LABORATORY MANUAL OF Medical Chemistry CHASE Cornell University Library RC 407.C48 A laboratory manual of m^j'SfuSiMl ''' 3 1924 000 306 583 The original of tiiis bool< 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/cletails/cu31924000306583 LABORATORY MANUAL OF MEDICAL CHEMISTRY A LABORATORY MANUAL • OF Medical Chemistry CONTAINING A SYSTEMATIC [ LIBbIr COURSE OF EXPERIMENTS Vu -^ Laboratory Manypulation and Chemical Action, the Non - Metallic Elements and the Medicinal Metals, Quantitative Processes a^fiied to Sanitary Water Analysis, Medicinal Organic Com;pounds, Proteids, Digestion, Blood, Milk, Urinalysis and Toxicology IRA CARLETON CHASE, A.M. Professor of Chemistry and Toxicology in the Medical Department, and Professor of Analytical Chemistry in the Scientific Department of Fort Worth University FORT WORTH, TEXAS: Published by W. W. UNDERHILL 1897 1i No. ruv V Copyright, 1897 By IRA C. CHASE ^W&-CC _ PRINTED AND BOUND BY J, Horace McFarland company MT. PLEASANT PRINTERY HARRI3SURQ, PENNSYLVANIA PREFACE. This volume is the outgrowth of several years of lab- oratory instruction given medical students. The exercises have been used in various forms by seven classes since the organization of the Medical Department. They have been gradually shaped by the peculiar needs of students of medi- cine, and by personal visitation and careful study of the laboratories and methods of instruction in nearly all of the leading medical schools of the United States during the years 1893-6. Since there is not known to the author a small volume of modern, progressive laboratory exercises, covering the various branches of medical chemistry here presented, it has been decided to put these exercises into permanent form for the use of our own students. The aim has been to give such a scope to the earlier exercises as to quickly furnish the student a broad chem- ical foundation necessary to an appreciative understanding of the complex field of medical chemistry; to escape, on the one hand, the danger of teaching pure theoretical chemistry, unadapted to the needs of a physician; and, on the other, the more dangerous error of mechanically drill- (v) VI PREFACE. ing tests and operations which the student has not the chemical knowledge to remember nor intelligently apply. The work has been kept within such limits that it can be completed in two sessions by students working two lab- oratory periods each week. The more difficult and elaborate experiments have been reserved for the lecture table. The author trusts that the errors, so easily creeping into the first edition of a work covering so large a field, will be charitably regarded. Fort Worth University, I. C. C. Fort Worth, Texas, October is, 1897. CONTENTS. PACK Laboratory Manipulation and Chemical Action .... 1-13 The More Important Non-Metallic Elements 15-53 Reactions, Analytical and Synthetical, of the Princi- pal Medicinal Metals 55-77 Gravimetric and Volumetric Processes Applied to Sanitary Water Analysis 79-89 Medicinal Organic Compounds 91-118 Physiological Chemistry, including Digestion, Blood, AND Milk 119-144 The Analysis of Normal and Morbid Urine 145-164 The Toxicology of the Common Irritant and Neurotic Poisons 165-198 Tables of Symbols, Valencies and Atomic Weights, and OF Equivalent Weights and Measures 199-200 Index . . 201-207 (vii) EXERCISES IN LABORATORY MANIPULATION AND CHEMICAL ACTION PRACTICAL MEASUREMENTS. I. Comparison of English and Metric Length Units. 1. Measure with the rule the length of several objects in both systems. 2. Familiarize the length of a centimeter and decimeter. 3. Note the number of centimeters in an inch. 4. Measure the area of a book cover in square inches and square centimeters. Express each of these areas in decimal parts of a square yard and a square meter^ respec- tively. Note how much more difficult it is to reduce the English than the metric measures. 2. Comparison of English and Metric Volume Units. 1. Using graduates, determine the capacity of the dif- ferent sized test tubes, beakers and evaporating dishes in ounces, drachms and cubic centimeters (c.c). 2. Note the number of cubic centimeters in a drachm. 3. Express the volume of your beaker in parts of a gallon and parts of a liter. Note the advantage of the metric system. 3. Comparison of English and Metric Weight Units. 1. Weigh out 5 grains of calomel, 5 grains of antifebrin and -^ of a grain of atropine hydrosulphate. 2. Note the number of grains required to balance a gram weight. 4 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 3. Weigh out I gram (g.) and 10 grams of salt. Remem- ber the quantity by noting how large a coin each quantity might just be piled upon. 4. Comparison of English and Metric Temperature Units. 1. Examine the Fahrenheit and Centigrade thermometers, noting the comparative freezing and boiling points. 2. Note the number of degrees F. corresponding to i" C. 5. Reduction of English and Metric Units. By the fol- lowing rules,! reduce the prescriptions and temperatures written on the blackboard to corresponding units in the other system. Rule I. To Reduce Metric to English Weights and Mea- sures. Multiply the quantity in grams or cubic centimeters by 10, add y^, and the result is grains or minims (nearly). Rule II. To Reduce English to Metric Weights and Measures. Divide the quantity in grains or minims by 10, subtract J^, and the result is grams or cubic centimeters (nearly). Rule III. To Reduce Fahrenheit to Centigrade Tempera- tures. Subtract 32, and multiply by f. Rule IV. To Reduce Centigrade to Fahrenheit Tempera- tures. Multiply by f, and add 32. NoTH I. Rules I and II are merely approximate, but near enough for physicians' use, and are easily remembered and mentally applied. CONSTRUCTION OF APPARATUS. CONSTRUCTION OF APPARATUS. 6. Construction of a Stirring Rod. Select a solid or hol- low glass tube 8 inches long. Hold the ends in a Bunsen ilame until the ends are melted together and neatly rounded. When cool, cap one end with a small piece of rubber tubing, to protect a glass beaker while stirring. 7. Construction of a Dropper. Select a piece of glass tubing 8 inches long, that will snugly fit the rubber nipple. Heat this tube at the center, with constant rotation over the Bunsen flame. When hot, draw it out an inch and a half. With a file, scratch it in the center and break it. If it has been evenly drawn, two droppers may be made. Lay one- half aside for experiment 8 {6). Take the other and, heat- ing the larger end very hot, stamp it quickly on a cool metal surface, thus forming a ring to hold the rubber cap. 8. Construction of a Generator. Select a 7-inch test tube and fit it with a i-hole and a 2-hole rubber stopper. This apparatus is referred to throughout the laboratory course as a " generaior." It is used for a variety of purposes, and may have two kinds of fixtures, referred to as (a) a delivery tube, and (iJ) an ignition jet. A thistle tube added is often useful. Examine these generators on the demon- strator's desk. Fit one up, and keep it constantly ready for use. (fl) Construction of a Delivery Tube. Select a glass tube a foot long. One-third the distance from one end heat it evenly, with constant rotation, over a fish-tail burner. When hot, remove the tube from the flame and bend it into syphon 6 LABORATORY MANUAL OF MEDICAL CHEMISTRY. form. The bend should be uniform, with no "buckle" in the glass. (J>) Construction of an Ignition Jet. Round the ends of the tube saved in experiment 7, in the flame. When cool, insert it in the 2-hole cork for an ignition jet. (f) Construction of a Thistle Tube. Select a glass tube 8 inches long. Heat one end very hot until closed with a bead of white-hot glass. Remove it from the flame and gently blow a bulb as large as a small marble. Next heat, with a blowpipe, a spot on the very top of the bulb until white-hot. Remove it from the flame and blow a quick, strong blast. The top will be blown out. Clean the edges from fine glass and round them by heating in the flame. This makes a small but serviceable thistle tube. Insert it in the cork by the side of the ignition jet, as in the "model" generator. 9. To Mend a Test Tube. Heat the lower portion of a bottomless test tube, and when hot draw it together with a glass rod and draw away superfluous glass. Reheat the end, remove from the flame, and gently blow it into shape. The bottom should be round and quite thin, to stand heating. To mend a large tube or generator, it will be necessary to use the bellows and blast lamp. CHEMICAL MANIPULATION. 10. Solution. Dissolve about i g. of common salt, NaCl, in a test tube of water, and preserve the solution. 11. Evaporation. Put a few drops of the above solution in a porcelain evaporating dish. Place this on a piece of CHEMICAL MANIPULATION. 7 wire gauze or asbestos board and heat it over a non-luminous Bunsen flame until dry. Water and salt can thus be sepa- rated by heat. 12. Precipitation. Take half the remaining salt solution. Add nearly an equal quantity of silver nitrate, AgNOg. A white precipitate falls. Set the tube aside for the next experiment. 13. Filtration. Prepare a funnel, filter paper, ring-stand, and beaker, as on the demonstration desk. Shake the tube prepared in the last experiment, and pour the contents down a glass stirring-rod into the filter. Reject the solution caught in the beaker below. Save the precipitate (ppt. ). 14. Decantation. Precipitate the remainder of the salt solution left in experiment 12, by silver nitrate, AgNOj. Heat the tube, and set it aside until the precipitate settles, then pour off or decant the solution, leaving the ppt. 15. Reduction. Bore a shallow hole in a piece of char- coal. In this place all the white ppt. formed in the above experiments. Turn down the Bunsen burner to a 2-inch luminous flame. Rest the blowpipe tip on the top of the burner. Blow gently and continuously from the cheeks, and heat the precipitate with the fine blowpipe flame until minute beads of metallic silver appear. 16. Distillation. In a generator place a salt solution colored with indigo. Boil it, and condense the steam by letting the delivery tube dip in a test tube cooled by being immersed in a beaker of water. Taste the distillate, and note that both the color and salt have been removed. 17. Dialysis. Make a small dialyzer by tying tightly a fresh, thin sausage skin over the mouth of a bottomless test 8 LABORATORY MANUAL OF MEDICAL CHEMISTRY. tube. In this place a mixture of salt, starch and water. Set this, skin down, in a beaker of water. After some time take some of the water from the beaker and add silver nitrate. A white precipitate indicates salt, which has pene- trated the membrane. The water in the beaker is not colored blue by tincture of iodine. The milky liquid in the dialyzer turns blue with iodine, showing that the starch has not passed through the membrane. i8. Electrolysis. Prepare a solution of copper sulphate, CUSO4, in a beaker. Attach a clean silver coin to the nega- tive electrode of the dynamo or battery. Suspend the coin and the positive wire in the liquid without their touching. Examine the coin after a few moments. LABORATORY QUESTIONS. 1. Is the force which holds salt in solution a weak or strong one? 2. How could you separate salt and water? 3. What is a filtrate? What is a precipitate? 4. How may a precipitate be washed on the filter paper? 5. How may a precipitate be washed by decantation? 6. What is a distillate? 7. Can silver be dissolved in water? 8. Where did the copper come from which plated the coin? 9. Why was the coin hung on the negative electrode? 10. What class of bodies passes through animal membranes ? CHEMICAL CHANGE. CHEMICAL CHANGE. 19. A Mechanical Mixture. Make a mixture in any proportion of fine iron filings and flowers of sulphur. Grind them together in a mortar. This is a mixture, because the iron and sulphur can be separated by mechanical means. (a) From a small portion of the mixture remove the iron by the use of a magnet. The sulphur will be left. (J>) To another small portion add carbon disulphide, CS2. Shake it up. The sulphur will be dissolved. The iron will remain. Pour the solutjon into an evaporating dish and set it aside. The sulphur will be left in crystals when the CS2 evaporates. 20. A Chemical Compound. Several students together weigh out 56 g. of iron filings and 32 g. of flowers of sulphur. Mix and grind them thoroughly in a mortar. Take an inch of this mixture .in a test tube. Heat it until the red-hot glow diffuses itself throughout the mass. Cool, break the test tube, and grind the mass in a clean mortar. This substance is now a chemical compound, FeS, because the iron and sulphur cannot be separated by mechanical, but only by chemical means. (a) From a portion of the compound try to separate the iron by a magnet. It is all slightly magnetic, but no iron can be separated and sulphur left. {b) With another portion of the compound try to dissolve out the sulphur by carbon disulphide, CSj. If the work has been carefully performed, no sulphur will be found on evapo- rating the CS2. 10 LABORATORY MANUAL OF MEDICAL CHEMISTRY. Thus iron and sulphur can unite only in definite propor- tion by weight, as 56 parts of iron with 32 parts of sulphur. Any extra iron or sulphur will remain as a mixture. EVIDENCES OF A CHEMICAL CHANGE. 21. An Elevation of Temperature. Powder separately a pinch of sugar and a pinch of potassium chlorate, KCIO3. Mix them on an iron plate, and touch the mixture cautiously with a drop of sulphuric acid, H2SO4. Brilliant deflagration ensues. In general, every chemical combination is attended by an elevation of temperature. 22. The Formation of a Precipitate. Make a solution of mercuric chloride, HgCl2, by dissolving a little of the salt in water. Do the same with a little tin chloride, SnCl2. The solutions should be clear. Mix the two, and a white ppt. falls, turning black, and sometimes by heating and stirring mer- cury will separate. The formation of a ppt. from two clear solutions indicates a chemical change, resulting in the forma- tion or liberation of a new insoluble substance. 23. The Evolution of a Gas. Cover marble dust with water in a test tube and add a little dilute hydrochloric acid, HCl. A gas is given off. The evolution of a gas from two solid or liquid bodies indicates a chemical change, resulting in the formation of a new gaseous substance. 24. A Change in Electrical Condition. Fill a beaker one- third full of water and add 20 drops of sulphuric acid, H2SO4. Connect a zinc and a copper wire to a delicate galvanometer. Dip the wires, without touching each other, into the beaker of diluted acid. The acid forms a. chemical compound with the zinc, and a current of electricity flows through the gal- CHEMICAL CHANGE. 1 1 vanometer from the copper to the zinc, as indicated by the movement of the needle. In general, it is supposed that every chemical change is attended by a change of electrical condition. 25. A Change of Color. Touch a crystal of sugar of lead, lead acetate, Pb(C2H302)2, with a drop of a potassium iodide, KI, solution. A new yellow-colored compound is formed. A change of color is often an evidence of a chemical change, resulting in the formation of a new and differently colored substance. 26. A Change of Reaction. Place in a beaker a little dilute hydrochloric acid, HCl. Add a small piece of blue litmus paper. The paper turns red and the: solution tastes sour. Add some sodium hydroxide, NaOH, solution, with constant stirring until the paper turns faintly blue. The solution loses its acidity, and tastes salty. A change in the reaction of a solution toward litmus paper indicates a chemical change, resulting in the union of an acid and an alkali to form a new compound. SOME CAUSES INDUCING A CHEMICAL CHANGE. 27. Heat. Place a few grains of sugar in the bottom of a test tube. Heat gently. Note the carbon remaining and the water on the sides of the tube. Heat in general favors chemical change by increasing the molecular motion, destroy- ing old and encouraging new combinations. 28. Light. Make a solution of salt, NaCl, and add some silver nitrate solution, AgNOs- A white ppt. falls. Filter the solution. Place the filter paper, with its white silver chloride, in the sun. It soon turns bluish black, from a chemical change due to the energy of the sunlight. 12 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 29. Electricity. Wet a piece of filter paper in starch mucilage test solution (see 57, Note i). Lay this on a metal plate. Firmly press the wire from the negative electrode of the dynamo or battery on the plate, and with the wire from the positive electrode write slowly upon the paper. A blue line is traced. The color is due to a chemical change, caused by electric energy. 30. Mechanical Energy. Notice that dry quicklime, CaO, and dry ammonium chloride, NH4CI, have no odor. Mix a little of each very loosely in a mortar. They have no odor. Press and rub them firmly with a pestle. An odor of ammo- nia is evolved, and the mixture grows moist. The intimate union caused by pressure has induced a chemical change. 31. Solution. Mix in a beaker some dry powdered sodium bicarbonate, NaHCOs, and powdered tartaric acid, H2(H4C408) . No action ensues. Add water. Gas is evolved. In general, solution favors chemical action by separating atoms and molecules, and bringing them into intimate association. LABORATORY QUESTIONS. 1. What is meant by an element; an atom; a molecule; valence, and bond? 2. What is a mechanical mixture? A chemical compound? 3. Name several mechanical mixtures and chemical compounds in nature. 4. What is meant by the indestructibiUty of matter ? 5. What is the law of definite proportion? How is it proved? 6. What is the law of multiple proportion ? 7. In the manufacture of FeS (20), what was the most striking evi- dence of a chemical change? 8. When ammonium nitrate and sodium sulphate are dissolved the solution becomes cold. What evidence is there of a chemical change? 9. Is ice cream frozen by a chemical change? CHEMICAL CHANGE. 1 3 10. From what force is most of the energy of the world obtained ? 11. What power unites chemical elements? How does it seem related to electricity or magnetism. 12. When a solid compound is dissolved, are the atoms separated? 13. Is solution a chemical change? 14. What is meant by atomic and molecular weight? 15. What is the molecular weight of one of the most complex sub- stances known, provided its exact formula be CeooHoeoNisiFeSsOno (oxyhaemaglobin) ? 16. What is a graphic formula ? LABORATORY EXERCISES IN THE MORE IMPORTANT NON-METALLIC ELEMENTS HYDROGEN. 1 7 HYDROGEN. Useful Data : Atomic Symbol, H ; Molecular Formula, Ha ; Valence, i ; Atomic Weight, i ; Electro-positive ; i liter at 0° C. and 760 m.m. weighs .0896 g. 32. Preparation from Water by Sodium. Fill a test tube one-third full of water. Drop into it a small piece of metallic sodium, Na. Quickly apply a flame to the mouth of the tube. H burns with a pale blue flame. When the action ceases, test the remaining liquid by the touch, taste and action on red litmus paper. Evaporate the solution. Ex- amine the salt remaining, which is sodium hydrate, NaOH, or solid caustic soda. + - + + - + Reaction— H20+Na=NaOH+H. 33. Preparation from Hydrochloric Acid by Zinc. In a generator ^ put lo g. of granulated zinc, Zn, covered with water. Add a little hydrochloric acid, HCl. Collect several test tubes of H by the displacement of water, as illustrated on the demonstration desk. Light the H in one test tube. Transfer the H from one test tube to another, remembering that it is lighter than air. Prove its presence in the last tube by igniting it. After the action in the generator entirely ceases, pick out the Zn remaining, and return it to the container. Filter the solution. Evaporate it to a small bulk, and set aside to cool. Crystals of zinc chloride, ZnClg, will separate. + +- + +- Reaction— Zn+2HCl=H2+ZnCl2. 1 8 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 34. Preparation by Electrolysis. Fill a beaker with water and add 15 drops of sulphuric acid, H2SO4. In this invert a test tube filled with water. Keep the two electrodes from the dynamo or battery separated, and insert them in the beaker. Collect the tube half full of the gas arising from the negative electrode, from which the greater number of bubbles arises. Cover the tube with the thumb. Invert it, apply a flame to the gas, and prove that it burns like H. + - + - Reaction — H20=H2+0. LABORATORY QUESTIONS. 1. Name a number of substances which contain H. 2. Can H be prepared from any substance which does not contain it? 3. How many atoms of H in two molecules? 4. What were the qualities of the solution remaining in 32 ? 5. How much will 120 c.c. of H weigh under standard conditions? 6. How do the experiments prove H electro-positive? 7. From the experiments, which would one judge more strongly electro-positive, H or Zn? 8. How much H is contained in a pound of water, H2O? In a pound of HCl? 9. Is H soluble in water ? 10. How does the H flame differ from that of illuminating gas? 11. Does H in the experiments act like a metal? 12. Were all these experiments examples of chemical change? Why? 13. Were any materials used in these experiments destroyed? 14. Would Zn and H be likely to form a stable chemical compound? Why? CHLORINE. 1 9 CHLORINE. Useful Data: Atomic Symbol, CI; Molecular Formula, CI2 ; Valence, i ; Atomic Weight, 35.5 ; Electro-negative ; i liter at 0° C. and 760 m.m. weighs 3.173 g. 35. Preparation from Hydrochloric Acid by Manganese Dioxide. Place a little manganese dioxide, Mn02, in a generator. Add hydrochloric acid, HCl, and warm gently. Note the color, odor and specific gravity of the gasi as it rises in the tube. Wash the material quickly into the sewer. , , , _j Reaction— MnO, + 4HCI := MnCl, + 2H,0 + 2CI,. 36. Preparation from a Chloride. In a generator mix two parts of sodium chloride, NaCl, and one part of man- ganese dioxide. Add some sulphuric acid, H2SO4, and gently warm. Pass the yellow gas into half a test tube of water, thus forming chlorine water. In a part of this a scale of gold leaf will dissolve. Set the rest away in the sun, and notice after some hours the change in color, odor, taste and action on litmus paper. It turns to hydro- chloric acid, HCl. Reaction — + -+- +- +- +- +-- (i) 2NaCl + Mn02 + 2H2 SO4 =Na2S04 + MnSO* -j-aHaO-f^a. Or it may be written in two stages, thus : [ +- + - + - + - I Mrsi. 4NaCl + 2H2 SO4 = 2Na2 SO4 + 4lfa. {2)4 +_+_ _!-_ +_ _ ( Second. 4HCI + Mn02 = MnCU + 2H2 O + 2O1, as above. Note i. Care should be taken not to allow any considerable amount of CI to escape in the laboratory, When inhaled it produces great irritation. 20 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 37. Preparation from Bleaching Powder by an Acid. Place some bleaching powder in a generator. Add dilute acid. Hold a little damp litmus paper, colored cloth, red ink writing or other organic colors in the gas. The colors are bleached. The color does not return when they are moistened with ammonia water, NH^OH. 38. Preparation from Hydrochloric Acid by Electrolysis. Put some hydrochloric acid into a test tube. Insert the carbon electrodes of the dynamo or battery, keeping them separate. Note that, as before, H is liberated from the negative electrode. CI is liberated from the positive elec- trode, and dissolves in the water, forming yellow chlorine water. Note the odor of chlorine. + - + — Reaction— HCl = H + CI. 39. Preparation of Chlorine Water. Place a few crystals of potassium chlorate, KCIO3, in a test tube. Add 3 c.c. of hydrochloric acid. Warm until CI escapes freely, then fill the tube with water. This is the most convenient way to prepare a little fresh chlorine water, so frequently needed in testing. The salt and acid present do not interfere with its use. Reaction— 2HCI + 2KCIO3 = 2KCI + 2H2O + CljO^ + 2CI. COMPOUNDS OF CHLORINE AND HYDROGEN. Hydrochloric Acid (Muriatic Acid), HCl. 40. Preparation by Synthesis. Fill two test tubes, one with H and the other with CI, from the generators in the hood. Join their mouths. Invert several times, to mix. Cautiously open the mixture to a flame. After combina- tion, shake the gas in the two tubes with 5 c.c. of water. CHLORINE. 21 Test this by taste and blue litmus and set it aside, and later prove the presence of HCl by 43. + - +- Reaction— H + CI = HCl. 41. Preparation from a Chloride by Sulphuric Acid. In a generator cover 5 g. of sodium chloride with sulphuric acid, H2SO4. Warm, and cautiously note the odor. Hold moist blue litmus paper in the gas until it changes to red. Pass the gas through 10 c.c. of water. A solution of HCl forms, which is the muriatic acid of commerce and medicine. Try to boil the HCl from the solution. + - + - + - +- Reaction— 2NaCl + H2SO4 = Na2S04 + 2HCI. 42. Formation of Chlorides. In an evaporating dish put 15 c.c. of HCl. Add a gram of iron filings and gently warm for ten minutes. Fill the dish with water. Filter, evaporate the clear filtrate to a small bulk, and set aside. After a few moments fine crystals of green ferrous chloride, FeCl2, will form. + +- + - + Reaction— Fe-f^2HCl = FeCl2+H2. 43. The Silver Nitrate Test for Chlorides. Silver nitrate, AgNOj, added to a solution of chlorides, precipitates white silver chloride, AgCl, which is soluble in ammonia but insoluble in nitric acid. 44. Detonation of Chlorates. Place in an iron mortar two small crystals of potassium chlorate, KCIO3. Add a few grains of dry sugar or flowers of sulphur. Very gently rub them to a powder and strike the mixture with the pestle. A sharp explosion will occur. 45. The Ignition Test for Chlorates. A solution of potas- sium chlorate, KCIO3, yields no precipitate with silver nitrate, 22 LABORATORY MANUAL OF MEDICAL CHEMISTRY. AgNOg. Ignite a crystal of KCIO3 on a piece of glass or porcelaini. Cool, and dissolve the residue. The chlorate has turned to a chloride, and its solution yields the test for chlorides with silver nitrate, 43. LABORATORY QUESTIONS. 1. Name several substances that contain chlorine. 2. Are all substances containing chlorine chlorides? 3. When is a solution neutral to litmus? 4. How many substances can you find that turn red litmus blue? Blue litmus red? 5. How do the experiments prove chlorine electro-negative? 6. Write the reaction which chlorine water undergoes in the sun ? 7. What two compounds form the pure hydrochloric acid of commerce ? 8. Why is HCl called "spirit of salt"? What are muriates? 9. How much CI in one pound of salt ? 10. How much NaCl is required to manufacture one pound of HCl ? 11. Outline a good method for bleaching white cloth? 12. How would one proceed to disinfect a room with chlorine? 13. If CI unites with water, to what substance is the bleaching due? 14. Write the symbols for K, Cu, Au, Sn, Hg, AI, Mn, Ca, Ba and NH4 chlorides. 15. How might all these chlorides be formed? 16. What chloride was formed in the preparation of H ? 17. What chlorides were formed in the preparation of CI ? 18. What results when chlorates are triturated with organic sub- stances ? 19. To what acid is sodium chloride related? 20. To what acid is potassium chlorate related? Note i. See the reaction occurring here, 55, where it will be noted that oxygen is at the same time liberated. BROMINE. 23 BROMINE. Useful Data: Atomic Symbol, Br; Molecular Formula, Bra; Valence, i ; Atomic Weight, 80 ; Electro-negative. 46. Preparation from a Bromide. In a generator, mix three well-powdered crystals of potassium bromide, KBr, with an equal quantity of manganese dioxide, MnOg. Add sulphuric acid, H2SO4, and warm gently. Note the color, odor, specific gravity and bleaching qualities of the gas as it rises in the tube. If possible, pass the gas through water, forming bromine water. Set this aside for future use. Reaction — ' 2KBr+Mn02+2H2S04=K2SOi+MnS04+2H20+Br2. COMPOUNDS OF BROMINE AND HYDROGEN. Hydrobromic Acid, HBr. 47. Preparation from a Bromide by Sulphuric Acid. Place a crystal of KBr in an evaporating dish. Add a few drops of sulphuric acid, H2SO4, and warm. Note the white vapors of HBr mixed with some yellowish brown vapors of free bromine. Reaction— 2KBr-fH2S04= K2S04+2HBr, and aHBr+HgSOi S02+2H20+Br2. 48. Preparation from a Bromide by Hydrogen Sulphide. Take a test tube of bromine water, containing a globule of carbon disulphide, CS2. Pass through this hydrogen sul- phide, H2S, from the generator in the hood. Shake well, 24 LABORATORY MANUAL OF MEDICAL CHEMISTRY. and pour off the clear solution containing HBr. Preserve this for the following tests. Reaction— 2Br + HjS = aHBr + S. 49. Tests for Bromides. 1. The Chlorine Test. To a solution of bromides add a globule of carbon disulphide, CSg (chloroform may be used). Next add some chlorine water and shake. Free Br is liber- ated, and tinges the CS2 yellow to brownish red, according to the amount present. 2. The Silver Nitrate Test. AgNOa ppts. in solutions of bromides yellowish white, AgBr. Insoluble in nitric acid, HNO3, and sparingly soluble in dilute ammonium hydroxide, NH4OH. 3. The Starch-bromide Test. When starch water is added to a solution of bromides and Br liberated by a few drops of chlorine water, yellow starch bromide is formed. IODINE. Useful Data : Atomic Symbol, I ; Molecular Formular, I2 ; Va- lenge, i; Atomic Weight, 127 ; Electro-negative. 50. Physical Properties. Note the characteristics of metallic iodine. Heat a crystal in a dry test tube. Notice the color of the vapor, and the minute crystals sublimed on the side of the tube. Try to dissolve the crystals in water. Next try alcohol. Next try carbon disulphide or chloroform, and note the differences. IODINE. 25 50 1-2. Preparation from an Iodide. In a generator mix a crystal of KI with manganese dioxide, MnOg. Add sulphuric acid and warm gently. Note the violet vapors of iodine as they rise in the tube. Reaction — 2KI + MnOj + 2H2SO4 = MnSO^ + KgSO^ + 2H2O + Ig. COMPOUNDS OF IODINE AND HYDROGEN. Hydriodic Acid, HI. 51. Preparation from Iodine. Prepare a test tube con- taining a globule of CSg, a crystal of I and some water. From the generator in the hood pass hydrogen sulphide gas, HjS, through the mixture. Hydriodic acid remains in solution. Save this solution for the following tests. Reaction— 2I + HgS^ 2HI + S. 52. Tests for Iodides. 1. The Chlorine Test. To a solution of iodides add a globule of carbon disulphide, CSg, or chloroform. Next add chlorine water, and shake. Free I is liberated and colors the globule violet to black, depending upon the amount present. 2. The Silver Nitrate Test. AgNOg, added to a solution of iodides, ppts. light yellow, Agl ; insoluble in nitric acid, HNO3, and sparingly soluble in dilute ammonium hy- droxide, NH4OH. 3. The Starch-iodide Test. When starch water is added to a solution of iodides, and free iodine liberated by a few drops of chlorine water, blue starch iodide is formed. 26 LABORATORY MANUAL OF MEDICAL CHEMISTRY. FLUORINE. Useful Data : Atomic Symbol, F ; Valence, i ; Atomic Weight, 19 ; Electro-negative. COMPOUNDS OF FLUORINE AND HYDROGEN. Hydrofluoric Acid, HF. 53. Preparation from a Fluoride. First cover one side of a glass plate with an even coating of paraffin, and with a sharp point scratch a name or design through the wax. Next moisten 10 g. of calcium fluoride, CaF2, with sul- phuric acid, H2SO4, in a lead dish. Warm this for a few seconds only. Note the irritating odors of HF which arise. Observe that they are acid to litmus paper. Place two match sticks across the top of the dish and cover it with the glass plate, wax side down. After half an hour remove the wax, and the design will be found etched in the glass. Reaction — CaFg + H2S04= CaSOi -f 2HF, and HF at- tacks the sihca of the glass thus : 4HF -)- Si02"SiF4+ 2H2O. LABORATORY QUESTIONS. 1. What per cent of KI is I? 2. What substances did you find would dissolve I? 3. Will I dissolve in KI solution? Try it. 4. Can iodine or iodides be detected by putting a drop of their solution on one's cuff? 5. Complete the equation NaBr + HgNOa = 6. Balance the equation KOH + 1 = KIO3 + KI + H2O. 7. What are the white fumes seen escaping in etching? 8. What becomes of the sulphur set free in experiments 48 and 51 ? OXYGEN. 27 9. How do you explain the change of color of the CSa globule in experiments 48 and 51? 10. Write the reaction for preparing I from Fel2. 11. Why is HCl a more useful acid than HF or HI? 12. Can you invent a method of preparing NaBr from Br and NaOH ? 13. Would Br be valuable for a disinfectant? Why? 14. How do the experiments prove Br, I and F electro-negative? 15. Arrange CI, Br, I and F in the order of their atomic weights. Is this the order of their chemical affinity ? 16. Taste the compounds of CI, Br and I, and assign a reason why Berzelius designated them as "halogens" (salt formers). 17. How many reasons can you assign for classifying CI, Br, I and F in one family? 18. Why cannot the hydrogen compounds of the halogens all be prepared by the action of sulphuric acid on their salts? 19. If the chlorine test be applied to a mixture of iodides and bromides, can these elements both be detected? Try it by adding small amounts of chlorine water and shaking until a large amount has been added. Note any characteristic colors in the globule. 20. Why could chlorides, bromides or iodides not be prescribed in solution with mercurous, silver or lead salts? 21. Compare the reactions for the preparation of CI, Br and I from their salts. What similarity is found? OXYGEN. Useful Data : Atomic Symbol, O ; Molecular Formula, O2 ; Valence, 2 ; Atomic Weight, 16 ; Electro-negative ; Specific Gravity, 1. 1056 (Air is i) ; Weight of i liter at 0° C and 760 m.m., 1.430 g. 54. Preparation for Mercuric Oxide. Place i g. of mer- curic oxide, HgO, in a small test tube. Heat it very hot and 28 LABORATORY MANUAL OF MEDICAL CHEMISTRY. insert a glowing match stick from time to time until it bursts into flame. Examine the sides of the t.t. Notice the change of color of the remaining oxide as it cools. Reaction — 2HgO = 2Hg -|- O2. 55. Preparation from Potassium Chlorate. Place 2 g. of potassium chlorate, KCIO3, in a test tube. Heat and test for O, as before. Continue to heat until the action ceases. Compare the taste of the residue with that of potassium chlorate. See 45. Reaction— 2KC103= 2KCI + 3O2. 56. Preparation from a Chlorate by Manganese Dioxide. Place 2g. of KCIO3 in a test tube and mix the crystals with a little manganese dioxide, Mn02 (fine dry sand is nearly as good). Heat and test for O, as before. Note that the gas is evolved with less heat and the Mn02 is unchanged. Reaction— 2KCIO3 + Mn02= MnOg + 2KCI + 3O2. OZONE— Molecular Formula, O3. 57. Preparation from Ether. Put a little ether in a test tube, and keep it away from flames. Over it hold a piece of damp starch test paperi. Heat the end of a glass rod very hot, and with it stir the ether. Ozone is formed, and turns the starch paper blue. 58. Preparation by Electricity. Notice the peculiar odor of ozone about a Holtz or other electric machine when in action. Detect its presence by holding starch test paper near the discharge. Note i. Starch test paper is made by boiling some thin starch water, dissolving in it a little KI, and in this dipping the paper. Any substance which will unite with the K will leave the I free to form blue starch iodide. OXYGEN. 29 COMPOUNDS OF OXYGEN AND HYDROGEN. Water, H2O. 59. Preparation by Synthesis. Prepare a hydrogen gen- erator with an ignition jet. In it cover Zn with water and add some sulphuric acid. After it has been acting some moments, and the air is all expelled, wrap it in a cloth and light the gas. Hold over the flame an inverted beaker kept cool, if necessary, by a damp cloth wrapped about it. Notice that the moisture gathers within the beaker. Reaction— H2+0=H20. 60. Composition of Water Proved by Analysis. Fill a beaker with water acidulated with sulphuric acid, H2SO4. Invert in this two test tubes filled with water. Insert in the beaker the electrodes of the dynamo or battery. Collect the gases in separate tubes until a quantity of each is obtained. Note that in the same time twice as much of one gas as the other is collected. Place the thumb under the mouth of each tube in succession, invert and test the one for hydrogen by igniting the gas, and the other for oxygen by the glowing of a smouldering match. Hydrogen Dioxide, H202. 61. Preparation from Barium Dioxide. Place 10 g. of barium dioxide, BaOg, in a beaker of cold water. Let it stand, and stir repeatedly for half an hour. Stir, and add slowly 20 c.c. of sulphuric acid, H2SO4. Let it stand, and finally filter or decant the clear solution. It contains H2O2. 62. The Chromate Test for Hydrogen Dioxide. Mix in a test tube i drop of potassium bichromate solution, 30 LABORATORY MANUAL OF MEDICAL CHEMISTRY. KgCraOT, 2 drops of H2SO4, and a little ether. Lastly, add a few drops of H2O2, and shake. The supernatant ether becomes blue from the presence of perchromic acid, H2Cr208, in solution. TWO CLASSES OF OXIDES. 63. I. Acidic Oxides. Place a piece of dry phosphorus in an evaporating dish. Ignite it and quickly cover it with a glass plate, admitting as little air as is necessary to totally consume it. When action ceases examine the white powder, phosphoric oxide, P2O5. It has no action on dry blue litmus paper. Add some water. The solution is acid toward litmus, and has a sour taste. Phosphoric acid, H3PO4, is formed. Electro-negative chemical ele- ments in general form oxides which unite with water, forming acids. Reaction— P2+05=P205;P205+3H20 = 2H3P04. 64. II. Basic Oxides. In an evaporating dish ignite a dry piece of metallic sodium. The white residue is sodium oxide, Na20. It has no action on dry red litmus paper. Add water. Sodium hydroxide, NaOH, is formed, which is alkaline to litmus, and has the biting taste and greasy touch of an alkali. Strong electro-positive elements form oxides, which unite with water, forming hydroxides. Reaction— 2Na+0 = NagO ; Na20+ HjO = 2NaOH. Most metallic oxides are insoluble in water. 65. The Union of Acids and Alkalies. Place in a beaker a solution of sodium hydroxide, NaOH. It turns red lit- mus paper blue. In another beaker put some dilute HCl. It turns blue litmus red. Add one mixture to the other OXYGEN. 31 until the resulting solution has no action on litmus. This is best accomplished by putting a few drops of phenol- phthalein in the acid and adding the alkali drop by drop until on stirring a faint pink color appears. Evaporate this neutral solution to dryness, and taste the crystals of sodium chloride, NaCl, remaining. Acids and alkalies always unite to form a salt and water. Reaction— NaOH + HC1= NaCl + HgO. LABORATORY QUESTIONS. 1. How much O can be obtained from 100 g. of KCIO3 ? From 100 g. of HgO ? 2. Can you invent a method of preparing HgO ? 3. Explain the action of MnOa in the preparation of O. 4. Why is O "smoky" when first prepared? 5. Why does ozone turn starch test paper blue and O not? 6. What is the compound formed when iron rusts? Give the formula. 7. We can ignite Na or Mg with a match. Why not coal or iron? 8. What is meant by oxidation ? What by reduction ? 9. Was P or Na acid or alkaline before oxidation? 10. Why is NaOH called sodium hydroxide? 11. What is the anhydride of an acid or alkali? 12. What is formed when an acid neutralizes an alkali? 13. Can you tell whether a compound is alkaline or not by its formula ? 14. Why is the amount of H collected in the electrolysis of water slightly greater than twice the O? 15. Why are O3 and H2O2 chemically active substances ? 16. Lime water, Ca(OH)2, turns green cloth yellow. How might the color be restored? 17. What should be done when acid is spilled on clothing? 18. What is an acid? What is a base? What is an alkali? What is a salt? 32 LABORATORY MANUAL OF MKDICAL CHEMISTRY. SULPHUR. Useful Data : Atomic Symbol, S ; Molecular Formula, S2 ; Va- lence, 2, 4 or 6 ; Atomic Weight, 32 ; Electro-negative. FORMS OF SULPHUR. 66. Precipitated Sulphur. To half a test tube of water, add I g. of flowers of sulphur and ^ g. of slaked lime. Boil and filter. Acidify the filtrate with HCl. Finely divided white sulphur is precipitated, called "milk of sulphur." 67. Crystalline Sulphur. Dissolve J^ g. of flowers of sul- phur in 5 c.c. of carbon disulphide, CS2. Set aside until evaporated, and examine the crystals. 68. Plastic Sulphur. Melt 10 g. of roll sulphur in an evaporating dish until it is thick and dark. Pour it slowly into a bowl of cold water. Chew the plastic sulphur. It quickly turns to the crystalline variety. 69. Test for Sulphur. Sulphur heated on a silver coin forms a black stain of silver sulphide, AggS, which may be dissolved by potassium cyanide, KCy, solution. COMPOUNDS OF SULPHUR AND HYDROGEN. Hydrogen Sulphide, HjS. 70. Preparation from Ferrous Sulphide. Prepare a gen- erator with an ignition jet. In it place several pieces of ferrous sulphide, FeS. Cover them with water and add a little sulphuric acid, HgSO^. Note the odor of the gas evolved, and its action on moist blue litmus paper. Ignite SULPHUR. 33 the jet. Note that the rotten-egg odor changes to that of burning matches. Cold porcelain, pressed down on the flame, has sulphur deposited on it. Reactions— FeS+H2S04= FeSO^+HaS. H2S+30 = H20+S02. 71. Formation of Metallic Sulphides. Prepare solutions of Pb, Cu and Sb salts in separate test tubes. Pass hydro- gen sulphide, H2S, through each. Precipitates of insoluble metallic sulphides are formed. Reaction— Pb(N03)2+H2S = PbS+2HN03. 72. Reducing Power of Hydrogen Sulphide. Pass H2S through a test tube of strong nitric acid, HNO3. Sulphur is deposited. Repeat the process, using acid solutions of ferric chloride, FeClg, and potassium chromate, KgCrO^. The colors change to those of reduced compounds. The hydro- gen unites with the O or CI in each case. Reaction— 2HN03+3H2S = 4H20+2NO+3S. 73. Tests for Sulphides. 1 . The Lead Acetate Test. Paper moistened with lead acetate solution, Pb(C2H302)2, turns black when held in HgS, or dipped into a solution of soluble sulphides. 2. The Sulphuric Acid Test. Sulphuric acid, H2SO4, on sulphides liberates H2S, as in 70. 3. The Silver Test. A solid sulphide, mixed with sodium carbonate, fused q-q. porcelain, cooled, placed on a bright silver coin and moistened with a drop of dilute HCl, yields a black stain of silver sulphide, Ag2S, soluble in potassium cyanide, KCy, solution. 34 LABORATORY MANUAL OF MEDICAL CHEMISTRY. COMPOUNDS OF SULPHUR AND OXYGEN. Sulphur Dioxide, SO2. 74. Preparation from Sulphur. Burn a little sulphur on the end of a match stick, and note the odor of SO2. Reaction — S+02= SO2. 75. Preparation from Sulphuric Acid by Copper. Place strong sulphuric acid, H2SO4, and copper turnings in a test tube. Heat until the white fumes of SOg arise. A moist wheat straw held in the gas will be bleached. Reaction— CU+2H2SO4 = CuSO^+aHgO+SOa. 76. Preparation from Sulphides. Heat on an iron plate a bit of lead sulphide, PbS (galena), or iron sulphide, FeSg (fool's gold). Odors of SO2 can be detected. Reaction— FeS24-202 = Fe+aSOg. Sulphurous Acid, H2SO3. 77. Preparation from Water and Sulphur Dioxide. In a generator place some sodium sulphite, Na2S03, and add a little hydrochloric acid. SO2 is evolved when strong acids are added to sulphites. Pass this gas through a test tube of water. Note the odor of the solution, its taste, and action on blue litmus paper. Save the solution for experiment 82. Reaction— SO2+H2O =H2S03. 78. Tests for Sulphites. I. The Acid Test. Treated with strong acids, sulphites liberate SOg. . 2. The Barium Test. BaCl2 added to a solution of sul- phites, ppts. white barium sulphite, BaSOg. Add HCl. SULPHUR. 35 BaSOg dissolves. Boil and filteri. Add HNO3 and heat. Any dissolved BaS03 is oxidized to insoluble barium sulphate, BaS04, and precipitated. Sulphuric Acid, H2SO4. 79. Dehydrating Effect. Use dilute sulphuric acid as ink, and write on paper. When dry warm over a flame and note the charred characters. Try other acids. Approximate Reaction — C6HioOg+H2SO, = Ce+5H20+H2S04. 80. Affinity for Water. In half a test tube of water pour concentrated sulphuric acid. Note the heat produced. In diluting H2SO4, acid should always be added to water, not water to acid. 81. Tests for Sulphates. I. The Barium Test. BaCl2 added to solutions of sulphates ppts. white BaS04^, insoluble in boiling HCl. 3. The Silver Test. Insoluble sulphates mixed with sodium carbonate, Na2C03, fused on charcoal, transferred to a bright silver coin, moistened with a drop of dilute acid, yield a black stain of AggS, soluble in potassium cyanide, KCy, solution. Compare 69 and 73-3. 82. Illustrative Preparation of Sulphuric Acid. Take a test tube of freshly prepared H2SO3, as in experiment 77. Test it for H2SO4 by 78-2, and set it away in an evaporating dish for some days, exposed to the air. No- tice that it loses its sulphurous odor, and the test shows Note i. There will nearly always be some BaS04 which is insoluble, and obscures the reaction, hence it is best to boil and filter. 36 LABORATORY MANUAL OF MEDICAL CHEMISTRY. H2SO4 increasingly present. This illustrates slow oxi- dation. Reaction— 2H2S03+02=2H2S04. 83. Formation of Sulphates. Neutralize dilute H2SO4 with NaOH. Evaporate the solution to a small bulk. Cool, and examine the crystals of sodium sulphate, Na2S04, or Glauber's Salt, formed. Note how iron and copper sulphates were formed in 70 and 75. Reaction— NaOH-|-H2S04=Na2S04+H20. 84. Tests for Thio-Sulphates— HjSgOg. 1. The Hydrochloric Acid Test. HCl, added to a solution of thio-sulphates (so called hyposulphites), ppts. S with evolution of SO2. 2. The Silver Chloride Test. Freshly prepared AgCl, 28, is dissolved by a few drops of strong sodium thio-sulphate solution, forming a very sweet syrup. Hence the use of "hypo" in photography to dissolve silver from photo- graphic plates. LABORATORY QUESTIONS. 1. Is sulphur soluble in water? 2. How could a solution of S be made ? 3. Describe the changes observed in melting sulphur. 4. Why do eggs discolor silver spoons? 5. Why is S deposited on porcelain in example 70 ? 6. Volcanoes eject H2S and SO2. Write the reaction showing why S is found in volcanic regions. 7. When H2S water stands in the air S is precipitated. Write the reaction, explaining why sulphur water often contains specks of solid sulphur. 8. Is H2S an acid ? What are its salts called ? 9. Complete this equation : H2S+CuS04= 10. Why are sulphide ores roasted in extracting the metals? NITROGEN. 37 11. What is fool's gold? 12. What are pyrites, and why so named? 13. What per cent of SO2 is S ? 14. Write the graphic formulas of HCl, H2O, SO2 and H2SO4. 15. Write the reactions for each of the tests for sulphates. 16. Analyze a piece of rubber, and report whether sulphides or sul- phates are present. 17. How much S in 100 g. of H2SO4? 18. What analogous compounds do S and O form ? 19. How might calcium sulphate, CaSOi, called gypsum, be prepared ? 20. Why should water not be poured into strong H2SO4? 21. Why are pocket match-safes not made of silver? 22. Which sulphides are soluble and which insoluble? NITROGEN. Useful Data : Atomic Symbol, N ; Molecular Formula, N2 ; Valence, 3 or 5 ; Atomic Weight, 14 ; Electro-negative ; Specific Gravity, 0.971 (Air is i). 85. Preparation from the Air. Place a piece of phos- phorus the size of a pea on a cork, which is covered with a piece of asbestosi. Float this in a bowl of water. Ignite the P and cover it quickly with a beaker, well pressed down, so that no heated air escapes. When the action ceases, and the white fumes are all gradually dissolved in the water, cover the beaker with the hand or a glass plate, invert it and set it out, with the water it contains, on the desk. Quickly thrust in a lighted match, and note that the flame is ex- NoTE I. Phosphorus is kept under water. It must not be handled with the fingers, nor objects contaminated by it thrown where spontaneous combustion would be likely to cause a conflagration. 38 LABORATORY MANUAL OF MEDICAL CHEMISTRY. tinguished at the surface of the gas. Observe what fractional volume is occupied by water which was formerly occupied by oxygen. Reaction— 2 P + O5 + N = PgOg + N. COMPOUNDS OF NITROGEN AND HYDROGEN. Ammonia, NH3. 86. Preparation from Ammonium Chloride. Gently heat in a generator a mixture of equal parts of ammonium chloride, NH4CI, and quicklime, CaO, taking care not to break the generator. Test the gas evolved by moist red litmus paper. Pass it for some moments through loc.c.of water, forming aqua ammonia. Reactions— 2 (N H^) Cl-f CaO = CaCl2+2 N H3+ HgO. NH3+H20 = NH40H. 87. Preparation from Animal Matter. In a test tube heat a few small pieces of hoof-parings from a blacksmith shop. Note the odor, and the power of the gas evolved to turn moist red litmus paper blue. 88. Formation of Ammonium Salts. Neutralize some ammonium hydroxide NH4OH, or aqua ammonia, by adding HCl. Notice the white smoky fumes of ammonium chloride that arise. Evaporate the solution to a small bulk, and cool. Examine the crystalline ammonium chloride. Reaction— NH4OH+HCI =NH4C1+H20. 89. Tests for Free Ammonia. 1. The Litmus Test. Moist red litmus paper is turned blue by the fumes of NH3. 2. The Ammonium Chloride Test. A glass rod dipped in HCl and held in NH3 yields white fumes of NH4CI. NITROGEN. 39 COMPOUNDS OF NITROGEN AND OXYGEN. Nitrous Oxide, N2O, Laughing Gas. 90. Preparation from Ammoaium Nitrate. Heat some ammonium nitrate (NH4)N03 in a generator. Nitrous oxide, NgO, is evolved. Catch several test tubes of the gas by the displacement of water, as in 33. Note that the gas has a sweetish odor. Insert a match, and prove it a supporter of combustion. Reaction— NH4N03 = N20+2H20. Nitric Oxide, N2O2, or NO. 91. Preparation from Nitric Acid by Copper. In a genera- tor gently warm a mixture of nitric acid, HNO3, and copper turnings. Note the color of the gas evolved. Catch a test tube full of this gas by the displacement of water. Note that it is colorless. Expose it to the air, and it turns reddish brown by uniting with oxygen to form N2O1, a reddish gas, one of the higher oxides of nitrogen. Reactions— 3Cu+8HN03=3Cu(N03)2+4H20+N202. N202+02 = N204. Nitrous Acid, HNO2. 92. Preparation of Nitrites from Nitrates. Heat some dry potassium nitrate, KNO3, in a test tube. Oxygen is evolved, and potassium nitrite remains. Save this for the following tests. Nitrates thus readily yield oxygen. Reaction— 2KN03 = 2KN02+02. 93. Tests for Nitrites. I. The Starch Test. Solutions of nitrites acidulated with H2SO4 turn blue a solution of starch-potassium-iodide mucilage (57, Note i). 40 LABORATORY MANUAL OF MEDICAL CHEMISTRY. Nitric Acid, HNO3. 94. Preparation from Sodium Nitrate. Heat in a genera- tor equal quantities of sodium nitrate, NaNOg, and HgSO^. Catch 10 drops of the distillate in a test tube cooled by being wrapped in a wet cloth or immersed in a beaker of water. Put a drop of the acid on the finger nail and notice the yellow stain. Observe that it is strongly acid. Reaction— 2NaN03+H2S04=Na2S04+2HN03. 95. Preparation of Gunpowder. Take i part of flowers of sulphur, 2 parts of powdered carbon, and 7 parts of potas- sium nitrate, KNO3, or nitre. Dampen the mass. Grind it in a mortar and dry it. When dry, pulverize and ignite small quantities on an iron plate. Approximate Reaction — S + 2KNO3 + 4C = K2S + 2CO2 + 2CO + N. 96. Preparation of Nitro-hydrochloric Acid. Make a mixture of i part HNO3 and 3 parts of HCl in a test tube. Warm, and note the odor of chlorine liberated. Dissolve in this a scale of gold. This acid with metals forms chlorides, and its power is due to the liberation of nascent chlorine. 97. Formation of Nitrates. Heat a bit of metallic silver with HNO3. Dilute with water, and use the silver nitrate, AgNOs, in testing for chlorides, as in 43. Note how cop- per nitrate was formed in 91. All nitrates are soluble in water. 98. Tests for Nitrates. I. The Ferrous Sulphate Test. Mix a little strong H2SO4 with a little fresh green ferrous sulphate solution, FeS04. Cool the mixture and overlay it, in a test tube, carefully with a solution containing nitrates. Tap the tube gently, and at NITROGEN. 41 the junction of the liquids a ring forms, usually brownish black, but sometimes violet, red or brown. 2. The Indigo Test. Indigo solution is decolorized by free nitric acid. 3. The Nitric Oxide Test. Any nitrate heated with HgSO^ and copper yields brownish red fumes of NO2. LABOEATORY QUESTIONS. 1. From the experiments, what per cent of air was shown to be N ? 2. What is the difference between ammonia and ammonium ? Aqua ammoniae and ammonium hydroxide? 3. What per cent of NH3 is H ? 4. How might thunder storms prepare ammonia in the atmosphere ? S- Why is nitro-hydrochloric acid called "aqua regia"? 6. Why is ammonia called "spirit of hartshorn " ? 7. What are spirits of ammonia ? 8. Write the equation when ammonia burns in air. 9. Is quicklime an allcali? Try it on moist red litmus paper. 10. Complete the equation KOH+(NH4)Cl= 11. In general, what distinctive difference is there in the action of H2S and HNO3 on compounds? 12. How many pounds of HNO3 can be prepared from a ton of NaNOa, or Chili saltpetre ? 13. Why are nitrates more easily prescribed than chlorides ? 14. Complete and balance the equation N20+C=C02 15. How might laughing gas be manufactured from nitric acid and ammonia ? Write the reactions in full. 16. Why is ammonia called the volatile alkali ? 17. What difference do you observe in the color of red litmus paper which has been dipped in a fixed alkali and dried, and that of a slip dipped in volatile alkali and dried? Explain. 42 LABORATORY MANUAL OF MEDICAL CHEMISTRY. PHOSPHORUS. Useful Data : Atomic Symbol, P ; Valence, i, 3 or 5 ; Atomic Weight, 31 ; Electro-negative. 99. Two Forms. (i) Waxy phosphorus should be handled with forceps. It emits a garlic odor and fumes luminous in the dark. Heat a small piece in a test tube and it burns, emitting white fumes of phosphoric pentoxide, P2O5. (a) Red or amorphous phosphorus is not luminous nor poisonous. It is an allotropic form of phosphorus. Heat a small amount cautiously in the bottom of a test tube. Some of it may then be removed on the end of a glass rod, and found to be ordinary waxy phosphorus. 100. Spontaneous Combustion. Dissolve a grain of waxy phos- phorus in a few drops of carbon disulphide, CS2- Pour the solution on a filter paper laid on a ring-stand. The paper will take fire after a few moments from the heat caused by the rapid oxidation of the phosphorus. loi. Reducing Properties. Add a solution of phosphorus in carbon disulphide, CSg, to a solution of copper sulphate, CUSO4. Shake the tube, and observe the formation of black metallic copper, copper phosphate, etc., due to the union of phosphorus with the oxygen of the CUSO4. Phosphoric Acid, H3PO4. 102. Preparation from Phosphorus. Place a small quantity of red phosphorus in an evaporating dish. Cover it with HNO3. Warm gently until the phosphorus is dissolved. Evapo- PHOSPHORUS. 43 rate the solution to a thick syrup. This is nearly pure phosphoric acid, H3PO4. Reaction— 3P+5HNO3+2H2O = 3H3PO4+5NO. 103. Preparation from Phosphoric Oxide. Refer to experi- ment 63, and note the method of formation from P2O5 and water. 104. Formation of Phosphates. Neutralize some H3PO4 with sodium hydroxide, NaOH. Evaporate the solution nearly to dryness, and examine the sodium phosphate, Na2HP04, formed. 105. Precipitation of Earthy Phosphates. Dissolve bone ash in dilute HCl. Render the solution alkaline with NH4OH, and note the precipitate of calcium phosphate, Ca3(P04)2. 106. Tests for Phosphates. 1. The Silver Test. AgNOs ppts. from neutral solutions of phosphates, yellow silver phosphate, Ag3P04, soluble in (NH4)0H and HNO3. 2. The Magnesium Test. A mixture, consisting of MgS04, (NH4)C1 and (NH4)OH ppts. white crystalline ammonio- magnesian phosphate, NH4MgP04. 3. The Molybdate Test. Ammonium molybdate, (NH4)2- M0O4, ppts. from solutions of phosphates, acidulated with HNO3, yellow ammonium phospho-molybdate (NH4)3P04- (Mo03)io, which increases on application of heat and on standing. 107. Tests for Hypophosphites. H3PO2.1 Note i. Phosphates are all insoluble in neutral solution, except those of the alkali bases. Hypophosphites are nearly all soluble, and are thus easily administered. 44 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 1. The Silver Test. AgNOj ppts. in neutral solutions of hypophosphites white silver hypophosphite, AgHgPOj, soon turning black from the reduction of metallic silver. 2. The Mercury Test. HgCl2 ppts. from solutions of hypophosphites white mercurous chloride (calomel, HgCl), when warmed with HCl. io8. Preparation of Sodium Pyrophosphate. Heat sodium phosphate, Na2HP04, in a test tube until the water of crys- tallization is given off. Gradually raise the tube to a red heat. Sodium pyrophosphate, Na^PgOY, is formed, which will differ from the original sodium phosphate by yielding a white ppt. with AgNOg. Reaction— 2Na2HP04=Na4P207+H20. LABORATORY QUESTIONS. 1. Is red phosphorus soluble in CS2 ? 2. Why should oils not be given in phosphorus poisoning? 3. What is spontaneous combustion? 4. What is the anhydride of phosphoric acid ? 5. Give the formulas of three acids that may be respectively neutral- ized by I, 2 and 3 atoms of Na. 6. Why is H3PO4 said to be tribasic? How many different sodium salts can it form ? Give their formulas and the reaction of each salt to litmus paper. 7. When earthy phosphates are in solution, what must be the reaction of the liquid ? 8. When would urine precipitate calcium phosphate in the bladder? 9. Give the graphic formula for H3PO4. 10. Why are hypophosphites useful in medicine? 11. Why is pyro-phosphoric acid so named? 12. Why are "hypo" and "ortho" phosphoric acids so named? 13. How may the latter two be distinguished by AgNOs? 14. Can there be one phosphate containing three distinct bases? 15. How much P is contained in a pound of bone ash, 85% of which is Ca3(P04)2, and the remainder CaCOs? 16. What are microcosmic salts? BORON. 45 BORON. Useful Data : Atomic Symbol, B ; Valence, 3 ; Atomic Weight, II ; Electro-negative. BORIC OR BORACIC ACID, H3BO3. 109. Preparation from Borax. Make a hot saturated solution of borax, NagB^O^. Add sulphuric acid until slightly acid. Cool, and collect the crystals of boric acid. Note that these crystals are soluble in glycerine. Reaction— Na2B407+H2S04+5H20 = Na2S04-f4H3B03. no. Borax Bead. Heat a small loop of platinum wire and dip it in powdered borax. Reheat and repeat the process until a glassy bead is obtained. Touch this with a tiny speck of cobalt compound and reheat. The bead becomes blue. Many metals may thus be distinguished by the color of the double salts formed with the borax bead. Ill, The Flame Test for Boric Acid. Boric acid, or borates moistened with H2SO4, dissolved in alcohol and the solution ignited, tinge the flame pale green. Best seen when the flame is gently blown by the breath. 46 LABORATORY MANUAL OF MEDICAL CHEMISTRY. CARBON. Useful Data : Atomic Symbol, C ; Valence, 4 ; Atomic Weight, 12 ; Electro-negative. 112. Forms of Carbon. Examine the prepared samples of graphite or plumbago, lamp black, bone black, char- coal, coke, gas carbon, peat, lignite, bituminous and anthra- cite coal. 113. Reduction by Carbon, Mix a small quantity of powdered charcoal and copper oxide in a test tube. Heat the tube very hot and reduced, reddish metallic copper ap- pears. Reaction— 2CuO+C = 2Cu-f COj. 114. Decoloration by Carbon. Arrange a funnel, fitted with a filter paper and half filled with bone black. Moisten this carbon with water, and through it filter a solution of any organic coloring matter, such as litmus, indigo or cochineal. The color is removed. Repeat with a mineral coloring matter, like potassium chromate, and observe that it is not removed. 115. Deodorization by Carbon. Saturate some water with H2S until it smells strongly of the gas. Filter this through bone black, and note that the odor is entirely removed. CARBON AND HYDROGEN. Methane, CH4. 116. Preparation from Sodium Acetate. Mix 2 g. of sodium acetate, Na(C2H302), which has been just pre- viously thoroughly dried, with 8 g. of NaOH and 2 g. of CARBON. 47 finely powdered quicklime, CaO. Heat this mixture very hot in a generator fitted with an ignition jet. Ignite the gas, which is methane, CH4, or marsh gas, the simplest of the many hydrocarbons, and the starting point in organic chemistry. Reaction— Na(C2H302)N+aOH = Na2C03+CH4. 117. Preparation of Illuminating Gas. Half fill a clay pipe bowl with powdered bituminous coal. Lute the top of the bowl with wet clay or a freshly mixed plaster paris paste. Heat the bowl very hot, and at last ignite the gas at the end of the stem. Note the water given off and the coal tar produced. The vapor will turn moist red litmus blue, showing the presence of ammonia. At the end of this "destructive distillation " open the pipe and examine the coke formed. CARBON AND OXYGEN. Carbon Monoxide, CO. 118. Preparation from Oxalic Acid. Put 5 grams of crystallized oxalic acid in a test tube and cover it with H2SO4. Strongly heat the tube. After a moment the escaping carbon monoxide, CO, may be ignited at the mouth of the tube, and burns with a characteristic blue flame. Reaction— C2H2O4+H2SO4 = H2SO4+CO2+H2O+CO. Carbon Dioxide, CO2. 119. Preparation by Combustion. Shake some lime water in a large test tube, and see that it remains clear. Hold a burning match in the tube until it is extinguished. 48 LABORATORY MANUAL OF MEDICAL CHEMISTRy. Close the tube with the thumb and shake it. Observe that the lime water becomes milky. Reaction — C-|-0 = CO2, and CO+2Ca(OH)2= CaCOg+HjO. 120. Preparation from Carbonates. In a generator cover small pieces of marble, CaCOg, with dilute HCl. Collect the heavy gas flowing from the delivery tube in a beaker, as though it were water. Insert a blazing match. Observe that it is extinguished at the surface of the gas. Pour some of this carbon dioxide, COg, into a tube containing lime water, shake, and confirm the presence of COg. Reaction— CaCOs+aHCl = COj+CaClj+HgO. 121. Preparation by Fermentation. At the close of the laboratory period place in a generator a dilute solution of molasses and a piece of yeast. Set it in a warm place and let the delivery tube dip in a test tube of clear lime water. The next day the lime water will be turbid from CO2 given off. Sugar. Alcohol. Reaction— CgHiaOg = 2(C2Hg)OH+2C02. 122. Detection of CO2 in the Breath. Blow the breath through a test tube of dilute lime water. The calcium carbonate, CaCOg, first formed will later redissolve. This precipitate will again fall when the solution is boiled, and the CO2, which holds the limestone in solution, escapes. 123. Preparation of Carbonic Acid. Pass CO2, from the generator prepared in the hood, through a test tube of water. The resulting solution turns blue litmus proper faintly red. Carbonic acid, HgCOg, is probably formed. CARBON. 49 When this is neutralized by NaOH, sodium carbonate, NagCOa, is produced. Reactions— CO2 + H2O = HgCOg. H2CO3 + 2NaOH = NajCOg + 2H2O. 124. Test for Carbonates. Any carbonate treated with HCl effervesces and liberates COg, which may be detected by passing the gas through lime water and noting any turbidity. CARBON AND SULPHUR. Carbon Disulphide, CSj. 125. Determination of Composition. Place a few drops of carbon disulphide in a bottle. Ignite it by inserting a glass rod heated to redness at the end. Note the odor of SO2 arising, proving the presence of sulphur. When action ceases, shake the contained gas with a little lime water. Observe the turbidity as a proof of the presence of C. Reaction— CS2 + 302= CO2 -\- 2SO2. 126. Solvent Power. Dissolve in CS2 small amounts of S, rubber, oil and paraffin. Remove a grease stain from cloth with CS2. CARBON AND NITROGEN. Hydrocyanic Acid, HCW, or HCy. 127. Preparation from a Cyanide. Place a pinch of po- tassium cyanide, KCy, in a test tube. Cover it with dilute H2SO4, warm, and note the odor of hydrocyanic acid, HCy, or Prussic acid, hke the odor of peach blooms. Reaction— 2KCy + H2S04= K2SO4 + 2HCy. 128. Preparation from a Ferrocyanide. In a generator place I g. of potassium ferrocyanide, K^FeCyg, and cover D 50 LABORATORY MANUAL OF MEDICAL CHEMISTRY. it with dilute HgSO^. Gently heat the generator and pass the gas formed into a test tube of water. Dilute hydro- cyanic acid, HCy, is formed. Save this for the following tests. Reaction — aK^FeCye + 6H2SO4 = 6KHSO4 + FejKgCye + 6HCy. 129. Tests for Cyanides. 1. The Silver Test. AgNOg ppts. from solutions of cyanides (except mercuric cyanide) white silver cyanide, AgCy, insoluble in dilute HNO3, soluble in NH4OH, recognized from silver chloride by evolving the odor of HCy when treated with strong HCl (see 43). 2. The Prussian Blue Test. To a solution of cyanides a few drops of KOH and FeS04 are added, and the mixture warmed, then two drops of FeClj are added. The whole is slightly acidulated with HCl, to dissolve ferrous and ferric hydroxides, when Prussian blue, Fe4- (FeCy6)3, will appear. 3. The Sulphocyanate Test {a). To a solution of cya- nides, add 2 drops of yellow ammonium sulphide, (NH4)2S2. Warm until colorless. Slightly acidulate with HCl, when a drop of FeClg will yield a blood-red colora- ation. This is the most delicate test. ((5) This test may be thus performed : Place in a beaker the material containing free HCy (if a cyanide, free HCy must be liberated by adding H2SO4). Cover the beaker with a glass on the under side of which are two drops of (NH4)2S2. With this the vapor forms ammonium sulphocyanate. Let the drop dry thoroughly. If HCy were present, the spot immediately turns blood- red when touched with a solution of FeCL. CARBON. 51 LABORATORY QUESTIONS. 1. How might cider vinegar be rendered white, like wine vinegar? 2. Write the reaction when iron ore, Fe203, is mixed with coke to obtain pig iron. 3. Write the reaction showing how H2S water was deodorized. 4. Write the reaction involved when "fire-damp" burns. 5. What is destructive distillation? 6. Write the reaction occurring when limestone is heated. 7. How is the CO2 of the breath formed ? 8. CO2 contains O. Why does it not support combustion? 9. The yeast plant uses neither alcohol nor CO2. For what purpose does it split the sugar into these substances? 10. Why is NaHCOa called bicarbonate or acid carbonate of soda? Is it acid? 11. What other substances beside CS2 will dissolve grease? 12. What is soda water, and how may it be prepared? 13. Name several elements that exist in allotropic forms. 14. Bone-black contains principally C, CaCOs, Ca3(P04)2, and tarry organic compounds. How may the C be obtained pure for filters? 15. If a ton of C is burned, what weight of CO2 is produced? 16. Does wood produce as much heat during decay as during com- bustion ? 17. Where was the C, now contained in coal, before the carbon- iferous age? 18. What causes beer to foam? 19. Explain the cause of a non-luminous Bunsen flame. 20. Why does burning alcohol, (C2H5)OH, not deposit soot? 21. Why do stoppers in KOH or NaOH bottles stick when left for some time ? Touch the neck of one with HCl, and observe the phe- nomena. 22. Of what are oyster shells composed? Test one for carbonates. 23. Why is KCy used in extracting gold? 24. Why was CaO used in the preparation of methane ? 25. How can a chloride be distinguished from a cyanide by AgNOs ? 26. Write the reaction showing how borax is a salt of tetraboric acid. 27. How may the NH3 given off in gas manufacture be utilized ? 28. What are the blue flames seen in a stove burning anthracite coal ? 29. Why do kettles have crusts deposited within them ? 52 LABORATORY MANUAL OF MEDICAL CHEMISTRY. TABLE I. Preliminary Tests for Common Acids. I. THE SUBSTANCE IS A SOLID. Place a small portion in a test tube, add i c.c. of con- centrated H2SO4, and warm gently. A. Efiervescence of a colorless gas. 1. Odorless. Carbonates, 124; Oxalates, 213. 2. Odor of rotten eggs. Sulphides, 73. 3. Odor of burning sulphur. Sulphites, 78 ; Thio- sulphates, 84. 4. Odor of peach blooms. Cyanides, 129. 5. Odor of vinegar. Acetates, 209. 6. Odor merely irritating. Nitrates, 98 ; Chlorides, 43 ; Fluorides, 53. B. Effervescence of a colored gas. 1. Violet color. Iodides, 52. 2. Yellowish brown color. Acrid odor. Bromides, 49. 3. Reddish brown color. Nitrous odor. Nitrites, 93. 4. Greenish yellow color. Chlorine odor, Hypo- chloriteg, 166. 5. Greenish yellow color. Chlorine odor, with de- tonation, Chlorates, 45. C. No action. Test for H2SO4, 81; H3PO4, 106; HNO3, 98, etc. II. THE SUBSTANCE IS A LIQUID. Neutralize a portion, evaporate it to dryness, and test the residue by I. TABLE II. 53 - OJ M „ Not Distinguished Oxalates, 213. Phosphates, 106. Borates, iii. Effervesce wit Acids. Carbonates, 124. Sulphites, 78. Char when Heate on Platinum. Tartrates, 214. Citrates, 215. II 1 i.5. K if re? a. &- .»>■ W M „ Char when Heate on Platinum. Tartrates, 214. Citrates, 215. Soluble in Borates, 11 Carbonates Citrates, 21 Insoluble Acid. Oxalates, 2 Tartrates 214. Sulphates 81. % .^1 OS ■ «i ►80 II no i 1^ «-l-"i P ? R. ^ ^ s- ft- w „ Chlor Cyani Brom lodid Sulph nHff 20 n« 3ft'gggE.s3-3 des, white, 43 des, white, 129 des, white, 49. 2s, yellow, 52. ides, black, 73. -^1 8. nates, white, tes, white, 213. hates, yellow, tes, white, 78. ites, white, 214. es, white,. 215. 1 b 1 '1 rn ■DO •o S-g » nS re & ." as s • re ■n "^ re s^ ■5'Q - re ■< S'tf 1 s 7S «^ ::: re ^ 5 R So ^3. ■ != - ^ e- OJ ^ ^ f _3" Nitrates, 98 Chlorates, 45. fi ! S£S.p]Sog w p „ B m ■ 13 • 3 B tn o !U n ^ ^ tP B ^ ■« E S 3 -, r, S g • ? — ft> !=i O • 2 2 o =; D. Q- s. ™ -^ P 03 ^ ^ rr* ' (1) H o > o p dUSd^- S3 g. ?-^ ST O- m o ^ I -g C!!^ c .-f, J 3 ^ CLTD ('^ p D" ri o G. f CA .&S 3 p "5 P O =1 (D ^ l=? rt. < (Tl fit ^ CL p y "" 1,3 g • a ff p* o ui p •-S si' o >-( P > o 2 3 X tl' 2 '-' "t p v; OK? THE MORE IMPORTANT REACTIONS ANALYTICAL AND SYNTHETICAL OF THE PRINCIPAL MEDICINAL METALS. THE METALS. 57 THE METALS. THE FIVE- METALLIC GROUPS. GROUP I. THE SILVER GROUP. Ag, Pb and Hg(ous). 130. Group Precipitant, HCl. Prepare a solution of the salts of each of these three metals in a separate test tube. Add a few drops of HCl to each, and note the white precipitates. GROUP II. THE COPPER GROUP. Hg(ic), Pb, Bi, Cu— As, Sb (Cd, Sn, Pt, Au). 131. Group Precipitant, HjS in Acid Solution. Prepare a solution of the salts of each of these six metals in sepa- rate test tubes. Add a drop of HCl to each; no precipi- tate falls. Pass H2S gas through each. Note the colors of the various sulphides precipitated. GROUP III. THE IRON GROUP. Fe, Al, Zn (Cr, Mn, Ni, Co). 132. Group Precipitant (NH4)2S in Alkaline Solution, Prepare a solution of the salts of each of these three metals in a separate test tube. Add first a few drops of NH4OH to each,- and follow with a solution of (NH4)2S. Note the color of the precipitated Fe and Zn sulphides and Al hydrate. GROUP IV. THE BARIUM GROUP. Ba, Ca (Sr). 133. Group Precipitant, (NH4)2C03 in Alkaline Solution. Pre- pare a solution of the salts of each of these two metals 58 LABORATORY MANUAL OF MEDICAL CHEMISTRY. in a separate test tube. Add first a few drops of NH^OH to each, and follow with (NH4)2C03. Note the white precipitates. GROUP V. THE POTASSIUM GROUP. Mg, K, Ka, NH4 (Li). 134. Group Precipitant, None. Prepare solutions of these salts, and note that they are precipitated by none of the preceding group reagents. LABORATORY EXERCISES. I. Test bottles Nos. i, 2, 3, 4 and 5, and determine to which group the one salt dissolved in each belongs. ANALYTICAL REACTIONS OF GROUP I. 135. Silver — Agf. — Confirm reactions, using Solution of AgNO^. 1. HCl ppts. white AgCl, soluble in NH4OH. 2. KgCrgOy ppts. red AggCrO^. 3. H2S ppts. black AgjS. 4. KI ppts. pale yellow Agl. 5. NH4OH ppts. black AgjO, soluble in excess. 6. NaOH ppts. brown AgaO, soluble in (NH4)OH. 7. Solids heated with NajCOj on charcoal, in the re- ducing blowpipe flame, yield bright metallic beads, solu- ble in HNO3. 136. Lead, Pb". — Confirm reactions, using solution of Fb{_C^H^O^\. 1. HCl ppts. white PbClg, soluble in hot water and strong acids. 2. KI ppts. bright yellow Pbl2. 3. H2S ppts. black PbS. 4. KOH ppts. white Pb(0H)2, soluble in excess. THE METALS. 59 5. H2SO4 ppts. white PbS04, in dilute solutions only on standing. 6. KgCrjOT ppts. bright yellow PbCrO^. 7. Solids heated with NajCOg on charcoal, in the re- ducing blowpipe flame, yield soft metallic beads, with a yellow incrustation of PbO on the charcoal. The beads are soluble in HNO3. 137. Mercury (ous),Hg" — Confirm reactions, using solution of HgNO^. 1. HCl ppts. white HgCl (calomel), blackened by (NHi)OH. 2. KI ppts. dark green Hgl. 3. HgS or (NH4)2S ppts. black HggS. 4. Bright copper, in a slightly acid solution, becomes covered with metallic Hg, made bright by rubbing. LABORATORY QUESTIONS ON GROUP I. 1. Name the three insoluble chlorides. 2. What is a group precipitant? 3. How may silver nitrate be prepared? 4. How might one prepare lead acetate? 5. Why are silver salts used in photography? 6. Write the reaction 2AgN03 + 2NaOH = 7. Why is MgS04, or Epsom salt, an antidote in acute lead poisoning? 8. With what acid could Pb be dissolved from an ore ? 9. Why does lead acetate added to tap water give a precipitate? 10. Try to dissolve silver stains from the hand with KCy. 11. Why is AgNOa an ingredient of indelible ink? 12. Make a solution containing salts of Group I. Turn to the analyti- cal Table, III, page 77, and separate each metal. 13. Test solutions i, 2, 3, 4, 5 and 6 on the side table, reporting the metals found in each. 14. Test ores A, B, C and D for first group metals by powdering the ore, dissolving in HNO3, filtering, neutralizing excess of acid, and pro- ceding by Table III. Report the metals found. 6o LABORATORY MANUAL OF MEDICAL CHEMISTRY. ANALYTICAL REACTIONS OF GROUP II. 138. Mercury(ic), Hg". — Confirm reactions, using solution of HgCl^. 1. HgS ppts. ultimately black HgS, soluble in aqua regia, insoluble in (NH4)2S. 2. SnCl2 ppts. white HgCl, turning black with liberation of minute globules of Hg, which appear when the ppt. is boiled with HCl and rubbed. 3. (NH4)0H ppts. white (NH2)HgCl (white precipitate). 4. KOH ppts. yellow HgO. 5. KI ppts. bright red Hgl2. 6. All mercury compounds sublime when heated, and yield reaction 4 under mercurous salts. 7. All mercury salts heated in a test tube with Na2C03 give a mercury mirror. 139. Bismuth, "Si"' . — Confirm reactions, using solution of BiCl^. 1. HgS ppts. black BigSg, soluble in HNO3, insoluble in (NHJ2S. 2. H2O ppts. white basic BiOCl, when the solution with HCl, evaporated nearly to dryness, is poured into much water. The ppt. is insoluble in tartaric acid (compare Sb test, 142-4). 3. KOH or (NH4)0H ppts. white Bi(0H)3, insoluble in excess, becoming yellow BigOs on boiling. 4. K2Cr04 ppts. yellow Bi2(Cr04)3, soluble in HNO3, insoluble in NaOH. 140. Copper, Q,^x". — Confirm reactions, using solution of CuSO^. I. HjS ppts. black CuS, soluble in HNO3 and KCy, insoluble in (NH4)2S. THE METALS. 5l 2. NH4OH ppts. greenish blue basic salts, soluble in excess to a dark blue solution of a double salt of copper and ammonium. 3. KOH or NaOH ppts. pale blue Cu(0H)2, insoluble in excess. 4. K^FeCyg ppts. brown CugFeCyg, insoluble in dilute acids, decomposed by KOH. 5. A bright steel needle dipped in a slightly acid solution of a copper salt becomes covered with metallic copper. 6. Cu salts moistened with HCl and heated in a Bunsen flame tinge it green. Solutions of Cu are always blue or green. 7. Solids fused with NajCOg + KCy on charcoal, before the reducing blowpipe flame, yield bright red metallic beads, soluble in HNO3. 141. Arsenic, As'". — Confirm 7-eactions, using solution of As^O^. 1. HgS ppts., from acid solution, yellow AS2S3, soluble in (NH4)2S, KOH, NH4OH and HNO3, nearly insoluble in hot HCl (compare Sb, 142-1). When As is in the form of arsenic acid, HgAsO^, it must first be reduced to arsenous acid, H3ASO3, by heating with HCl and NajSOg. 2. Alkalies produce no ppt. 3. AgNOg ppts. from neutral solutions pale yellow Ag3As03. To a solution of AsgOg forming arsenous acid, H3ASO3, add a few drops of AgN03. Touch the surface of the solution with a glass rod moistened with dilute (NH4)0H. The yellow Ag3As03 forms. With arsenic acid, the ppt. is chocolate color (151). 4. Marsh's Test. Fit a generator with a bent ignition jet and thistle tube, like the one on the demonstration desk. In 62 LABORATORY MANUAL OF MEDICAL CHEMISTRY. it generate H from pure zinc, water and H2SO4. Observe the usual precaution and light the jet. The flame leaves no spot when cold porcelain is pressed in it, and no stain forms on the jet when the tube is heated to a dull redness by a Bunsen burner. Through the thistle tube add a few drops of an arsenic solution. Spots may now be deposited on porcelain, and a mirror on the tube when heated. Tests for Arsenic Spots. {a) They dissolve in a solution of bleaching powder. {b^ They turn yellow, touched with (NH4)2S and evapo- rated. The residue is insoluble in HCl, soluble in NH4OH. (f) They dissolve in HNO3, and when this is evaporated and touched with AgNOg and NH^OH and again evaporated, a brick-red spot remains (arsenic acid, 141-3). Tests for Arsenic Mirror. Cut out a section of the tube containing the mirror, so that about 3 inches of clean tube extends beyond the stain. Hold the tube at an angle. Gently heat the mirror and drive the As into the clean portion of the tube. It will be deposited as AS2O3, seen under the microscope as octahedral crystals. These may be dissolved in 10 drops of hot water and tested by 141-3, or acidulated and tested by 141-1. 5. ReinscKs Test. Bright copper boiled with an acid solution of arsenic will be coated with metallic arsenic. Fold this up and heat it in a glass tube open at both ends. The As sublimes as AS2O3, which may be detected by the microscope as octahedral crystals. These may be dissolved in H2O and tested by 141-1 and 3. This test is a delicate and useful one, as it is unim- paired in solutions containing organic matter. Arsenates act more slowly than arsenites. Antimony and mercury THE METALS. 63 both give a deposit on copper, but Sb sublimes as an amorphous powder and Hg as minute metallic globules. 6. Solids heated in a tube with carbon, or K^FeCyg, or a mixture of NagCOg and KCy, deposit metallic As on the sides of the tube. 7. Metallic As heated on charcoal burns, yielding copi- ous white fumes of AS2O3, having a garlic odor. 142. Antimony, Sb'". — Confirm reactions, using solution of Tartar Emetic. 1. HgS ppts. orange SbgSg, soluble in (NHi)3S, KOH, and hot concentrated HCl (compare As, 141-1), insoluble in NH^H. 2. KOH ppts. Sb203, soluble in excess. 3. NH4OH ppts. SbgOs, insoluble in excess. 4. Solutions concentrated with addition of a few drops of HCl and poured into water ppts. white SbOCl, soluble in tartaric acid (compare Bi, 139-2). 5. Marsh's Test. Performed like test 141-4, for arsenic. Tests for Antimony Spots. (a) They are insoluble in a solution of bleaching powder. (^) They turn orange when touched by (NH4^)2S, and evaporated. The orange spots are insoluble in (NH4)OH, but soluble in HCl. (f) They dissolve in HNO3, but yield no red color with AgNOg and NH4OH (compare, 141-4, c). Tests for the Antimony Mirror. It is often blacker than the As mirror, and yields no crystals under the mi- croscope when treated as directed under As. The white amorphous residue is insoluble, and the attempted solution does not yield tests with AgNOg or H2S, as under arsenic. 64 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 6. Solids fused with NagCOj on charcoal, before the reducing blowpipe flame, yield brittle metallic globules of Sb and white incrustations of Sb203. SOME IMPORTANT SYNTHETICAL REACTIONS OF GROUP II. 143. Preparation of Bismuth Nitrate and Sub-Nitrate. Dissolve a little powdered metallic Bi in a few drops of HNO3. Bismuth nitrate Bi(N03)3 is formed. Pour this into a beaker of water. Bismuth sub- nitrate is precipitated. Reaction— Bi(N03)3+H20=BiON03+2HN03. 144. Preparation of the Two Mercury Nitrates. (a) Mercurous Nitrate. Heat a globule of Hg with 2 drops of HNO3. Hg is in excess, and mercurous nitrate, HgN03, is formed. Dissolve this in a little distilled water, and save for 145. (i) Mercuric Nitrate. Heat the Hg as above with more HNO3 than is required to dissolve it. Mercuric nitrate, HglNOsJs, is formed. Dis- solve in distilled water, and save for 145. 145. Preparation of the Two Mercury Iodides. (a) Mercurous Iodide. To the HgNOs formed above, add KI. A yellowish green ppt. of yellow mercurous iodide, Hgl, falls. {b) Mercuric Iodide. To the Hg(N03)2 formed above, add KI. A bright red ppt. of red mercuric iodide, Hgl2, falls. 146. Preparation of the Two Mercury Oxides. (a) Mercurous Oxide. To calomel, HgCl, add a solution of KOH, or Ca(OH)2. Black mercurous oxide, Hg202, is formed. (5) Mercuric Oxide. To a solution of mercuric chloride, HgCla, add KOH, or Ca{OH)2. Yellow mercuric oxide, \i%0,iaSSs. Heat a few grains of Hg(N03)2 in a test tube until fumes cease. The red •inercuric oxide, HgO (red precipitated, remains. 147. Preparation of the Two Mercury Chlorides. (a) Mercurous Chloride. Heat in a test tube equal parts of finely powdered dry NaCl and mercurous sulphate, Hg2S04. Calomel, HgCl, sublimes in the cool part of the tube. Tested by NH4OH, it turns black. (i) Mercuric Chloride. Repeat the above experiment, using mer- THE METALS. 65 curie sulphate, HgSOi. Mercuric chloride, HgCl2, or corrosive subHmate, sublimes. It is tested by dissolving some in water and adding H2S. A black ppt. of HgS falls. 148. Test for Corrosive Sublimate in Calomel. By the above. processes the products are usually mixed. Calomel is insoluble, HgClj is soluble. Boil the suspected calomel, filter and pass iH2S through the filtrate. A black ppt. indicates HgCl2, by the formation of HgS. 149. Formation of Arsenous Acid. Note that AS2O3 dissolves slowly in cold water, faster in hot. and also when HCl or alkalies are added. Reaction — As203+3H20=2H3As03, or arsenous acid. 150. Preparation of Fowler's Solution, Boil a grain of AS2O3 with water containing a little K2CO3, and filter if necessary. This solution of potassium arsenite, KH2ASO3, colored with compound tincture of lavender and diluted to i % of As , is Fowler's Solution. Only alkali arsen- ites are soluble. 151. Preparation of Arsenic Acid. Dissolve a little AS2O3 in water, forming arsenous acid. Filter and divide into two parts. Boil one part nearly to dryness, with a little strong HNO3, which will oxidize arsenous acid, H3ASO3, to arsenic acid, H3ASO4. Arsenous acid, on addition of AgN03 and neutralizing with NH4OH, yields a yellow ppt., while arsenic acid yields a chocolate ppt., 141-3. Arsenous acid is also oxidized to arsenic acid by heating thus with KCIO3 and HCl. Arsenic acid is reduced to arsenous acid by the addition of Na2S03 and HCl, when SO2 is liberated. 152. Preparation of Tartar Emetic. Precipitate a solution of SbCU with Na2C03. Boil the well-washed ppt. Sb203 with an equal weight of cream of tartar in solution. Evaporate, and collect the crystals. Reaction-Sb203+KH(C4H40o)=2K(SbO)(C4H408)+H20. LABORATORY QUESTIONS ON GROUP II, 1. How many of the sulphides of the copper group are soluble in (NH4)2S? 2. Why does Pb appear in the second group? 3. Why does Hg appear in the second group ? 4. What difference is there in the addition of KOH and (NH4)OH to copper solutions ? 66 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 5. How does bleaching powder distinguish between As and Sb ? 6. Which contains the more acid, mercurdus or mercuric nitrate? 7. How might one detect piclcles turned green by copper? 8. Give the formulas for arsenous and arsenic acids. Which is the more common? 9. Is there any difference in Valence in Hg(ous) and Hg(ic) com- pounds ? 10. Write the reaction when CuO is heated with C. When AS2O3 is heated with C. 11. H3As03+NaOH= ? Complete and balance. 12. What is an amalgam? 13. Why is calomel apt to contain corrosive sublimate? 14. Why is Reinsch's test useful ? 15. What are the substances called red and white precipitates? 16. Write the reaction involved in the formation of mercuric chloride, as prepared above. 17. What substance would one get from a druggist if arsenic were called for? 18. What substance would one get from a druggist if antimony were called for? 19. Is there a difference in solubility among the varieties of white arsenic ? 20. When a dime is dissolved in HNO3, what evidence is there that it is an alloy of Cu ? 21. Make a solution of CUSO4, Pb(N03)2, Bi(N03)3, HgCU, AS2O3 and K(SbO)(C4H40e), acidulate with HCl, precipitate the sulphides, and separate by Table III, page 77. 22. Make a solution of the first two groups, containing AgNOs, PbNOs and CUSO4. Turn to the table, and separate and identify each metal. 23. Examine bottles 7, 8, 9, 10, 11 and 12, on the side table, for metals of the first two groups, and report the metals in each. 24. Dissolve lin HNO3 the metals E, F, G and H, filter, nearly neu- tralize, separate by the table, and report the metals of the first two groups found. ANALYTICAL REACTIONS OF GROUP III. 153. Iron(ous) Ffc''. — Confirm reactions, using solution of FeSO^. I. NH4OH ppts. Fe(0H)2, white when pure, usually dirty THE METALS. 67 green from the presence of ferric salts, gradually turning, by absorption of O, to reddish brown ferric hydroxide, Fe(0H)3. 2. (NH4)2S ppts. black FeS, soluble in HCl, insoluble in alkalies. 3. (NH4)2C03 ppts. white FeCOg, soluble in the pres- ence of free CO2. 4. Tannic acid and tannates ppt. black iron tannate. 5. Ferrous salts are oxidized to ferric salts, slowly by standing, rapidly by heating with oxidizing agents, HNO3, KCIO3, etc. 154. Iron(ic), Fe'". — Confirm reactions, using solution of FeCl^. 1. NH4OH or KOH ppts. reddish brown Fe(0H)3, in- soluble in excess. 2. (NH4)2S ppts. black FeS mixed with S, soluble in HCl and HNO3. 3. H2S, in acid solutions, ppts. S and reduces "ic" to " ous" iron. 4. KCyS, KgFeCye and K^FeCyg have the following reac- tion on "ous" and "ic" iron salts: Reairent. Ferrous Salts Ferric Salts. KCyS. KsFeCye. KiFeCys. No change. Dark blue ppt. Fe3{FeCy6)2. Pale blue ppt. (K2Fe)(FeCy6). Red solution Fe(CyS)3. No ppt. Reddish brown solution. Dark blue ppt. Fe4(FeCy6)3. 155. Aluminum, kV". — Confirm -reactions, using solution of Al^(_SOi)s- 1. KOH and (NHJOH ppt. white A1(0H)3, soluble in excess, but re-pptd. by the addition of ammonium salts. 2. (NH4)2S ppts. white Al(OH)3. 3. NagHPO^ ppts. white Al2(P04)2. 4. Solids fused on charcoal by the blowpipe, moistened with C0CI2 and reheated, yield an infusible 6lue mass. 68 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 156. Zinc, Zn". — Confirm reactions, using solution of ZnSO^. 1. (NH4)2S ppts. white ZnS, soluble in acids. 2. KOH ppts. white Zn(0H)2, soluble in excess, repre- cipitated on boiling. 3. (NH4)OH ppts. white Zn(OH)2, soluble in slight excess. 4. Na2C03 ppts. white basic carbonate. 5. Solids fused with NajCOg on charcoal, in the reducing blowpipe flame, become incrusted with ZnO, yellow while hot, white when cold. 6. Solids fused on charcoal with the blowpipe, and moistened with C0CI2 and reheated, yield an infusible green mass. SOME IMPORTANT SYNTHETICAL REACTIONS OF GROUP III. 157. Preparation of the Two Iron Sulphates. (a) Ferrous Sulphate. Dissolve iron filings or small nails in dilute H2SO4 until all action ceases. Filter, evaporate to small bulk and crystallize the green ferrous sulphate, FeS04, or iron proto sulphate. (b) Ferric Sulphate. Dissolve 6 parts by weight of FeS04 in water, and add i part of H2SO4. Heat the mixture and drop in HNO3 until the black color first formed disappears. The resulting solution contains ferric sulphate Fe2(S04)3 or ferric tersulphate. When less acid is used the solution contains basic ferric sulphate, Fe40(S04)5. Reaction— 6FeS04+3H2S04+2HN03=3Fe2i.S04)3+2NO+4H20. 158. Preparation of the Two Iron Chlorides. (a) Ferrous Chloride. Dissolve iron filings in HCl until action ceases. Filter, evaporate and crystallize the ferrous chloride, FeCb. (b) Ferric Chloride. Acidulate a solution of FeCb with HCI, heat and add HNO3 until the black color first formed disappears. The resulting fluid is a solution of ferric chloride, FeCls. Reaction— 6FeCl2+2HN03+6HCl=6FeCl3+2NO+4H20. 159. Preparation of Ferrous Iodide. Heat a few crystals of iodine with iron filings and water in a test tube. When the iodine has disappeared, THE METALS. 69 filter the greenish solution. On evaporation the solid ferrous iodide, FeCU, remains. 160. Preparation of Ferric Acetate. Precipitate a solution of FeCls with NH4OH and wash the precipitated Fe(OH)3. Dissolve this ferric hydroxide in glacial acetic acid. The solution contains ferric acetate, Fe(C2H302)3. 161. Preparation of the Citrate of Iron and Ammonium. Prepare ferric hydroxide as above, and quickly and thoroughly wash. Dissolve it in as little of a solution of citric acid as possible, thus forming ferric citrate. Next add one-third its volume of NH4OH. No ppt. falls, due to the formation of a double citrate of iron and ammonium. Evaporated over a water bath to a small bulk and spread on glass to dry, it forms scales, and is one of the so-called "scale" compounds of iron. 162. Preparation of Reduced Iron. Prepare a glass tube 18 inches long. In one end loosely push some CaCl2, in the other partly fill 4 inches of the tube with freshly prepared and dried Fe(OH)3. From a H generator pass the gas first over the CaCl2 to dry it, and then over the Fe(OH)3. Heat the Fe(OH)3 by placing a Bunsen burner under the tube, and keep up the evolution of H and heat until no more moisture comes from the iron. The iron is left in the tube in a finely divided state, and when hot burns spontaneously in the air, forming Fe203. Reaction— Fe(0H)3 + 3H =Fe + 3H2O. 163. Preparation of Dialyzed Iron. Dissolve in a warm solution of FeCl3 freshly prepared Fe(OH)3. Filter the solution, and place in a dialyzer and dialyze through parchment paper until the fresh difiusate gives no ppt. with AgNOs. Only colloid iron compounds then remain in the dialyzer, composed of about 5 per cent of basic ferric-oxy- chloride, Fe20Cl4. LABORATORY QUESTIONS ON GROUP III. 1. What is the diflference in Valence in ' ' ous ' ' and " ic " iron ? 2. What is FezCle? 3. What is blue vitriol? White vitriol? Green vitriol? 4. What action has the aircon ferrous compounds? 5. How much iron in 10 grains of Fel2 ? ^0 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 6. What members of Group III are pptd. by the group reagent as sulphides? Which as hydroxides? 7. How is carbonate of iron formed ? 8. Why does KsFeCyo afford the best distinguishing test for ' ' ous " and "ic" iron? 9. Why do solutions of "ous" iron usually contain "ic" iron? 10. What compound of iron is found dissolved in chalybeate waters ? 11. What is the composition of black ink? 12. What is galvanized iron? Prepare by dipping bright iron wire in melted zinc. 13. How would one prepare ZnCls for "soldering fluid" ? 14. Why would sewer gas turn white paint black ? 15. In what experiment did you prepare ferrous sulphide? 16. If KOH be added to excess to a solution of Fe(ic), Al and Zn, what would occur? 17. If ammonium salts were added to an alkaline solution of Al and Z, what would occur? 18. What are the blowpipe tests for Zn and Al ? 19. Could dialyzed iron in the stomach enter the circulation? 20. What are the "scale" compounds of iron? 21. Write the reactions which occur when (NH4)2S is added to solu- tions of third group metals. 22. Make a solution of FeCla, Al2(S04)3, and ZnSOi ; turn to Ana- lytical Table III, page 77, and separate by methods of the third group. 23. Make a solution of HgNOa, K{SbO)C4H406, Al2(S04)3, and ZnS04, separate and identify each by Table III, page 77. 24. Test bottles 13, 14, 15, 16, 17 and 18 on the side table, and re- port the metals found in each. 25. Report the minerals of the first three groups found in ores I, J, K and L. ANALYTICAL REACTIONS OF GROUP IV. 164. Barium, Ba". — Confirm reactions, using solution of BaCl^- 1. (NH4)2C03 ppts. white BaCOg, soluble in acids. 2. KgCrO^ ppts. yellow BaCrO^, insoluble in acetic acid, THE METALS. 7 1 soluble in HCl or HNO3. Ca yields no ppt. with this reagent. 3. H2SO4 ppts. white BaS04, insoluble in hot water, acids or alkalies. 4. Ammonium oxalate, (NH4)2C204, ppts. white BaCgO^, soluble in HCi and HNO3. 5. Solids moistened with HCl and heated in the Bunsen flame impart to it a yellowish green tinge. 165. Calcium, Ca". — Confirm reactions, using solution of CaCl^. 1. (N 114)2003 ppts. white CaCOs, becoming crystalline on heating. 2. H2SO4 ppts. white CaSO^ from strong calcium solu- tions, soluble in water and acids. 3. (NH4)2C204, ammonium oxalate, ppts. white CaC204, even in very dilute solutions, soluble in HCl and HNO3, insoluble in acetic acid. 4. Solids moistened with HCl and heated in the Bun- sen flame tinge it dull red, not seen in the presence of Ba. SOME IMPORTANT SYNTHETICAL REACTIONS OF GROUP IV, i6e. Preparation of Bleaching Powder. Place some damp slaked lime in a bottle and pass into the bottle CI gas. The gas is ab- sorbed, forming probably a loosely united compound of calcium hypo- chlorite and calcium chloride, thus: 2Ca(OH)2 + 4CI ^ Ca(C10)2' CaCl2 +H2O. When treated with acids, the Ca(C10)2 breaks up thus: 2HCl + Ca(C10)2^CaCl2 +2HCIO, and in the presence of more acid ; 2HC10 + 2HC1 = 4H20-|-2C1. When hypochlorites are treated with acids they liberate CI, and chlorides remain in solution. Their so- lutions rendered acid bleach litmus solution. 167. Preparation of Quicklime. Heat a small piece of CaCOs or marble in the blowpipe flame, keeping it white hot for some time. Cool it, and notice that a portion of it does not effervesce with HCl. 72 LABORATORY MANUAL OF MEDICAL CHEMISTRY. It dissolves in H2O, and gives the water an alkaline reaction. It is CaO, or quickime. Reaction— CaCOa + heat = CaO + CO2. 168. Preparation of Plaster Paris. Heat CaS04-2H20 (gypsum) in a test tube until the moisture is all given off. Considerable heat is required, but overheating must be avoided. Anhydrous CaS04, or exsiccated calcium sulphate, or plaster paris, remains, which when wet with a little water quickly hardens. LABORATORY QUESTIONS ON GROUP IV. 1. Write the reactions occurring when lime is slaked. 2. How may calcium chloride be formed ? Write the reaction. 3. Name two important uses you have thus far made of Ba com- pounds. 4. What is gypsum? Is it soluble in water? 5. Will moist bleaching powder give off CI ? 6. What is the composition of slaked lime ? 7. What is anhydrous gypsum called? 8. Make a solution of gum arable (calcium gummate), and test for calcium by test 165-3. 9. What is whiting? Whitewash? Why does mortar harden when exposed to the air? 10. Will lime harden under water or protected from the air? Why? 11. How were the marbles and limestones formed? 12. How are stalactites and stalagmites formed? 13. What test best distinguishes Ba from Ca? 14. Make a solution of BaCl2 and CaCla, and turn to Table III, page 77, and separate according to the scheme for Group IV. 15. Make a solution of CUSO4, AS2O3, FeCls, BaCU, and CaCl2, and separate and identify each metal by the table. 16. Test bottles 19, 20, 21 and 22 on the side table for metals of the first four groups, reporting the metals found in each. ANALYTICAL REACTIONS OF GROUP V. 169. Magnesium, Mg". — Confirm reactions, using solution of MgCh. I. (NH4)0H and (NH4)2C03 give no ppts. in presence of ammonium salts. THE METALS. 73 2. (NH4)2HAs04 ppts. white NH4MgAs04 in presence of (NH4)C1 and (NH4)0H. 3. NaaHPOi ppts. white crystalline Mg(NH4)P04 in presence of NH4CI and (NH4)OH, hastened by stirring. 4. Solids fused on charcoal, moistened with C0CI2 and reheated yield a pink mass. Metallic Mg burns with a brilliant white light, leaving MgO. 170. Potassium, K'. — Confirm reactions, using a solution of KCl. 1. PtCl4 ppts. yellow crystalline double chloride PtCl4- 2KCI from moderately concentrated solutions, hastened by presence of alcohol. Test thus : To i drop of a solution of KCl, "bn a micro- scope slide, add i drop of PtCl4 and i drop of alcohol ; stir, set aside, and examine with the microscope for yellow octa- hedral crystals. 2. H2(C4H40e), tartaric acid, ppts. white crystalline KH- (C4H4O6), promoted by stirring. 3. Solids heated on platinum wire in the non-luminous Bunsen flame tinge it violet, seen through blue glass. 171. Sodium, Na'. — Confirm reactions, using a solution oj NaCl. 1. Neither PtCl4 or H2(C4H406) yield ppts. 2. Solids heated on platinum wire in the non-luminous Bunsen flame tinge it intensely yellow, not seen through thick blue glass. 172. Ammonium, NH4'. — Confirm reactions, using solution of NH^Cl. 1. PtCl4 and H2(C4H40e) ppt. as with K. 2. Heated with KOH or NaOH, fumes of NH3 are evolved. 74 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 3. Nessler's reagenti ppts. brown NHggI, or in dilute solutions a yellow coloration. 4. Solids heated on platinum foil volatilize completely. K or Na salts do not. SOME IMPORTANT SYNTHETICAL REACTIONS OF GROUP V. 173. Preparation of Potassium Hydroxide. Boil in a test tube with water equal parts of K2CO3 and Ca{OH)2, set aside, and when subsided decant the KOH solution. 174. Preparation of Potassium Iodide. Into 10 c.c. of a warm solution of KOH stir crystals of I until the solution is permanendy faint yellow. Evaporate to dryness. Reaction— 6KOH+3l2=5KI+KI03+3H20. Heat the resulting product containing potassium iodide and iodate with an equal part of finely powdered charcoal until a slight incandescence ensues. 2KI03+3C2=2KI+6CO. Treat the mass when cool with water, warm, filter, evaporate and crystallize out the KI. 175. Preparation of Potassium Permanganate. Mix equal parts of solid KOH, KCIO3 and Mn02. Fuse the mixture on platinum foil until it is dark green. Reaction— 6KOH+3Mn02+KC10=3K2Mn04+KCl+3H20. Boil this mass in water, and a purple solution of potassium perman- ganate results. 3K2Mn04+2H20=K2Mn208+Mn02+4KOH. 176. Preparation of Cream of Tartar. To 10 c.c. of a strong cold solu- tion of KNO3 add an excess of a strong solution of tartaric acid. A white ppt. falls of acid potassium tartrate, KH{C4H40c), or cream of tartar. This method merely illustrates a possible method of preparation. It is not the method of commerce. 177. Preparation of Sodium Bi-Carbonate. Prepare a concentrated solution of Na2C03 in less than its own weight of water. Cool this, and slowly let CO2 bubble through the solution for some time. Set aside and Note i. To a solution of HgCl2 add KI until tlie ppt. is nearly all redissolved. Render strongly alkalipe with KOH. Set aside. Decant the clear straw-yellow solu- tion for use. THE METALS. 75 decant, drying the precipitated NaHCOs on filter paper. Note that an equal quantity of this acid sodium carbonate is alkaline in reaction, and effervesces more violently with acids than the original NazCOs. Reaction— Na2C03+C02+H20=2NaHC03. 178. Preparation of Rochelle Salts. To a strong hot solution of Na2C03 add cream of tartar (KH){C4H406) until no further effervescence occurs. Set the solution aside to cool, and the double salt, potassium and sodium tartrate, KNa(C4H40o) will crystallize out. Reaction— Na2C03+2KH(C4H406)=2KNa(C4H406)+C02+H20. 179. Preparation of Sodium Hypochlorite. Malce a clear solution of bleaching powder. Dissolve in a little hot water twice as much Na2C03 as the weight of bleaching powder used. Mix the two solutions and filter. Add water to the filtrate until the hydrometer, immersed in it, stands at 1.059. The solution (Labarraque's Solution) contains sodium hypochlorite, NaClO, a useful disinfectant, liberating CI. Reaction— Ca(C10)2+2Na2C03=2NaC10+CaC03. 180. Preparation of Ammonium Sulphide. Pass H2S through a solu- tion of NH4OH until the solution no longer gives a ppt. when a few drops are added to a little magnesium sulphate solution. Ammonium hydrosulphide is formed, NH4HS. (Reaction NH4(OH)+H2S= NH4HS+H2O). To this solution add nearly an equal volume of NH4OH. Ammonium sulphide, much used in analysis, is formed thus : NH4HS+(NH4)OH=(NH4)2S+H20. Yellow ammonium sulphide is usually more efficient, and is prepared by adding to ammonium sulphide a pinch of flowers of sulphur, or letting the reagent stand for some time until yellow. (NH4)2S2 is thus formed. LABORATORY QUESTIONS ON GROUP V. 1. What peculiar action has ammonium salts on solutions of Mg? 2. How may both K and NH4 be detected in a mixture ? 3. What is the original source of K2CO3, used in 173 ? 4. What is cream of tartar ? Write its reaction on bicarbonate of soda. 5. How might you manufacture baking powder? 6. Why does soda sweeten sour batter cakes? 7. Why is NaHCOs called bicarbonate of soda ? 76 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 8. What are Rochelle salts? Epsom salts? Glauber's salts? 9. In what operations has yellow {NH4)2S or (NH4)3S2 been used? 10. How might ammonium acetate be prepared? 11. For what purpose is NaHCOs more useful than Na2C03? 12. What is pearlash ? Sal soda? Saleratus? Soda? 13. What is calcined magnesia (169-4) ? Is it soluble in water? Is the attempted solution alkaline? Write the reaction. 14. Make a solution of K, Na and NH4 salts, and separate by Table III, page 77, under Group V. a ^ ■ p wi C II , ^ "^ n !> = 1 OE. • ntnan P.S-": rrlB. a>ao.5- q, ";o3?g in £.n B M ^ — »< T -f^ 3 O E= i^nSi '^'■aBi'ia^Pf ■2-691 -3^1 = siiqyw '9 ffis a.2 focnSW If If? 5-3 3 W- it O ^ ° « HI a Q & 3 " p g; Do a CD t^ g w si H S, o ■I ff a Pi I o EXERCISES IN GRAVIMETRIC AND VOLUMETRIC PROCESSES APPLIED TO SANITARY WATER ANALYSIS EXERCISES IN GRAVIMETRIC PROCESSES. EXERCISES IN GRAVIMETRIC PROCESSES. i8i. Determination of the Per Cent of Soluble Solids. Weigh a clean, dry evaporating dish on a good balance. i Fill the dish three-fourths full of the liquid under examination, as, for example, drinking water. Carefully reweigh. Evapo- rate the solution to dryness over a slow fire, or better, a water bath, avoiding loss by "spitting." Cool the residue by placing the dish under a desiccator. 2 Remove the resi- due, when cold, to the balance and quickly weigh. From these data compute the per cent of solids. 182. Determination of the Approximate Weight of a Precipitate. Determine the amount of lead in a solution thus : Precipi- tate the solution by passing H2S through it for some time. Weigh a well-dried filter paper. Filter the solution through this paper, collecting the well-washed precipitate on it. Dry the funnel and its contents in a drying oven. Cool the dry paper and its dry precipitate in a desiccator. Weigh, and subtract the weight of the paper. The remainder is the weight of lead sulphide, PbS. From this compute the actual amount of lead in the solution. 183. Determination of the Exact Weight of a Precipitate. To determine the per cent of SO4 in sodium sulphate, weigh accurately a small porcelain crucible. Place in it a large pinch of dry, powdered sodium sulphate. Reweigh the Note i. A good balance must be carefully handled. Special instruction in the methods of weighing, handling weights, determining centers, care of the balances, etc., will be given by the demonstrator. Note 2. This is a bell-jar containing a dish of strong H2SO4, which absorbs all the moisture, and keeps the material from gaining in weight. 82 LABORATORY MANUAL OF MEDICAL CHEMISTRY. crucible and contents. Carefully wash the salt into a beaker with about 25 c.c. of distilled water. Add a few drops of HCl, and heat to boiling. Next add BaClj until no further ppt. is formed, boil and filter through a fine Swedish filter paper, which has a known ash. Wash the ppt. with distilled water until the washings no longer ren- der AgNOg solution turbid, showing absence of chlorides. Dry the filter in an air bath. Fold up the dry paper com- pactly. Wrap a platinum wire about it. Hold it over the crucible. Burn it and shake the ash into the crucible. Heat the crucible to redness for a few minutes, until all carbon is burned and the residue is white. Cool in a desic- cator. Weigh when cold. Subtract from this weight the weight of the crucible and the known filter ash. The re- mainder is the weight of barium sulphate. Knowing the weight of NagSOi taken and BaSOi found, (i) From the weight calculate what % of the BaS04 is SO4 (2) From the weight calculate what % of NagSOi was SO4 _^^==^ Error ^„^^„ EXERCISES IN VOLUMETRIC PROCESSES. Volumetric solutions are solutions of a definite weight of a given substance dissolved in a definite amount of dis- tilled water. They are designated by the U. S. P. "as normal (5) when they contain in one litre the molecular EXERCISES IN VOLUMETRIC PROCESSES. 83 weight of the active reagent expressed in grams and reduced to the valency corresponding to one atom of replaceable hydrogen." A decinormal solution (^"^) is -jV the strength of a normal one, and is preferable in many operations. ACIDIMETRY. 184, The Process of Determining the Amount of Free Acid in a Solution. The Solution. A decinormal volumetric solution of KOH is prepared by dissolving 5.599 grams of solid KOH in distilled water, and making up the solution to 1,000 c.c. at 15° C. In accurate work this must always be standardized. The Process. The process is known as titration. A burette is filled with decinormal KOH solution up to the 0° mark. Place in a beaker a known quantity of the clear solution in which the free acid is to be estimated. (Use 30 c.c. of acidum hydrochloricum dilutum, and calculate the per cent of acid.) Add to the beaker 5 drops of a solution of phenolphthalein as an indicator, which turns red with the slightest excess of alkali. Place the beaker on white paper, and with constant stirring add the alkali from the burette, drop by drop, until a faint pink tinge becomes permanent. Read off the number of c.c.'s of alkali required to neutralize the acid. From the following table the exact weight of the acid may be computed. Each c.c. of alkali neutralizes exactty — Acetic Acid, absolute H(C2H302) 0.005986 g. Citric Acid, H3(C6H507)H20 0.006983 Hydrobromic Acid, absolute HBr. 0.008076 Hydrochloric Acid, absolute HCl 0.003637 84 LABORATORY MANUAL OF MEDICAL CHEMISTRY. HydriodicAcid, absolute HI o.oi2753g. Hypophosphorous Acid, H(H2P02) 0.006588 Lactic Acid, absolute H(C3Hg03) .... 0.008979 Nitric Acid, absolute HNO3 0.006289 Oxalic Acid, crystallized H2C204-H20 0.006285 Phosphoric Acid, absolute H3PO4 0.004890 Sulphuric Acid, absolute H2SO4 0.004891 Tartaric Acid, crystallized H2(C4H40g) 0.007482 ALKALIMETRY. 185. The Process of Determining the Amount of Free Alkali in a Solution. The Solution. A decinormal volumetric solution of oxalic acid, H2C204-2H20, is prepared by dissolving 6.285 grams of pure crystals in distilled water, and making up to 1,000 c.c. at 15° C. In accurate work this solution must always be standardized. The Process. Fill a burette with decinormal oxalic acid solution up to the 0° mark. Place in a beaker a known amount of any clear solution in which the amount of free alkali is to be determined. (Use 100 c.c. of liquor calcis, and prove that it contains 14% of Ca(OH2), as required.) Color the solution yellow by a few drops of methyl-orange solution for an indicator, which turns red with the slightest excess of acid. With constant stirring, drop in the oxalic acid from the burette until the pink color appears. Read off the number of c.c.'s of acid used, and compute the absolute weight of alkali. Each c.c. of oxalic acid neutral- izes exactly — Ammonia Gas, NH3 0.001701 g. Ammonium Carbonate, (NH4')2C03 0.004793 EXERCISES IN VOLUMETRIC PROCESSES. 85 Calcium Hydroxide, Ca(0H)2 0.003691 g. Potassium Hydroxide, KOH 0.005599 Sodium Bicarbonate, NaHCOg 0.008385 Sodium Carbonate, NagCOg 0.005292 Sodium Hydroxide, NaOH o. 003996 186. Estimation of Chlorine. The Solution. A standard volumetric solution of silver nitrate convenient in water analysis is prepared by dissolv- ing 4.79 grams of pure AgNOg in distilled water and making it up to 1,000 c.c. at I5°C. This solution is of such strength that i c.c. of it precipitates exactly one milli- gram of chlorine (or i c.c. is equivalent to .00165 g- of NaCl). The solution must be kept in a dark container. The Process. Fill a clean burette to the 0° mark with the standard solution of silver nitrate. Place a beaker on a piece of white paper. In it put a known quantity of the clear solution in which chlorides are to be determined. (Take 70 c.c. of drinking water. i) Add a crystal of K2Cr04 to color the water lemon-yellow. With constant stirring, drop in the silver nitrate from the burette until a faint red color is per- manent. Read the number of c.c. 's used. Refill the burette and repeat the process with an equal volume of distilled water exactly matching the red color. Read the number of c.c.'s used, and deduct this from the first reading, as the amount required to produce the color. The number of c.c.'s remaining represents grains of chlorine per gallon, or may be calculated as sodium chloride. Note i. Seventy c.c. of distilled water weighs 70,000 milligrams. The imperial gallon contains 70,000 grams, hence the number of milligrams of solids in 70 c.c. will likewise express the number of grains in a gallon. 86 LABORATORY MANUAL OF MEDICAL CHEMISTRY. SANITARY ANALYSIS OF POTABLE WATERS. 187. Detection of Poisonous Minerals. Poisonous metals most commonly present in potable waters are lead and copper, sometimes arsenic, zinc, etc. A quantity of water is evaporated to a small bulk. The metals may be then identi- fied by the general methods of qualitative analysis given in Table III, page 77. 188. Determination of Total Solids. Find the total solids in 70 c.c. of water, according to the method in 181. The num- ber of milligrams of solids likewise expresses grains per imperial gallon. River waters usually contain less solids than well waters. Solids ought not to run over 30 or 40 grains per gallon (Wanklyn). Artesian water often runs higher. The presence of carbonates in the residue is evinced by effervescence when touched with acid. 189. Determination of Chlorine. Find the number of grains of chlorine per gallon, as in 186. Pure waters, unless charged from some uncommon mineral deposits, contain little chlorine. Urine and sewage contaminated from animal excretions are highly charged with chlorides. A water free from chlorides could not have been contaminated. Water containing much chlorine would be looked upon with sus- picion until its albuminoid ammonia had been determined. River water contains not far from \% of chlorides. Well water contains more. Five or six grains of chlorine per gallon does not injure water, but is a reason for suspicion. SANITARY ANALYSIS OF POTABLE WATERS. 87 190. Determination of Free and Albuminoid Ammonia. I. Distillation. Thoroughly cleanse a Liebig condenser and a glass retort holding i litre. Arrange them for distil- lation. In the retort place 500 c.c. of the water under examination. Add a handful of broken, well-washed glass, to keep from "bumping." Slowly distill, catching each 50 c.c. of distillate in a large tube, and set each aside in the order of distillation until four tubes are filled, representing 200 c.c. These contain any free ammonia in the original portion, and are to be examined later by («). Stop the distillation. Add 50 c.c. of permanganate solu- tioni through a funnel, and begin to distill very cautiously to avoid bumping. Catch each 50 c.c. of distillate in a large tube, and set each aside as before until three are obtained. These contain ammonia from any organic matter present in the original solution, and are to be examined later for albu- minoid ammonia by {F). II. Nesslerizing. Prepare two solutions of ammonium chloride. 1. 3.15 g. of NH4CI dissolved in i litre of distilled water. 1 c.c. contains i milligram of ammonia. 2. Dilute I c.c. of the above with gg c.c. of water, i c.c. contains y^ of i milligram of ammonia. (a) Estimation of Free Ammonia. Take the tubes con- taining the free ammonia distillate. Into the first drop 2 c.c. of Nessler's reagent (172-3, Note i), and stir. It will strike a yellow or brownish color if traces of ammonia are present, assuming its deepest permanent hue in 3-5 minutes. The next process consists in making an artificial solution, Note i. The permanganate solution is prepared by dissolving 200 grams of solid potassium hydroxide and 8 grams of crystallized potassium permanganate in water, boiling for 20 minutes and making up to 1,000 c.c. with distilled water. This solution oxidizes organic matter and, in the presence of an alkali, liberates N in the form of NHs. 88 LABORATORY MANUAL OF MEDICAL CHEMISTRY. with a known amount of ammonia, that strikes the same color with Nessler's reagent. Take a tube containing about 45 c.c. of water free from ammonia. Add for the first trial, say 3 c.c. of the dilute ammonium chloride solution. Stir. Add 2 c.c. of Nessler's reagent. Stir, and compare the two colors by looking down through the solutions on a white surface. If the colors do not exactly agree in depth of shade, make another solution with more or less ammonium chloride, which will exactly match the distillate. The amount of ammonia added, obviously, is the amount con- tained in the first 50 c.c. of distillate. The remaining tubes might be Nesslerized, but as the first tube invariably contains three-fourths of the entire free ammonia, it is easier to add one-third for the total amount of free ammonia contained in the 500 c.c. of water examined. (iJ) Estimation of Albuminoid Ammonia. Nesslerize, in the manner described, each tube containing albuminoid ammonia, and add the results for total albuminoid ammonia in the 500 c.c. of water taken. (^) Discussion of Results. EXAMPLE OF ANALYSIS. IN 500 c.c. Free ammonia 01 milligram. Correction, % °°3 " Total 013 " Albuminoid ammonia 035 milligram. 015 " " " 000 " Total 050 " IN ONE LITRE. Free ammonia 026 milligram. Albuminoid ammonia 10 " SANITARY ANALYSIS OF POTABLE WATERS. 89 If water contains o.oo parts of albuminoid ammonia per million (that is, milligrams per litre), it is organically pure. If it contains 0.02-0.05 it is classed with good waters, o. 10 begins to be a suspicious sign, and 0.15 ought to condemn the water (Wanklyn). In river waters, where the organic matter is from vegetable origin, even a somewhat higher amount of albuminoid ammonia might exist, however, without serious effects. 191. Determination of Other Ingredients. Many other deter- minations are sometimes made in sanitary analysis of water, which may serve for valuable deductions concerning water supply, etc., such as the determination of nitrites, nitrates, hardness, bases, acids, oxygen-consuming power, etc. The above, however, are all that are usually required in determin- ing the purity of water. The residue from concentration may be used in determining the bases and acids present by the ordinary processes of analysis. See Tables I, II and III. For further work, the student is directed to Wanklyn's Water Analysis. LABORATORY EXERCISES IN THE MEDICINAL ORGANIC COMPOUNDS. ORGANIC CHEMISTRY. 93 ORGANIC CHEMISTRY. 192. Determination of the Fusing Point. Select a fine capil- lary tube, of an internal diameter just sufficient to insert a fine wire. Seal one end. Push a very small amount of the substance under examination into the lower end of the tube. Fasten the tube to a good chemical ther- mometer by a rubber band, and insert the whole in a beaker of water. Apply a gentle heat to the beaker, and as the temperature rises carefully note the reading of the ther- mometer the moment the solid melts. Whenever the fusing point is above 100° C, a beaker of melted paraffin may be used. Any impurity will account for variations in the fusing points. Determine the fusing points of the following sub- stances : Substance. Ohserved. Given. Chloral hydrate . . . ... Naphthalene . Stearic acid . 57° £■ 79° C. 69° C. 193. Determination of the Boiling Point. Place about 10 c.c.'s of the liquid under examination in a dry, clean 7-inch test tube, fitted with a 2-hole rubber stopper, holding a ther- mometer, which should nearly touch the surface of the liquid, and be free from the sides of the tube. Place a few bits of broken glass in the liquid to assist in the boiling. Place the whole in a beaker of water or melted paraffin, as required. Apply a gentle heat, and carefully note the reading of the 94 LABORATORY MANUAL OF MEDICAL CHEMISTRY. thermometer when the liquid is boiling vigorously and the thermometer stem is thoroughly surrounded by vapor. Determine the boiling point of the following substances. When the determination is made, return the liquid to its con- tainer. Stbstancs. Observed. Given. Ethyl Alcohol 78° C. 66° C. 35° C. Methyl Alcohol Ether 194. Optical Activity of Organic Chemicals. Examine with the polariscope the prepared solutions of cane sugar, which is dextro-rotary, and morphine hydrosulphate, which is laevo- rotary. 195. Detection of Carbon and Hydrogen in Organic Compounds. Dry and powder the substance under examination, as, for example, tartaric acid. Mix it with twice its bulk of dry finely pulverized cupric oxide, CuO. Place the mixture in a test tube, fitted with a delivery tube, constructed so that any gas formed may pass through a glass tube, in which is loosely placed a few pieces of dry CaClg, and then bubble through clear lime water. Heat the test tube to a dull red. The moistening of the CaCl2 indicates HgO, and a precipitate in the lime water, CO2, both formed from the union of the H and C of the tartaric acid with the O of the CuO. Note that metallic copper remains in the test tube. 196. Detection of Nitrogen in Organic Compounds. Place in a dry test tube a small piece of metallic sodium. Cover this with the dry powdered solid under examination, as, for example, urea or any alkaloid, and heat. Cyanides are ORGANIC CHEMISTRY. 95 formed if the compound contains nitrogen. Test the resi- due for cyanides by adding water, filtering and boiling with a drop each of FeSOi and NaOH, then adding a drop of FeClj and acidulating with HCl. Cyanides yield a ppt. of Prussian blue, which is best seen on the filter paper when the solution is filtered (129). 197. Detection of Sulphur in Organic Compounds. (i) Take some lead acetate solution in a test tube. Add KOH until the precipitate first formed dissolves. In this drop a little finely powdered organic substance containing sulphur, like albumin. Boil, and notice the blackening of the sulphur compound. A still more delicate test is the following : (2) In a dry test tube place a small piece of metallic sodium. Cover this with the dry powdered substance under examination, like albumin, and heat. If the compound con- tains sulphur, sodium sulphide, NagS, will be formed. After prolonged ignition, dissolve this residue in water. Filter, and add a few drops of sodium nitro-prusside, which yields with sulphides a purple coloration. 198. The Determination of Formula. Examine the combus- tion furnace with the combustion tube, containing an intimate mixture of CuO and .46g. of acetic acid. Note the arrange- ment of the weighed CaCl2 tubes to retain the H2O. Evi- dently -g- their increase in weight will be the weight of H in the compound. Note the weighed potash bulbs arranged to absorb all COg. Evidently ^ of their gain in weight will represent the C in the acetic acid. The remaining weight of the substance we know to be O, and its amount can be determined by the method of difference. g6 LABORATORY MANUAL OF MEDICAL CHEMISTRY. Result of the Combustion of .46 g. of Acetic Acid. C02=.6765g.XT^.i845g. of C, or 40.11% H20=.28i7g.Xi=.03i3g. of H, or 6.82% O by difference 53.07% 100.00 This does not give us the' formula. The weight of one molecule is determined by the Victor Meyer vapor density ap- paratus. Examine this instrument, with its outer casing in which aniline has been boiled, and the inner tube surrounded by hot aniline vapor. Note the little bottle in the inner tube, which was dropped in when the apparatus was hot. It con- tained .ogg. of acetic acid, which vaporized and crowded out 33.48 c.c. of air into the graduated tube. 33.48 c.c. of acetic acid vapour, then, weighs .og gram. 33.48 c.c. of H weighs .002gg gram (i c.c. = .oooo8g6 g.). Hence the acetic acid va- pour is about thirty times as heavy as an equal volume of H (ro^oiss ^ 3°)' Avogadro's law states that equal volumes of gases under equal conditions contain equal numbers of mole- cules. Then a molecule of acetic acid weighs thirty times as much as a molecule of hydrogen (H2), and 60 times an atom of hydrogen (H). The molecular weight of acetic acid, then, is 60, and the formula is thus calculated : Computation of Formula, Given Percentage Composition and Molecular Weight. C =40. 11% X 60 = 24.06 -=- 12 = 2 H = 6.82% X 60 = 4.og -=- 1=4 O = 53.07% X 60 = 31.84 -=- 16 = 2 Thus the formula of acetic acid is C2H402^6o. ORGANIC CHEMISTRY. 97 LABORATORY QUESTIONS. 1. Of what value is the estimation of the boiling and fusing points ? 2. In the determination of boiling points, why should the thermometer not touch the liquid? 3. What are the causes of inaccuracy in determinations of boiling and fusing points ? 4. Of what value is a knowledge of the optical activity of substances ? 5. Why is the detection of N often valuable ? 6. Deduce the formula for alcohol with a molecular weight of 46, when C = 52.18% ; H = 13.04% ; O = 34.78%. 7. Sulphuric acid has a molecular weight of 98. Deduce the formula when H = 2.04% ; S = 36.65% ; O =65.31%. 8. Chloroform has a composition C^ 10.04% ; H = .84%; Cl = 89.12%. What are some possible formulas ? How would its formula be decided? 9. What are some of the difficulties which keep compounds from having their formulas accurately determined? 199. Determination of the Flashing Point of Kerosene. Fill a 7-inch test tube one-third full of kerosene. Insert a ther- mometer, also a bent glass tube so arranged that air can be blown through the oil and bubbles of foam kept constantly upon its surface. Place the test tube in a beaker of water. Apply a gentle heat, so that the temperature rises i° in two or three minutes. At each rise of i", blow for a few seconds through the glass tube until foam stands on the oil, then apply a flame to the mouth of the tube. When the flame flashes down to the oil, the reading of the thermometer gives the flashing point. A second determination, with a fresh sample, will enable a more accurate result to be secured. The laws should require oil to have a flashing point above the maximum temperature of the atmosphere in the locality where it is used. 200. Preparation of Iodoform. In half a test tube of water put 2 CO. of alcohol and 3 c.c. of KOH. Warm and add G g8 LABORATORY MANUAL OF MEDICAL CHEMISTRY. crystals of metallic iodine, with constant stirring, until a fine yellow powder becomes visible. Filter the solution, and examine the crystals of iodoform, CHI3, on the filter paper. Reaction — (C2H5)OH + 4I2+ 6K0H = CHI3 + KCCHOg) + 5KI + 5H2O. Alcohol Iodoform Pot. Formate 20i-a. Preparation of Chloroform. In a generator put about 3 grams of chloral hydrate, cover with KOH solution, and warm and condense the vapor in a test tube immersed in a beaker of ice water. Some water will be condensed with a globule of pure chloroform, CHCI3, which will have the chloroform odor, and, when poured on the hand, will quickly evaporate. The purest chloroform is prepared by this method. It may also be prepared by the action of alcohol and bleaching powder. Reaction— CCI3COH + KOH = CHCI3 + KCCHOg) + H2O. chloral Chloroform Pot. Formate 20i-b. Tests for Chloroform. I. The Iso-nitrile Test. In a test tube take a little alcohol and KOH. Add a drop of chloroform and a drop of anihne. Warm, and the peculiar disagreeable, persistent odor of benzyl-iso-cyanide is developed. Reaction — CHCI3 + 3KOH + C6H5NH2 = QHjCy + 3KCI + 3H2O. 2. The Flame Test. Prepare and light a H generator. A rod, wet with ammonia, above the flame gives no fumes, and a copper wire in the flame does not color it, except momentarily. Through the thistle tube add a drop of chloroform. The CI escapes as HCl, and yields fumes of ammonium chloride, with the rod wet with ammonia, and the wire heated colors the flame green from the constant formation of CuClj. ORGANIC CHEMISTRY. 99 202. Preparation of Alcohol by Fermentation. Alcohol may be prepared from the fermentation of molasses or sugar, but the following process illustrates the usual commercial method. In a saucepan, or large evaporating dish, boil a quart of water. Into this slowly sift a mixture of flour and cornmeal, with constant stirring, until a thin, even mush is produced. Cool this to 60° C, and put into a large flask or bottle. Next pound up some dry malt (see 254). Add water and filter out the husks. Add the solution to the mush. Stir, and notice how rapidly the mush liquefies as the diastase of the malt converts starch to malt sugar. Next add some yeast to the solution. Close the bottle with a perforated cork, in which a delivery tube is so fitted that it dips into a beaker of lime water. Notice that COj is soon evolved, as indicated by the turbidity of the lime water. When the action nearly ceases, after about two days, examine a few drops of the solution under the microscope, and note the form of the yeast plant {Saccharomyces cerevisia). Next filter the solution into the distillation flask, and distil in the apparatus prepared by the demonstrator. Note the tem- perature at which distillation begins, and continue until the thermometer rises to about 96°. Test this distillate by the iodoform test, 204. This distillate can be shaken with quick lime, which removes water, unites with any acetic acid formed, and decomposes any etherial salts. It can then be filtered and redistilled. The strongest alcohol obtained by mere distil- lation is called "rectified spirits," and contains about 7% of water. 203. Absolute Alcohol. Heat a few grams of copper sul- phate (CuS04-5H20) until it loses its water of crystalliza- tion, and is white. Shake the powder with alcohol. After lOO LABORATORY MANUAL OF MEDICAL CHEMISTRY. Standing, if water is present, the blue color is restored to the salt. This test for water in alcohol affords a convenient method of preparing nearly absolute alcohol, but it contains some CuSO^ in solution, which, however, does not interfere with its use in preserving and hardening specimens, etc. It may be purified by distillation. Commercially, absolute alcohol is prepared by distilling alcohol from quicklime, adding a little metallic sodium and redistilling. 204. The Iodoform Test for Alcohol. Warm a portion of the suspected alcohol with KOH. Add crystals of I, and stir as long as the iodine is decolorized, or until a yellow powder appears. Cool and set aside. If alcohol were present, yellow crystals of iodoform, CHI3, precipitate, and can be seen by the microscope as hexagonal stars and ro- settes. Some other organic compounds, like acetone, alde- hyde, etc., yield the iodoform test, but do not cause confu- sion in testing potable liquids. 205. Detection of Alcohol in Beer, Etc. Distill a small quan- tity of beer, or any alcoholic solution under examination. Test the first portion of the distillate for alcohol by the iodo- form test, 204. Beer contains from 1-4% alcohol. The quantity of alcohol is usually determined from the specific gravity of a measured quantity of the distillate. LABORATORY QUESTIONS. 1. What would be a safe flashing point for Texas oil ? 2. Why is a lamp in which oil is low more liable to explode than when filled? 3. Is kerosene explosive ? Insert a lighted match into a beaker of oil, and see. 4. Give the chemical names of chloroform and iodoform? 5. What is a substitution product? 6. What relation does alcohol bear to an alkali? ORGANIC CHEMISTRY. lOI 7. For what purposes is absolute alcohol useful? 8. When alcohol is in contact with specimens, what change does it undergo ? How would you preserve a specimen in alcohol ? 9. What is the cause of the blue color of CUSO4? 10. Write the reaction showing how lime assists in the preparation of absolute alcohol? 11. How does the preparation of whisky differ from that of alcohol? 12. What are alcoholic " tears," observed in the distillation of alcohol, and why formed ? 206. Preparation of Ether. Into a generator put 20 c.c. of alcohol. Keep it cool, and slowly add 10 c.c. of strong H2SO4. Mix, gently warm, and catch the distillate in a test tube surrounded by a beaker of ice and water. The distillate will be ether, (C2H5)20, mixed with a little alcohol. Pour the distillate on the hand. Note the odor and the cold produced. Reactions— (C2H5)OH+H2S04 = (C2H5)HS04+H20, and Acid ethyl sulphate (C2H5)OH + (C2H5)HSO, = (C,H,)02+H3SO,. Alcohol Acid ethyl sulphate Ether 207. Preparation of Aldehyde. In a generator mix 20 c.c. of a solution of potassium bichromate, KgCrgOy and 2 c.c. of H3SO4. Cool the mixture. Next add 5 c.c. of alcohol, heat, and catch several c. c. 's of the distillate in a test tube, cooled by being immersed in a beaker of ice and water. The distil- late has the sharp odor of ethyl aldehyde, C2H4O, and is mixed with some alcohol. Reactions — KzCt^O^j and H2SO4 liberate O. Alcohol Aldehyde Test a few drops by 209-2 for acetic acid. There is none present. Set the remainder away, exposed to the air, for some 102 LABORATORY MANUAL OF MEDICAL CHEMISTRY. time. Aldehyde is slowly oxidized to acetic acid. Test the solution for acetic acid by 209-2, and prove its presence. Reaction — CjH^O + O ^ C2H4O2. ^dehyde Acetic acid This process illustrates the oxidation of alcohol to acetic acid, which is performed in one stage by the mother of vine- gar {bacterium aceti). 208. Tests for Aldehydes. 1. The Silver Test. Take a solution of AgNOg in a test tube. Add 3 drops of NH4OH and a few drops of an alde- hyde, as, for example, formaldehyde, CH2O, called formalin; a powerful non- poisonous preservative and antiseptic. Next warm the solution and set it aside. The aldehyde will reduce the silver nitrate and deposit a bright mirror of metallic silver on the sides of the tube. 2. The Potassium Test. To a solution of KOH add a few drops of an aldehyde and heat gently. A yellow alde- hyde resin, having a peculiar odor, is precipitated. 209. Tests for Acetates. 1. The Sulphuric Acid Test. Heated with H2SO4, acetates evolve fumes of acetic acid, H(C2H302). 2. The Ferric Chloride Test. FeClg forms in exactly neutral solutions a deep red liquid, due to the presence of red Fe(C2H302)3, which when boiled precipitates red oxace- tate of iron, FeOCCaHgOa). 3. The Acetic Ether Test. Strong solutions of acetates mixed with H2SO4 and a Httle alcohol, when heated evolve the peculiarly fragrant odor of ethyl acetate, or acetic ether, C2H5(C2H302). 210. The Detection of Mineral Acids in Vinegar. Take a dilute solution of methyl violet. Add a drop of acetic acid, or ORGANIC CHEMISTRY. lOJ Other organic acid. No change occurs. Now add a few- drops of acetic acid containing a minute amount of H2SO^> or any mineral acid. The color immediately changes from violet to blue, and to green when the mineral acid is in excess. 211. Preparation of Valerianic Acid. Into a generator put 20 c.c. of KgCrjO^ and 2 c.c. of H2SO4. Cool the mixture. Next add 7 c.c. of amyl alcohol. Heat, and catch the dis- tillate in a cool tube. The substance is composed of a mix- ture of valerianic aldehyde, CgHijO, valeric acid and water. Cool the generator, pour back the distillate, and distil again. The second distillate is purer valerianic acid, H(C5H902), with its characteristic odor and acid reaction. Reactions— CgHu(OH)+0=C5HioO-|-H20. Amyl Alcohol Valerianic Aldehyde CgHioO + = C5Hio02. Valerianic Aldehyde Valerianic Acid 212. Preparation of Oxalic Acid. Heat a mixture of 10 parts of strong HNO3 and 2 parts of cane sugar until ni- trous fumes cease to be evolved. On cooling, crystals of oxalic acid, HgCgO^, are deposited. On a large scale, oxalic acid is prepared by heating sawdust and soda. 213. Tests for Oxalates. 1. The Barium Test. BaClg slowly ppts. from neutral solutions a white ppt. of BaCgO^, slightly soluble in acetic acid, soluble in HNO3, HCl and NH.Cl. 2. The Calcium Test. CaClg slowly ppts. white CaCgO^, insoluble in acetic acid, soluble in HCl and HNO3. 3. The Silver Test. AgNOs ppts. white Ag2C204, soluble in HNO3 and NH4OH. 4. The Carbon Monoxide Test. Heated in a test tube with H2SO4, 104 LABORATORY MANUAL OF MEDICAL CHEMISTRY. oxalates without charring evolve CO, which, when ignited, burns at the mouth of the tube with its characteristic blue flame. Heated alone in a test tube, oxalates without charring evolve CO, and the white residue effervesces on addition of HCl. 214. Tests for Tartrates. 1. The Calcium Test. CaClj, after a few moments, ppts. white crystalline CaCC^H^Og). NH^Cl prevents this pre- cipitation. The ppt. is soluble in KOH, but reprecipitated on boiling. It is slightly soluble in acetic acid. 2. The Barium Test. BaCla ppts. white Ba(C^H406), soluble in HCl and NH^Cl. 3. The Silver Test. AgNOs ppts. in neutral solution white AgaCC^HiOo). This ppt. blackens on boiling, and is soluble in HCl and NH4CI. 4. The Heat Test. Tartrates char when heated alone or with H2SO4. 215. Tests for Citrates, I. The Calcium Test. CaCl2 in neutral solution precipi- tates perfectly, on boiling, Ca3(CgHg07)2. Unlike calcium tartrate, this calcium citrate is insoluble in KOH. 2. The Silver Test. AgNOa ppts. white Ag3(C6H507), and, unlike silver tartrate, this ppt. does not blacken on boiling. 3. To detect a Mixture of Tartrates and Citrates. Precipitate the mixture with CaCla, heat, filter and digest the well-washed ppt. with cold KOH. Dilute and filter. The filtrate contains the tartrate, which is precipitated on boiling. The undissolved residue on the filter con- tains the citrate. Dissolve this in NH4CI solution, filter, and the cal- cium citrate will precipitate when boiled. LABORATORY QUESTIONS. 1. How would you detect aldehyde in alcohol? 2. Can you mention an insoluble normal acetate? 3. Is acetic ether a proper name for ethyl acetate? 4. What is an ester? 5. What is an ether? What is an alcohol? ORGANIC CHEMISTRY. I05 6. What is an aldehyde? 7. How may you distinguish between HCl and H2C2O4 by AgNOs ? 8. What acid does a Seidlitz powder contain? Test one. 9. What substance is the best for use to preserve pathological speci- mens? 10. Show the steps in the change from alcohol to vinegar. 11. Of what value is brown paper or tea leaves added to dilute molasses in vinegar making? 12. How does oxalic acid aid in removing inkstains? 13. Why does wine sour when the bottle is not well corked ? 14. Is oxalic acid in pieplant poisonous? Why? 15. Is there any reason why vinegar should contain a mineral acid? 16. What relation does chloral bear to aldehyde? Will it give the mirror test for aldehydes, 208-1 ? Try it. 216. Preparation of Nitrous Ether. In a generator mix 1 CO. of H2SO4 and 2 c.c. of HNO3. Cool the mixture thoroughly and add 10 c.c. of alcohol. Add a few pieces of broken glass. Boil gently, and collect the distillate in a test tube cooled by ice water. Note the odor of ethyl nitrite or nitrous ether (C2H5)N02. A solution of this in alcohol forms sweet spirits of nitre. Reaction — 2(C2Hg)OH+HN03 = (C2H5)N02+C2H40+2H20. Alcohol Ethyl Nitrite Aldehyde 217. Preparation of Acetic Ether. In a generator mix 2 g. of dry sodium acetate, 10 c.c. of alcohol and 2 c.c. of H2SO4. Add a few pieces of broken glass. Heat gently, and collect the distillate in a test tube surrounded by ice and water. Note.the odor of ethyl acetate or acetic ether, (C2Hb)CC2H302), mixed with some alcohol. Reaction — 2(C2Hg)OH+2Na(C2H302)+H2S04=2C2H5(C2H502)+Na2S04+2H20 Alcohol Sodium Acetate Ethyl Acetate 218. Preparation of Amyl Acetate. In a generator mix 2 g. of dry sodium acetate, 5 c.c. of amyl alcohol and 2 c.c. of I05 LABORATORY MANUAL OF MEDICAL CHEMISTRY. H2SO4. Add a few pieces of broken glass. Heat gently, and collect the distillate in a test tube cooled by ice and water. Note the odor of the distillate or amyl acetate, CgHii(C2H302), like pears. It is an example of the fruit essences, and is used by confectioners. Reaction — 2(C5Hn)OH+2NafC2H302)+H2S04=2C5H„(C2H302)+Na2S04+2H20 Amyl Alcohol Sodium Acetate Amyl Acetate 219. Preparation of Amyl Nitrite. Put 3 c.c. of amyl alcohol in a test tube. Through this pass red nitrous fumes of N2O3 from a generator containing hot HNO3 and starch. Keep the alcohol cool by immersing it in ice water. Impure amyl nitrite is formed (C5Hii)N02. Note its characteristic chok- ing odor. Reaction— 2(C5Hii)OH+N203 = 2(C5Hii)N02+H20. Amyl Alcohol Amyl Nitrite 220. Preparation of Benzoic Acid. Place in a dry test tube a small piece of gum benzoin. Gently heat the tube, and note the small white needles of benzoic acid which sublime on the cool part of the tube. 221. Preparation of Benzene. Place in a test tube an intimate mixture of equal parts of benzoic acid and quicklime, and apply heat. Note the odor of the benzene vapor evolved, and compare it with the commercial article. On a larger scale these vapors may be easily condensed. Reaction— HCC^HsOa) + CaO = CgHe + CaCOg. Benzoic Acid Benzene 222. Preparation of Nitro-benzene. In a large test tube mix 10 c.c. of H2SO4 and 5 c.c. of HNO3. Cool the mixture, add a drop of benzene, shake, and cool the tube. Continue this process until 20 or 30 drops are added, keeping the tube cold. Next slowly pour the whole into a beaker of water. Impure ORGANIC CHEMISTRY. I07 nitro-benzene, CgHgNOg, will sink to the bottom in the form of a brownish yellow oil, called the "essence of mirbane," or, from its odor, artificial oil of bitter almonds. Reaction— CeHg + HNO3 = CgHsNOa + HOg. Benzene Nitro-benzene 223. Preparation of Aniline from Acetanilid. Powder a piece y of solid NaOH the size of a grain of corn. Mix this with ' ^ about I g. of acetanilid and put it into a test tube. Apply heat. Notice the oily globules that arise as the mixture melts. Sudden solidification will occur. Then tip up the test tube so that the aniline oil may run out. Catch it in another test tube containing about 3 c.c. of water. Keep heating the solidified substance as long as the oil continues to drip from the tube. In the second tube a good sized glob- ule of white aniline oil, CgHgNHa, is obtained. Note the odor, and compare it with that of the commercial article. Reaction — (C6Hg)(NH)(C2H30) + NaOH = CeHgNH^ + NaCC^HgOa). Acetanilid Aniline Sodium Acetate 224. Color Reaction of Aniline. Take half a test tube of water. Into it put a drop of aniline prepared in the above experiment. Shake well and add 3 drops of a clear solution of bleaching powder in water. A violet color develops, due to the formation of an aniline dye. 225. Preparat on of Rosaniline. Take about 2g. of HgCl2 in a test tube. Add 3 drops of aniline oil. Heat gently until the mass assumes first a green and then a dark purple hue. Cool the tube, and add a little alcohol and two drops of HCl. Stir it up and pour it into a beaker of water. The purple color resulting is due to the presence of rosaniline hydrochloride, one of the most important aniline dyes. I08 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 226. Tests for Carbolic Acid or Phenol, CgHgOH. 1. The Ferric Chloride Test. FeClg colors moderately concentrated aqueous solutions dark violet. 2. The Bromine Test. A drop of bromine, even in very dilute solutions, ppts. yellowish white tri-bromo-phenol, CeHgBrgCOH). 3. The Hypochlorite Test. A pinch of bleaching powder added to an ammoniacal solution of phenol yields a green coloration. 4. Certain Pine Shavings moistened with HCl, and touched with a solution of phenol, turn blue when exposed to the air for a time. 227. Preparation of Picric Acid. Take 5 c.c. of dilute HNO3 in a beaker. Add 2 c.c. of carbolic acid. Cool the mixture. Add 10 c.c. of strong HNO3, drop by drop. Boil the solution for five minutes, adding a little strong HNO3, drop by drop, if the dark, oily phenol seems to float after some boiling. Pour the whole solution finally into 25 c.c. of water. Set aside and decant the liquid from the yellow crystals de- posited. Wash these with a few drops of water. Decant, and notice the yellow solution of picric acid. Dry the crys- talline picric acid, C6H2(N02)3(OH). Notice that these dry crystals explode when dropped on a very hot surface or introduced into a flame. Reaction— CgHgCOH) + 3HN03 = C6H2(N02)30H + 3H2O. Phenol Picric Acid 228. Tests for Antipyrin. 1. The Ferric Chloride Test. FeCls turns solutions of antipyrin blood red. 2. The Nitrous Acid Test. Dilute HNO3, having dissolved in it a little KNO2, turns solutions of antipyrin green. 3. The Iodine Test. I dissolved in KI solution forms, in solutions of antipyrin, a brick-red ppt. ORGANIC CHEMISTRY. I09 LABORATORY QUESTIONS. 1. Why is sweet spirits of nitre apt to be acid? 2. How might nitrous acid be present in sweet spirits of nitre ? If possible, test for its presence by 93, and write the reactions showing its formation. 3. Why would old sweet spirits of nitre turn green on addition of antipyrin ? 4. To what class of chemical compounds do fruit essences belong ? Do they occur in nature ? 5. Ethyl butyrate is pineapple flavor. Write its formula, and invent a method for its preparation. See 237. 6. For what is benzene used? 7. What is the difference between benzene and benzine? 8. Of what two substances is acetanilid composed ? Outline a method of preparing acetanilid. 9. For what is amyl nitrite valuable? 10. What compounds of benzoic acid are used in medicine ? 11. Why is H2SO4 used in the preparation of nitro-benzene ? 12. Of what is smokeless powder composed? 13. Why does HNO3 turn the skin yellow? 14. Is nitro-benzene poisonous ? 15. Is carbolic acid an acid? Try litmus paper. 16. What is a carbolate? 17. What is the significance of the termination "ol " in phenol. 18. Is aniline poisonous ? FATS AND OILS. 229. Solution of Fats. Dissolve a small amount of the dif- ferent oils and fats found on the side table in small amounts of the following solvents : Boiling alcohol, benzol, gasolene carbon disulphide, ether, etc. 230. Emulsion of Fats. Shake a few drops of fat with half a test tube of water. Note that on standing the oil rapidly rises to the surface. Repeat the process with I. A few drops of alkali. A white emulsion is formed. no LABORATORY MANUAL OF MEDICAL CHEMISTRY. 2. A few drops of soap solution. A white emulsion is formed. 3. A solution of white of egg. A white emulsion is formed. 4. A solution of gum arabic. A white emulsion is formed. Examine an emulsion under the microscope. 231. Saponification of Fats — Hard Soap. Take 10 c.c. of castor oil in a beaker. Add an equal quantit}' of a solution of NaOH in alcohol. Stir it until a firm mass of soap is produced. This may be melted and poured into a mold, ■or used as it is. It will produce a good lather. When a solution of alkali in water is used, several hours' boiling is required. The same action occurs in a few moments with an alcoholic solution of alkali. Hence this is the quickest laboratory method. Illustrative reaction — C3H5(Ci8H3502)3+3NaOH = C3H5(OH)3+3Na(Ci8H3502). stearin Glycerine Sodium Stearate 232. Saponification of Fats — Soft Soap. Repeat the above process, using KOH in place of NaOH. A potassium, or soft soap, is formed. 233. Preparation of Insoluble Soaps. Make up a beaker full of a strong solution of the soft soap formed above. Half £11 several test tubes with this. Add the following reagents, and shake the tubes. 1. CaCl2 ppts. an insoluble white lime soap. No lather forms when shaken. 2. Repeat No. i, first adding a solution of NagCOs before CaCl2. The lime is first pptd. as CaCOg. No lime soap forms in the softened water, and it lathers when shaken. Reaction— CaCl2+Na2C03 = 2NaCl+CaC03. 3. MgS04 ppts. insoluble white magnesium soap. ORGANIC CHEMISTRY. Ill 4. Pb(C2H302)2 ppts. insoluble white lead soap (lead plaster). 5. FeClj ppts. insoluble brown ferric soap. 6. FeS04 ppts. insoluble greenish ferrous soap. 7. CUSO4 ppts. insoluble blue copper soap. 234. Preparation of Lead Plaster. Heat in an evaporating dish, with constant stirring, 3 parts of olive oil, 5 parts of water, and i part of litharge, PbO, until a yellowish white, tenacious mass of lead soap is obtained. It is the same substance formed iu 233-4. 235. Preparation of Fatty Acids. Boil in an evaporating dish a mixture of soft soap and water. Add HCl to excess. Reheat and set aside. Remove the light, fatty substance from the surface. Wash it in water. It is neither fat nor soap. It has no soapy taste. It is a mixture of fatty acids used in the manufacture of candles. Glycerine, KCl and HCl remain in the solution. 236. Preparation of Butter Soap. Put a teaspoonful of butter in a beaker. Boil it with a little dilute aqueous solution of KOH, adding repeatedly small amounts of alkali of increasing strength. After boiling and stirring for some time, remove from the fire and examine the soft butter soap. It will produce a lather. 237. Separation and Identification of Butyric Acid. Mix the butter soap formed above with a little water. Place the mixture in a generator. Add 10 c.c. of H2SO4. The fatty acid may be seen rising to the top. Heat, and collect 5 or 10 c.c. of the distillate in a cooled tube. It will contain butyric acid, have a rancid odor and an acid reaction to litmus paper. Heat this distillate with a little alcohol and 2 c.c. of H2SO4. Note the odor of ethyl butyrate, C2H5- 112 LABORATORY MANUAL OF MEDICAL CHEMISTRY. (C4H7O2), which is used as artificial pineapple flavor. "Oleo" treated thus contains only traces of volatile fatty acids in the distillate, and yields no ethyl butyrate. 238. Identification of Butter and "Oleo." The principle of the above reaction may be more rapidly applied thus : Place in two test tubes equal quantities of butter and oleo- margarine. Add 5 c.c. of a strong alcoholic solution of KOH to each. Warm gently, and notice the difference in odor. The "oleo" has the odor of alcohol simply, the butter the odor of pineapple, by the formation of ethyl butyrate from the alcohol and butyric acid. "Oleo" which has been churned with milk in the process of manufacture may yield slight traces of butyric ether. 239. Decomposition of Glycerine. Mix a few drops of glycer- ine and H2SO4 in a test tube. Heat, and note the sharp odor of acrolein. Any fat containing glycerine yields the same reaction. Reaction— (C3Hg)(OH)3=C3H40+2H20. Glycerine Acrylic Aldehyde 240. The Borax Test for Glycerine. Mix with a pinch of borax a few drops of suspected glycerine, which must be neu- tral and free from ammonium salts. Heated on a platinum wire, the mixture immediately tinges the Bunsen flame pale green, due to the liberation of boric acid by the glycerine. LABORATORY QUESTIONS. 1. Can castor oil be dissolved in benzine? 2. Can olive, castor and cod-liver oils be distinguished by being touched with H2SO4 ? 3. Why does an alkali emulsify fats ? 4. Why does albumin water emulsify fats ? 5. Does the character of the fat have any influence on the hardness of soap ? ORGANIC CHEMISTRY. II 3 6. Why does an alcoholic solution of an alkali more quickly saponify fats than an aqueous solution ? 7. Write the reaction involved in the formation of a calcium soap. 8. Why does hard water not lather soap ? 9. Why, when an alkali is added to water, does it often become soft ? 10. Why is one soap hard and another soft? 11. Is soap a salt? 12. Why is NaCl often added to a kettle of soap after it is made? 13. What lye is usually used in preparing "home-made" soaps? 14. Resin contains 3 fatty acids. Would it saponify with NaOH ? 15. What is lead plaster? 16. Which are more stable, the fats of " oleo " or butter? 17. Write the reaction involved in the preparation of fatty acids. 18. In what is oleomargarine superior to butter? 19. Chemically, what is a fat ? 20. What is glycerol ? What does the termination "ol " signify? CARBOHYDRATES. 241. Preparation of Corn Starch. Finely pulverize some corn in a mortar. Add water from time to time, and strain the milky liquid off through a cloth. Set it aside until the white sediment can be separated by decantation. Collect and dry this white corn starch, (CeHiQ05)n, on filter paper. 242. Appearance of Starch Granules. Examine under the microscope a thin section of potato. Notice how the starch granules lie packed within each cell. Touch the section with a drop of dilute iodine solution. Note that the granules turn blue, but the cellulose wall appears unchanged or slightly yellow. Examine the prepared microscopical specimens of potato, arrow root, corn, rice, and other starches, noting the difference in form and size of the granules. 243. Preparation of Starch Paste. Mix a little powdered starch with water to a thin milk. Pour this slowly into a beaker of boiling water with constant stirring. The milki- 114 LABORATORY MANUAL OF MEDICAL CHEMISTRY. ness disappears, and the whole forms apparently a translucent homogeneous solution. 244. The Iodine Test for Starch. A few drops of a solution containing free iodine strike with starch paste a deep blue color, from the formation of starch iodide. This color disap- pears on heating and reappears on cooling, unless the heat has been high enough to volatilize the iodine. , 245. Properties of Glucose. Examine commercial glucose, or grape sugar (dextrose). It is not as sweet as cane sugar. Make a solution of glucose, and set it aside for further experi- ments. Add H2SO4 to cold solutions of cane sugar and grape sugar. Heat to boiling, and note the charring of the cane sugar, while the glucose remains nearly unaltered. Repeat the experiment, using KOH. The cane sugar is only slightly affected, while the glucose assumes a dark brown color, due to the formation of caramel. 246. Fehling's Test for Glucose.^ In a test tube take some Fehling's glycerine solution. Heat it just to boiling, and add a few drops of liquid containing glucose. Again bring the mixture to the boiling point, and set it aside. If glucose was added in any appreciable quantity, a yellow or brick-red ppt. of cuprous oxide, CU2O, will be found in the bottom of the tube. When sufficient glucose is added the copper will all be thrown down and the supernatant liquid will be colorless. 247. Detection of Glucose in Fruits and Candies. Soak a bruised -raisin in water and test the solution for glucose. Dissolve a bit of candy in water and test the solution for glucose. Glu- cose will be present in both instances. 248. Action of Various Sugars on Fehling's Solution. Examine Note i. The composition of Fehling's glycerine test will be found under 348, Note i, also other tests for glucose under urinalysis, 346-353. ORGANIC CHEMISTRY. II5 the following sugars. Make a solution of each, and test as for glucose by Fehling's test. Verify the following conclu- sions : (i) Sucrose (cane sugar). No reducing power. (2) Dextrose (grape sugar). Reducing power. (3) Lactose (milk sugar). Reducing power. (4) Maltose (malt sugar). Reducing power. 249. Preparation of Invert Sugar. Make a dilute solution of cane sugar. Add a few drops of acid, and warm for lo min- utes. Test some of this solution for glucose. Cane sugar takes up a molecule of water and splits into two molecules, one of which is dextrose, or dextro-rotary, the other laevulose, or laevo-rotary. Reaction— Cj^HaaOii+HjO = 2(C6Hi206) Cane Sugar Invert Sugar 250. Action of Sugar on Polarized Light. Examine with the polariscope a solution of cane, malt, milk, or grape sugar. Note that they are dextro-rotary. Examine a solution of in- vert sugar, and decide on its action on the polarized ray. 251. Preparation of Barley Sugar and Caramel. In an evapo- rating dish very slowly heat some dry cane sugar. It melts to a clear yellowish liquid. When melted, stir and cool. Barley sugar is formed. Save a sample. Heat the remainder until it is of a dark brown color. Caramel is formed, which is sugar less a molecule or two of water. It is used for color- ing liquors, etc. Test solutions of both these products for glucose, and prove its presence. 252. Conversion of Starch to Dextrine. Heat a pinch of dry, finely powdered starch in an evaporating dish over a very slow fire, with continuous stirring until it all turns brown, but does not char. This occurs at about 250° C. Dextrine or British Il6 LABORATORY MANUAL OF MEDICAL CHEMISTRY. gum is formed, CgH^QOg. When cool, boil with a little water, and filter. The solution should be of a reddish color, and if all starch has been transformed to dextrine, the solution strikes a pinkish and not a blue color with iodine. This is an example of the splitting of the starch molecule into soluble starches or dextrines, which action is induced by heat, acids and digestive ferments. It may be continued through a series of products, ending at last in the formation of sugar, as seen in the next exercise, and in 282, under Saliva, where the color reactions of the intermediate products with iodine are given. _ 253. Conversion of Starch to Glucose. Boil a beaker of starch paste, prepared as in 243, with 10 c.c. of H2SO4 for an hour or more, adding water from time to time, if necessary, until a drop of it ceases to give a color with iodine, showing the conversion of all the starch. Add marble dust until effer- vescence ceases, which neutralizes any free acid remaining. Filter the solution. Notice its sweet taste. Test it, and prove the presence of glucose. It may be evaporated, and will sometimes crystallize. This is the method of preparing commercial glucose from corn, and is much used in syrups, candies, etc. 254. Conversion of Starch to Maltose. In a test tube warm to blood heat a dilute solution of starch paste. Pound up some malted barley — /. e., barley which has begun to sprout, and dried and heated until the germ is killed. Mix this with water. Grind it up and filter. Add some of this clear solution to the warm starch paste. After some minutes, test and prove the presence of maltose by Fehling's test. Prove the conversion of all the starch by a drop of iodine solution. Illustrative reaction— (C6Hio05)3+2H20=Ci2H220u+C6Hi206. Stai'ch Maltose Dextrose ^ 255. Preparation of Parchment Paper. Dip a piece of unsized ORGANIC CHEMISTRY. II 7 paper in strong H2SO4 for about 15 seconds, then wash it thoroughly with water. Notice that the paper has become stronger and more translucent. Its fibers have become col- loidal, and it may be used for dialysis, as it will allow only the molecules of crystalline substances to penetrate its pores, and the simpler proteid molecules, like peptones, etc. 256. Preparation of Collodion. Make a mixture of 15 c.c. of strong H2SO4 and 20 c.c. of strong HNO3 in a beaker. Cool the mixture and immerse in it a piece of shredded absorbent cotton. Macerate 15 minutes. Wash the cotton free from acids under the tap. Press it between filter paper and dry it. Notice that when thoroughly dry a portion burns in- stantly, leaving little ash. If the nitration has been exactly performed, di-nitro-cellulose is produced. This is pyroxylin, or soluble cotton. It dissolves in a mixture of 3 parts of ether and i of alcohol, forming collodion. Illustrative reaction — C6Hio05+2HN03=C6H803(N03)2+2H20. Cellulose Di-nitro-cellulose Gun-cotton is more explosive, and consists of a mixture of higher nitrates of cellulose, produced by more prolonged action of HNO3. It is insoluble in a mixture of alcohol and ether. Illustrative reaction— C6Hio05+3HN03=C8H702(N03)3+3H20. Cellulose Tri-nitro-cellulose LABORATORY QUESTIONS. 1. Can the source of starch be determined by the form of its granules ? 2. Write the formulas and show the relations of cellulose, starch and the various sugars. 3. What is a carbohydrate? 4. What is the nature of the change when cane sugar becomes glucose ? 5. Why is vinegar, cream of tartar, or lemon juice used in candy making ? 6. In how many ways may starch be converted into glucose ? 7. Ought glucose to be unhealthful ? Il8 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 8. For what is dextrine used ? 9. How does a potato sprout bring its insoluble starch up to its leaves ? 10. Why is maple sap sweet in the spring ? Why does it lose its sweet- ness when the leaves start ? 11. What does malt extract contain? Is its nutritive value diminished by fermentation ? 12. Could saw-dust be made available for food ? 13. Why do fruits sweeten as they ripen, and lose their sweetness when they decay ? 14. Why is invert sugar so called ? 15. What is malt? LABORATORY EXERCISES IN PHYSIOLOGICAL CHEMISTRY INCLUDING THE DIGESTION, BLOOD AND MILK PHYSIOLOGICAL CHEMISTRY. 121 PHYSIOLOGICAL CHEMISTRY. THE DETECTION OF PROTEIDS IN SOLUTION. 257. The Xanthoproteic Color Reaction. Take half a test tube of albumin solution. Add 5 c.c. of concentrated HNO3. Warm, and notice the yellow precipitated albumin. Divide in two test tubes. Add, until alkaline, to the one NH4OH, to the other KOH. The yellow solution turns orange in both instances. 258. The Biuret Color Reaction.^ Take half a test tube of albumin solution. Add two drops of dilute CUSO4 and 5 c.c. of KOH. A violet tint appears, which deepens on boiling. With peptones the color is rose-red. 259. Millon's Precipitation Reaction. Take half a test tube of albumin solution. Add a few drops of Millon's reagent. 2 A white ppt. falls, becoming reddish on boiling. If only traces of albumin are present, the solution becomes slightly red. 260. The Sodium Sulphate Precipitation Reaction. Take one- third of a test tube of albumin solution. Render it acid with acetic acid. Add an equal volume of concentrated Na2S04 solution. When boiled, a white ppt. falls. This test throws down all proteids except peptones. The salt and acid do not interfere with the tests for peptones in the solution. 261. The Ferrocyanide Precipitation Reaction. Render a Note i. Called Biuret because this color also forms from biuret, C2H5N3O2, a substance obtained by heating urea. Note 2. Millon's reagent is prepared by dissolving i part of metallic mercury in 2 parts of strong HNO3, adding twice its volume of distilled water, setting aside until settled, and decanting the clear solution. 122 LABORATORY MANUAL OF MEDICAL CHEMISTRY. solution of albumin strongly acid with acetic acid. Add a few drops of a fresh, strong K^FeCyg solution. A milk- white precipitate falls. This reagent ppts. all proteids except peptones and some forms of albumose. When acetic acid is added after the K4FeCyg, mucin is not precipitated. 262. The Tannic Acid Precipitation Reaction. Take half a test tube of albumin solution. Add i c.c. of a strong solu- tion of tannin. A white precipitate falls. Tannic acid reacts with all proteids. 263. The Picric Acid Precipitation Reaction. Take half a test tube of albumin solution. Add a few drops of a strong solution of picric acid. A yellow ppt. comes to view, dis- solving to a dark color when heated with KOH. Picric acid ppts. all • proteids. '' 264. The Absolute Alcohol Precipitation Reaction. Take one- third of a test tube of albumin solution. Render it acid with acetic acid. Add an excess of absolute alcohol. A ppt. falls. This reagent precipitates all proteids, including peptones. EXAMINATION OF THE VARIOUS PROTEIDS. CLASS I.— ALBUMINS. Albumins are proteids soluble in water and coagulated by heat. 265, Preparation of Egg Albumin. Take an egg. Break a small hole in one end of the shell. Pour the white into a beaker, leaving the yolk for future examination. Half fill the beaker with water. Stir with a glass rod, breaking up the albumin thoroughly. Filter through a piece of muslin, and keep the solution for the following reactions. 266. Precipitation of Albumin by Heat. Boil half a test tube PHYSIOLOGICAL CHEMISTRY. 1 23 of albumin water. Notice the coagulation. The coagulum does not dissolve on addition of HNO3. 267. Temperature of Coagulation. Extract a little more of the undiluted white of egg from the shell. Place it in a test tube, insert a thermometer, and set the whole in a beaker of water. Warm the beaker slowly. Note the tem- perature at which the albumin shows the first signs of opal- escence (about 59° C), and the point of total coagulation (about 73° C). 268. Precipitation of Metallic Albuminates. Prepare three test tubes of albumin solution. Add to separate tubes solu- tions of CUSO4, AgNOg, and HgCl2. Copper, silver and mercury albuminates fall. Compare them with the soap ppts. formed in 233. CLASS 11.— PEPTONES. Peptones are proteids soluble in water, but not coagulated by heat. Albumoses are substances between albumins and peptones. 269. Action of Heat and Reagents. Using the prepared solution of peptones, confirm the following reactions : -• I. Not coagulated by heat. -■z. Not pptd. by adding NaCl. 3. Not pptd. by acids or alkalies. 4. Not pptd. by Na2S04, as in 260. 5. Not pptd. by K^FeCyg, as in 261. 6. Pptd. by tannic acid, as in 262. 7. Pptd. by absolute alcohol, as in 264. 8. Yields the rose-red Biuret reaction, 258. CLASS III.— GLOBULINS. Globulins are proteids insoluble in water, soluble in dilute NaCl. Solutions are coagulated by heat. Among the most important are : 124 LABORATORY MANUAL OF MEDICAL CHEMISTRY. Vitellin, crystallin (globulin), myosin, fibrinoplastin (paraglobulin), and fibrinogen (metaglobulin). 270. Preparation and Detection of Vitellin. Stir up the yolk of the egg saved, after washing it as carefully as possible from the white. Put 3 c.c. of yolk in a test tube. Shake well with half a test tube of ether several times. Allow the yellowish ether to rise each time, and pour it into an evapo- rating dish. The ether largely dissolves the fatty matters. Impure vitellin remains. Very gently warm the tube, and set it aside until the smell of ether disappears. Add water. The vitellin does not dissolve. Next add a pinch of NaCl. Vitellin dissolves to a milky solution. Filter, and apply the Heat (266), the Xanthoproteic (257), and the Biuret (258), reactions. 271. Examination of the Fatty Matter of the Yolk. Evapo- rate the ether solution in the evaporating dish very cau- tiously, lest the ether take fire. In case it should, cover the dish quickly with a damp cloth, which will extinguish the flame. A yellow oil is left. Pour a few drops into water. Globules of fat are formed. Heat the remainder with a few drops of HNO3. The solution is first colored blue, then fades to green, and becomes colorless. Now add a little water and a drop of KCyS solution. A reddish color proves the presence of iron. 272. Preparation and Detection of Crystallin. Extract the crystalline lens from the eye of an ox. Thoroughly grind it up in a mortar with 10 c.c. of water, which will dissolve albumin. Filter, and grind the residue with 10 c.c. of a 10% salt solution. The crystallin dissolves. Filter, and confirm its presence by Millon's test (259), the Biuret test (258), and the Heat test (266). 273. Preparation and Detection of Myosin. Soak a table- PHYSIOLOGICAL CHEMISTRY. I25 spoonful of finely chopped lean meat in a beaker of water for 10 minutes. Squeeze out the juice through a wet cloth, and test the liquid for albumin by Millon's reaction (259.) Wash the cloth. Repeat the digestion several times, until the water no longer yields much, if any, ppt. with Millon's reagent, showing the removal of serum albumin. Now soak the washed meat in a 10% salt solution. Strain the liquid. Filter, and test for myosin by Millon's reagent and a few other tests for proteids (257-264). Fibrinogen and Fibrinoplastin (310), may be similarly pre- pared, but the process is too tedious for the time allowed. CLASS IV.— DERIVED ALBUMINS. Derived albumins are proteids insoluble in water, soluble in dilute acids or alkalies. Solutions not coagulated by heat. 274. Preparation of Acid Albumin. Take half the white of an egg. Add 10 drops of glacial acetic acid. Whip them well together for some minutes until a gelatinous mass of acid albumin forms. Dissolve this in half a beaker of warm water. Filter, and confirm the following reactions for solu- tions of acid albumin : 1. Not coagulated by heat. 2. Precipitated when exactly neutralized by NaOH. 3. This last ppt. is dissolved by an excess of alkali. 275. Preparation of Alkali Albumin. Take half the white of an egg. Add 15 drops of strong KOH solution. Whip them well together until a gelatinous mass of alkali albumin is formed. Dissolve this in warm water. Filter, and confirm the following reactions for solutions of alkali albumin : 1. Not coagulated by heat. 2. Precipitated when exactly neutralized by HCl. 3. This last ppt. is dissolved by an excess of acid. 126 LABORATORY MANUAL OF MEDICAL CHEMISTRY. 276. Preparation of Syntonin (muscle acid albumin). Wash a tablespoonful of lean chopped muscle until proved by Millon's reagent to be free from serum albumin, as in 273. Digest the washed meat in- a o.i