tained free ammonia, and more HjS is required. If no blue color is produced, the solution then consists of (NH4)aS. If on testing the first filtrate with KOH it had shown an acid reaction, this would indicate the presence of NH^HS. The following equations explain the reaction taking place under the different conditions : {a} SCuSO, + 2NH,HS = 2CuS + (NH4),S04 + H,S04 + CuSO,. (h) 2CUSO4 + (NH^oS - CuS + (NH4)3S04 + CuSO,. (c) 3CuS0, + (NHO^S + 2NH4OH = CuS + Cu(OH),+ 2(NH4),SO, + CuSO.. From the above equations it is evident that when the filtrate has an acid reaction, as in (a), the solution of sulphide contains NH4HS. If the reaction is neutral and the precipitate gives no blue solution on treatment with ammonia, then the solution must be one of (NH4)2S (b). It is 288 REAGENTS. better that the solution contain some NH4HS rather than free ammonia. The polysulphide is made from the normal one by addition of a little free sulphur, which soon dissolves and the solution becomes of a red color. Ammonimn oxalate (NH4)aC204, 1 part to 24 parts of water. Ammonium chloride^ NH4CI, 1 part to 8 parts of water. Annmoniumj sulphate, (NH4)3S04, 1 part to 5 parts of water. Ammonium molyMate, (NH4)6Mo7024. Dis- solve 1 part of molybdic acid MoOg in 4 parts of ammonium hydroxide solution, filter quickly, and pour into 15 parts by weight of nitric acid, sp. gr. 1.2. Ammonium acetate, NH4C2H3O2. Add acetic acid to ammonium hydroxide until solution is neutral. Ammonium thiocyamate, NH4CNS, 1 part to 10, parts of water. Potassium ferrocyanide, 'K^^FqC^^, 1 part to 10 parts of water. Barium chloride, BaCl2,2H80, 1 part to 10 parts of water. REAGENTS. 289 Calcium Jiyd/roxide (lime-water), Ca(0H)3, a saturated solution. Oalcium chloride, CaClgjeHgO, 1 part to 5 parts of water. Calcium sulphate, CaSO^, saturated solution. ;' Magnesium sulphate, MgSOijVHjO, 1 part to 10 of water. Ferric chloride, Yq^(u\, 1 part to 10 parts of water. Lead acetate, ^^{G^^O^z, 1 part to 10 parts of water. Siher niPrate, AgNOg, 1 part to 20 parts of water. Mercuric -chloride, HgClg, 1 part to 16 parts of water. Stannous chloride, SnClj, 1 part to 10 parts of water. Some metallic tin and a little free hydrochloric acid should be kept in this solution to prevent oxidation of the stannous chloride. Hydrochloroplatinic acid, HgPtClg, 1 part to 10 \ parts of water. Cobaltous niPrate, Co(N03)3, 1 part to 10 parts s of water. Potassium pyroantdmoniate, KgHgSb^O^, dis- 290 EEAGENTS. solved as required in hot water and the solution filtered. Potassium ferricycmide, ]L^(Fe,G^^z. It is best to keep this reagent in the dry state and dis- solve in water when needed, as its solution when kept is liable to partial reduction to fen*ocyanide. Potassiwm cymiide, KCN, dissolved in water as required. Barium ca/rbonate, BaCOg. Prepared by pre- cipitating a solution of barium chloride with ammonium carbonate, washing the precipitate thoroughly, and then placing it in a bottle and covering with water. When wanted for use it is shaken up with the water and the milky-looking solution used. Ferrous sulphate, FeS04,7H20, dissolved in cold water as required. Sodium potassium carhonate. Used in prefer- ence to either carbonate alone, as it is more easily fusible. Made by thoroughly mixing 13 parts of anhydrous potassium carbonate with 10 of dry sodium carbonate. Borax, Na2B4O7,10Hs,O. Sodium, ammonium Tiyd/rogen phosphite (phos- phorus salt; microcosmic salt), ]S'aNH.HP0„4H,0. REAGENTS. 291 When fused it becomes sodium metaphosphate, NaPOg. Potassium chlorate, KCIO3. Lead dioxide, PbOg. INDEX. Acid: boracic 302 carbonic 213 chloric • 247 chromic 188 hydriodie 226 hydrobromic 222 bydrochloric 218 hydrocyanic 230 hydroferricyanic 235 hydroferrocyanio 234 hydrofluoric 205 hydrosulphuric 236 nitric 242 phosphoric 196 silicic 209 sulphuric 192 Aluminium 52 Ammonium 25 Antimony 152 Arsenic 157 Bases — acids— salts, defini- tions of 5 Barium 33 Bismuth 127 Cadmium 134 Calcium 37 Chromium 57 Cobalt 82 FAOE Copper 131 Equations 9 Gold 165 Group I (bases). 19 •' II " 32 " III " 52 " IV " 64 V " 110 " VI " 143 I (acids) 188 II ■' 218 "III " 342 Iron 64 Lead..- 115 Magnesium 39 Manganese 77 Mercury 120 Nickel 87 Notes to scheme 1 48 II 96 " IV 107 V 175 " "for acids. . . 255 Oxidation and reduction ... 12 Phosphates— separation in presence of 104 Platinum I'^'O Potassium. . , 20 Reagents 284 393 294 INDEX. PAGE Scheme of separation : Groups I and II 47 " III " IV 95 " m " IV(BaCos) 103 " III " IV (Phos- phate). . . 104 " V "VI 174 Scheme of separation for acids 253 Separations: Group 1 29 " II 43 " IllandIV 92 PAGE Separations: Group V 136 Silver Ill Sodium 23 Strontium 35 Table of reactions : Groups I and II 47-48 Scheme II 95-96 " III 103-104 " IV 106-107 V and VI 174 Tin 145 Titanium • 60 Zinc 73 SHORT-TITLE CATALOGUE OF THE PUBLICATIONS OF JOHN WILEY & SONS, New York. London: CHAPMAN & HALL, Limited. AREANGED UNDER SUBJECTS. Desciiptive cireulars sent on application. Boolfs marked with an asterisk are sold at net prices only. All books are bound in cloth unless otherwise stated. AGRICULTURE. Cattle Feeding— Daiky Practice— Diseases of Animals- Gardening, Etc. Armsby 's Manual of Cattle Feeding 12mo, $1 75 Downing's Fruit and Fruit Trees Svo, 5 00 Grotenfell's The Principles of Modern Dairy Practice. 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