GIFT OF Prof. M. E. Jaffa C^C^y]AJ^^ PRACTICAL CHEMISTRY. THE OWENS COLLEGE JUNIOR COURSE PRACTICAL CHEMISTRY BY FRANCIS JONES, F.R.S.E., F.C.S, Chemical Master in the Grammar School, Manchester WITH A PREFACE BY SIR H. E. ROSCOE, F.R.S. MACMILLAN AND CO. AND NEW YORK 1889 The Right of Translation and Reproduction is Reserved 357 RICHARD CLAY AND SONS, LONDON AND BUNGAY. First Edition printed 1872. Reprinted 1873, 1875, 1877, 1878, 1880, 1882, 1884, 1885, 1887, With Addition 1880. PREFACE. THIS little book contains a short description of a course of practical chemistry, which an experience of many years has proved suitable for those com- mencing the study of the science. The work is, however, not intended to supplant, but rather to sup- plement instruction given by the Teacher, as it points out the principles of the various processes and names the various reactions without entering into the details of the necessary manipulation, which after all can only be satisfactorily learnt by practical demonstra- tion. The subject-matter has been carefully compiled, under my supervision, by Mr. Francis Jones, F.R.S.E., formerly Junior Demonstrator in this Laboratory, and many improvements in the analytical tables are due to Mr. Schorlemmer, F.R.S.. the Senior Demonstrator. To those about to use the bcok, I would give the following advice :- 769795 vi PREFACE. 1. Be sure you understand the theoretical explana- tion (for which you may have to refer to other books) as well as the practical part of the experiment, or the reaction, which you perform. 2. Keep careful notes of each day's laboratory work, and write out answers to the questions found at the end of these pages. 3. One of the first virtues in the practical chemist is cleanliness. Learn to work neatly, and you will soon obtain exact views of the science. Those who work in a mess not unfrequently get their minds in a muddle. R ROSCOE . THE OWENS COLLEGE, MANCHESTER, October, 1872. PREFACE TO THE SEVENTH EDITION. To this edition I have added a series of examples on quantitative analysis, forming Part VI. of the book. I have chosen only such as are capable of easy and accurate determination, and trust they will serve as a sufficient introduction to the more complete study of the subject. FRANCIS JONES. MANCHESTER, August, 1882. CONTENTS. PART I. PAGE Preparation of Apparatus, Gases, &c. I PART II. Blow-pipe Analysis 41 Bunsen's Flame Reactions , . <. 46 Preliminary Examination 55 PART III. Grouping of the Metals 60 Reactions of the Metals of the Silver Group .... 63 Separation ,, ,, ,, .... 65 Reactions ,, ,, Copper Group ... 66 Separation ,, ,, ... 7 Reactions ,, ,, Arsenic Group ... 72 Separation ,, ,, ,, ... 80 Reactions Iron Group . . . 81 Separation ,, ,, ,, ... 88 Reactions Barium Group ... 90 Separation ,, ,, ,, ... 92 Reactions ,, Potassium Group . . 94 Separation ,, ,, ,, ... 96 viii CONTENTS. PART IV. PAGH Grouping and Reactions of the Inorganic Acids ... 97 Separation of Inorganic Acids 115 Grouping and Reactions of the Organic Acids . . .118 Tests for Organic Alkaloids 125 Separation of Organic Acids ......... 134 PART V. Grouping and Reactions of the Rare Metals . . . . 136 PART VI. Examples of Quantitative Analysis 15 APPENDIX. A. Table of the Elements '73 B. Table of Weights and Measures . 176 C. Treatment of Silver Residues . ...... 177 D. Treatment of Platinum Residues 180 Questions and Exercises l l INDEX. '99 JUNIOR COURSK, ,QF ; , LABORATORY PRACTICE. PART I. PREPARATION OF APPARATUS, GASES, &c. 1. Glass stirring-rods. Divide a piece of glass rod into several pieces about t\vo decimetres in length. This is done by filing the glass rod at each place where it is to be cut off, with a three-cornered file, and then snapping it across. Knock off any projecting pieces of glass which may be left at the newly-cut edges, and then hold each end of the rod in the flame of a Bunsen lamp untrl the sharp edges are fused and rounded. The glass rods thus made serve for stirring liquids. &c. 2. A wash bottle. Soften a cork* by gently rolling it under the foot, and fit it air-tight into the neck of a flask about one litre capacity. Then, by means of a round file, bore two holes in the cork about three millimetres in diameter, * A doubly-bored india-rubber stopper may be used instead of a cork. 2 PRACTICAL CHEMISTRY. and running parallel to each other and to the longer axis of ,the cpr]<. Next obtain two pieces of glass tubing oi', the .same diameter, one three decimetres long and the.otte* half that length. Hold one end o the longer tube in the Bunsen flame until the opening contracts considerably (but take care not to seal it up entirely), and then bend it about half a decimeter from the end, as shown in the figure. This is done by holding the glass tube horizontally in a common batswing gas jet flame, turning it round so as to heat all parts equally, and bending to the pro- per angle as soon as it feels sufficiently soft. Now round the edge of the wide end of the tube by holding it in the flame till it softens, and when cold fit it into the cork. In a similar way round both ends of the shorter piece of tubing, bend to the angle shown in the figure, and fit into the other hole. Clean out the flask and tubes thoroughly, rinse with distilled water, and then fill up with distilled water. 3. Preparation of oxygen from mercuric oxide. When mercuric oxide is heated it is decomposed into mercury and oxygen. Mercuric oxide yields mercury and oxygen. * HgO = Hg + O 216 == 200 + 16 * For the explanation of these symbols a larger work must be con- sulted. Seep. 13 of Roscoe's " Elementary Chemistry." PREPARATION OF OXYGEN. 3 Place a. small quantity of mercuric oxide in a dry test tube, and heat it over a Bunsen lamp. The sub- stance darkens in colour, and a ring of minute globules of mercury soon forms on the cool part of the tube. That the tube contains oxygen may be shown by plunging a glowing chip of wood into it, and observing that the wood will be rekindled. It is advisable to keep the thumb loosely on the mouth of the test tube to prevent the escape of oxygen by diffusion. 4-. Preparation of oxygen from potassium chlorate. Potassium chlorate when heated is decomposed into oxygen and potassium chloride. Potassium chlorate yields potassium chloride and oxygen : KC1Q 3 = KC1 + O 3 . I22'6 = 74'6 -f 48. Place a few crystals of potassium chlorate in a dry test tube, and heat gently. The salt soon fuses and then begins to effervesce, giving off oxygen, which may be recognized, as before, by its power of rekindling a glowing chip. When oxygen ceases to be evolved, the residue is a white salt called potassium chloride. 5. Tests for potassium chlorate and chloride. A solution of potassium chlorate is not precipitated by a solution of silver nitrate ; but potassium chloride is precipitated, silver chloride being formed. Potassium chloride and silver nitrate yield potassium nitrate and silver chloride : KC1 + AgN0 3 = KNO 3 + AgCl. 74-6 + 170 == ion + i43'S- B 2 4 PRACTICAL CHEMISTRY. Dissolve a crystal of potassium chlorate in distilled water, then add solution of silver nitrate, no precipitate will be formed. Dissolve also the residue of potassium chloride (obtained in 4-) in distilled water, and add silver nitrate solution. A curdy white precipitate of silver chloride will at once be formed. 6. Preparation of oxygen from potassium chlorate and manganese dioxide. Potassium chlorate when mixed with about one-fifth of its weight of manganese dioxide, gives off its oxygen at a much lower temperature than when heated although the oxide itself remains quite unaltered. Place in a small flask (about 100 cubic centimetres capacity) a mixture of potassium chlorate and manga- nese dioxide in the proportions already mentioned. Then fit into the neck of the flask a cork through PREPARATION OF OXYGEN. 5 which a bent conducting tube passes, the lower end of which is placed under water in the pneumatic trough. Fill some gas jars with water and invert them in the trough. Support the flask on a retort stand, and heat the mixture (Fig. 2) ; as soon as heat is applied, bubbles of gas will begin to rise through the water, consisting of air expelled from the flask by the heat. These are allowed to escape, and the oxygen, which soon begins to come off abundantly, is collected by placing the in- verted bottles over the end of the conducting tube, and thus allowing the bubbles of oxygen to ascend into the bottles and displace the water. As soon as the first bottle is filled with oxygen, place a shallow earthen- ware tray under the neck, and remove it from the trough, taking care that the tray contains enough water to prevent the escape of the gas. When four bottles have been filled in this manner, remove the flask and conducting tube and allow the former to cool. 7. The residue from the preparation of oxygen from potassiitm chlorate and manganese dioxide is potassium diloride, and ^maltcred manganese dioxide. Dissolve in water the residue in the flask from the preparation of oxygen, filter from the insoluble manga- nese dioxide, and evaporate the filtrate in a porcelain basin to a small bulk ; on cooling, crystals o f ootassium chloride will separate out. Pour off the mother liquor, dry the crystals between filter paper, and keep them in a small bottle for future experiments. 8. Combustions in oxygen. When a taper is burned in oxygen, the carbon which 6 PRACTICAL CHEMISTRY. it contains unites with the oxygen, forming carbon di- oxide (carbonic acid gas), and the hydrogen also unites with the oxygen, forming hydrogen monoxide (water). Carbon and oxygen yield carbon dioxide : C + 2 = C0 2 . 12 + 32 = 44. Hydrogen and oxygen yield water : H 2 + O = H 2 0. 2 + 16 = 18. Place a small piece of taper on an iron wire support, Sight it, and plunge it into one of the jars of oxygen (prepared in 6). Observe that it burns much more brightly than in air. Withdraw it, blow the light out, and if only a spark remain on the wick, observe that when again plunged into the oxygen it is at once rekindled. This serves as a convenient test for the presence of oxygen. When solution of calcium hydrate (lime water) is placed in contact with carbon dioxide, calcium carbonate (chalk) and water are formed, and the former, being insoluble in water, is precipitated as a white powder. Calcium hydrate and carbon dioxide yield calcium carbonate and water : Ca(HO) 2 -f CO 2 = CaCO 3 +' H 2 O. 74 + 44 = loo + 18. Prove that the jar in which you have burned the taper contains carbon dioxide, by adding a little clear lime water to it and shaking so as to bring the gas and liquid in contact. A white precipitate will be formed, consisting of calcium carbonate. COMBUSTIONS IN OXYGEN. 7 9. When charcoal is burned in oxygen carbon dioxide is produced. (Compare 8.) Place a few pieces of charcoal, about the size oi peas, in the deflagrating spoon, hold this in the lamp- flame till the charcoal is just kindled, and then plunge the spoon into a jar of oxygen. The charcoal will burn with great brilliancy, and, if enough oxygen be present, nothing will remain on the spoon but a little white ash (the inorganic matter in the charcoal). Prove the presence of carbon dioxide (as in 8) by adding lime water. A similar union of carbon and oxygen takes place in animals when breathing, but the combustion is slow, not rapid. A portion of the carbon of their bodies unites with the oxygen which they inhale from the air, and carbon dioxide is .produced. Observe this by blowing the air from the lungs through a glass tube into some clear lime water, a white precipitate of cal- cium carbonate soon forms, caused by the union of the carbon dioxide of the breath with the lime. (See Roscoe, p. n.) 10. When sulphur is burned in oxygen, sulphur lioxide (sulphurous anhydride} is produced. Sulphur and oxygen yield sulphur dioxide : S + 2 = S0 2 . 32 -f 32 = 64. Place a few pieces of sulphur in the deflagrating spoon, heat until the sulphur is melted and takes fire, and then plunge it into a jar of oxygen. Observe that the sulphur, which only burned feebly in air, burns 8 PRACTICAL CHEMISTRY. with considerable brightness in oxygen. When it has ceased to burn, remove the spoon, and observe the suf- focating odour of the gas which has been produced. Free acids redden blue vegetable colouring matters. Show that sulphur dioxide when dissolved in water forms an acid, and on pouring a solution of blue litmus into the wet bottle in which the sulphur was burned, observe that the blue colour is changed to red. 1 1 . When phosphorus is burned in oxygen, phosphorus pentoxide (or phosphoric anhydride'} is produced. This is not a gaseous body like the dioxides of sul- phur and carbon, but is a white solid substance, which very readily unites with water, forming tri-hydrogen phosphate (tribasic phosphoric acid). Phosphorus and oxygen yield phosphorus pentoxide : P 2 4- 5 = P 2 6 . 62 -f 80 = 142. When the oxygen has been collected over water, there is always enough moisture in the gas-jar to unite with the pentoxide, and form a solution of phosphoric acid (H 3 P O 4 ), thus : Phosphorus pentoxide and water yield tri-hydrogen phosphate : P 2 5 + 3 H 2 = 2 H 3 P 4 . 142 + 54 = 196. Place a small piece of phosphorus (having care- fully dried it with filter-paper) on the deflagrating spoon, light it by touching it with a hot wire, and place it in a jar of oxygen. Observe the intensely brilliant light with which it burns, and the dense white PREPARATION OF HYDROGEN. 9 fumes of phosphorus pentoxide which are produced. Observe too that in a short time these fumes disappear for the reason already mentioned ; and show that an acid is present by pouring in some litmus solution and observing the change of colour from blue to red. 12. Presence of oxygen in air. Oxygen is present in the air diluted with another gas called nitrogen. In one hundred volumes of air, about twenty-one consist of oxygen, and seventy-nine of nitrogen. When substances burn in air they unite with the oxygen in it and produce oxides, just as they do when burned in the pure gas. Burn a taper, sulphur, and phosphorus in three jars of air, and prove that carbon dioxide, sulphur dioxide, and phosphorus pentoxide are respectively produced. Ascertain this by applying the same tests as in 8. 1O, and 11. 13. Preparation of hydrogen. By the action of acids on certain metals, such as zinc, magnesium, and iron, hydrogen gas is evolved, a salt of the metal with the particular acid employed, being formed at the same time, thus : Zinc and sulphuric acid yield hydrogen and zinc sulphate : Zn + H 2 S O 4 = H 2 -|- Zn S O 4 65-2 -f 98 = 2 -f- 161-2; and iron and hydrochloric acid yield hydrogen and ferrous chloride Fe -f 2 H Cl = H 2 + Fe C1 2 56+ 73 = 2 -f 127. ro PRACTICAL CHEMISTRY. Place a few pieces of zinc in a wide-mouthed flask provided with a cork, through which pass two tubes, one a bent conducting tube for leading the gas into the pneumatic trough, the other a funnel tube for sup- plying acid to the zinc. Fill four jars with water, and collect the gas (which is given off without the aid of heat) in the same way as in the case of oxygen (Fig. 3). Having placed the end of the conducting tube in FIG. 3. the pneumatic trough, pour down the funnel lube enough water to cover the zinc completely, and then add sulphuric acid by degrees and cautiously, shaking the flask so as to mix the acid and water. The gas will soon come off with increasing rapidity, but the first portion must be rejected, as air and hydrogen form an explosive mixture. It is therefore necessary to wait until all the air has been expelled from the flask by the hydrogen. Before filling the jars, collect a test- tube full of the gas over the pneumatic trough, and apply a light to it (holding the mouth downwards) ; if EXPERIMENTS WITH HYDROGEN. \ \ it burn quietly, you may safely proceed to fill the gas jars, but if a slight explosion occur, you must wait until another test-tube is collected and found to burn quietly. As soon as this is the case, fill four jars with the gas, and remove them from the trough in the same way as described in 6. 14. Production of water by the burning of hydrogen in air. Whenever hydrogen burns in air or in oxygen, water is produced. Hydrogen and oxygen yield water : H 2 + O = H 2 2 -f 16 = 1 8. Remove the conducting tube from the hydrogen flask above described (but without removing the cork), and substitute for it a straight piece of glass tube of the same diameter, drawn out at the upper end so as to form a jet (Fig. 4). As all the air has long ago been expelled, the hydrogen issuing from the jet may be safely lighted. Hold over this flame a per- fectly dry beaker, or test-glass, and observe the instant deposit of mois- ture on the sides of the glass. This is caused by the union of the hy- drogen with the oxygen of the air contained inside the beaker. 15. The residue from the preparation of hydrogen from zinc and sulphuric acid is zinc sulphate. 12 PRACTICAL CHEMISTRY. Filter a portion of the contents of the flask used for preparing hydrogen from the undissolved zinc, and evaporate the nitrate to a small bulk. On cooling, crystals of zinc sulphate will be formed, which are drained from the mother liquor, dried between filter paper, and kept in a corked tube or bottle for future examination. 16. Hydrogen is an inflammable gas. Apply a light to one of the jars of hydrogen, and observe that it burns. Notice also the deposit of moisture on the sides of the gas jar caused by the union of the hydrogen with the atmospheric oxygen. 17. Hydrogen is lighter than air. Show this by pouring the contents of one jar, upwards, into a jar of air held above the mouth of the hydrogen jar. The hydrogen in virtue of its light- ness will leave the lower jar and displace the air from the upper jar. Apply a light to each jar, and observe that the one originally filled with hydrogen no longer contains any, while the other originally filled with air contains hydrogen. 18. Hydrogen does not support combustion. Take another jar of hydrogen, held mouth down- wards, and push up into it a lighted taper, supported on a straight wire. The hydrogen will burn at the mouth of the jar, but the taper will be extinguished. The taper may be withdrawn, relighted, and re-extin- guished two or three times. 19. Preparation of nitrogen. When phosphorus is burned in air, it unites with the PREPARA TION OF NITROGEN. 13 oxygen, forming phosphorus pentoxide (P 2 O 6 ), which dissolves in water, whilst the residue is nitrogen. Place a small piece of dry phosphorus on a porcelain crucible lid, and fix this on a flat piece of cork, so that the lid may be floated on the water in the pneumatic trough, or other tray containing water. Light the phos- phorus, and quickly surround it with a bell-jar the tubulure of which is closed with a cork (Fig. 5;. A FIG. 5. portion of the air will first be expelled by the heat, and the phosphorus will continue to burn as long as any oxygen remains in the air of the bell-jar ; when that is exhausted it will cease to burn. Observe that the water rises inside the bell-jar, replacing the oxygen, which is no longer gaseous, but has combined with the phos- phorus to form phosphorus pentoxide. Wait a few minutes to allow the pentoxide to be completely dis- solved by the water. The bell-Jar then contains nitrogen. 2O Nitrogen does not support ~comb^tst^on, and is not combustible. Place a lighted taper in the bell-jar, and observe i 4 PRACTICAL CHEMISTRY. that the flame is at once extinguished, and that the gas does not burn. 21. Preparation of nitric acid. All nitrates when heated with sulphuric acid are decomposed, nitric acid and a sulphate being formed. Thus, when sulphuric acid acts on potassium nitrate (nitre or saltpetre), nitric acid, and hydrogen potassium sulphate (bisulphate of potash) are produced : K N O 3 -f H 2 S O 4 = H N O 3 + H K S O 4 ici'i -\- 98 = 63 -f 136'!. Put about twenty grams of potassium nitrate (nitre) in a small stoppered retort, place in the tubulure FIG. 6. a small funnel, and through this pour the same weight of strong sulphuric acid. Withdraw the funnel without soiling the neck of the retort, place the stoppex in the tubulure, and support the retort upon a piece of wire TESTS FOR NITRIC ACID. 15 gauze placed across the ring of the retort stand. Let the neck of the retort pass into a clean flask, so ar- ranged that a stream of water may fall on it, and thus aid the condensation of the acid (Fig. 6). Now heat the retort, and observe the formation of red fumes (lower oxides of nitrogen) and soon after, the con- densation of the nitric acid on the neck of the retort. When the contents of the retort cease to boil, with- draw the lamp, separate the receiving flask containing the nitric acid from the retort, and pour out the con- tents of the latter into a dry porcelain dish. When cold, break up the cake of hydrogen potassium sulphate thus formed, and preserve it for future examination. 22. Tests for nitric add. (a] Add a few drops of the nitric acid prepared as above to a solution of indigo contained in a test-tube. The blue colour will speedily disappear owing to the oxidizing action of the nitric acid. () Place a few bits of copper turnings in a test- tube, cover them -with water, and add a little nitric acid. The copper soon begins to dissolve, forming a blue solution, and at the same time brown vapours fill the test-tube. In this case also the nitric acid acts as an oxidizing agent, forming copper nitrate, and red fumes of the oxides of nitrogen are given off.- For explanation see 3O and 31. (c) Add a few drops of nitric acid to a little water contained in a test-tube, and then add some strong sulphuric acid, and shake until the liquids are tho- roughly mixed Allow this mixture to cool completely, 1 6 PRACTICAL CHEMISTRY. and then pour gently on to the surface of the liquid a solution of ferrous sulphate prepared by dissolving a few crystals of the salt in water. This solution is lighter than the other, and if poured on gently will form a layer of liquid resting on the heavier sulphuric acid, and a black ring will form where the two liquids meet. This is caused by the liberation of nitrogen dioxide (by the action of the FeSO 4 on the HNO 3 ), which forms a dark-coloured compound with the ferrous sulphate. On shaking the tube, nitrogen di- oxide will escape with effervescence, and the black ring will disappear. Retain a portion of nitric acid for 27. 23. Tests for nitrates. The same reactions may be used to test for com- bined nitric acid (i.e. nitrates), but sulphuric acid must first be added to liberate it. Repeat the tests rt, , and c (22), substituting for the nitric acid a solution of potassium nitrate in water, to which a few drops of strong sulphuric acid have been added. 24. Preparation of ammonia. When caustic alkalies act on salts of ammonium, ammonia gas is liberated. Thus : Ammonium chloride and caustic lime yield am- monia, calcium chloride, and water : 2 (NH 4 ) Cl + CaO = 2 NH 3 + CaCl 2 + H,O 107 + 56 = 34 + TIT -f 18. Place about ten grams of powdered ammonium chloride (sal-ammoniac) and the same weight of PREPARATION OF AMMONIA. 17 powdered lime in a test tube provided with a tightly- fitting cork, through which passes a tube bent as shown in the figure. Place over this an inverted dry gas-jar, and heat the mixture in the' test tube. After the air has been expelled, the ammonia gas will come over, and in virtue of its lightness be retained in the jar. When the jar is filled, remove it from the Fia. 7. upright tube, place the palm of the hand on the mouth, and immerse the jar (mouth downwards) in the pneu- matic trough filled with water : the gas will be rapidly absorbed, and the water will rise so as to fill the jar. 1 8 PRACTICAL CHEMISTRY. Now reverse the conducting tube, and allow the remaining gas to pass into some distilled water contained in a beaker. The gas bubbles will be completely absorbed, and when the gas ceases to be evolved, be careful to withdraw the conducting tube from the liquid, so as to prevent it being sucked back into the hot tube. 25. Tests for ammonia. (a) Observe the pungent and very characteristic smell which the solution of ammonia in water possesses. (fr) Place a piece of reddened litmus paper above the solution, and observe that the vapour given off is able to change the colour from red to blue. Place the litmus paper in the solution, and the change will be still more marked. (c) Hold a glass rod which has been dipped in fuming hydrochloric acid over the solution of ammo- nia : it will at once form white fumes of ammonium chloride, caused by the union of the acid and 'volatile alkali,, as ammonia is sometimes called. Hydrochloric acid and ammonia yield ammonium chloride : HC1 + NH 3 = NH 4 C1. 36-5 + 17 = 53-5. 26. The residue from the preparation of ammonia Jrom ammonium chloride and lime is calcium chloride. Dissolve a portion of the residue contained in the flask used to prepare ammonia, in water, filter from excess of lime, and evaporate the solution to dryness. A white salt results, which is detached from the evapc- PREPARATION OF NITROUS OXIDE. \$ rating basin, and preserved in a well-stoppered bottle for future examination. 27. Preparation of ammonium nitrate. Nitric acid, when neutralized with ammonia, yields ammonium nitrate and water : HN0 8 + (NH 4 ) HO = (NH 4 ) NO 3 + H 2 O. 63+35 = 80 + 1 8. Place some nitric acid (21) in an evaporating basin, and dilute with twice its bulk of water j add am- monia solution (24) cautiously, and with constant stir- ring until a drop of the liquid ceases to colour litmus paper red. Observe that as the ammonia is added the red colour caused by the nitric acid disappears, and a point is reached when the liquid neither reddens litmus paper nor turns it blue. When this is the case the acid is said to be exactly neutralized, and the addition of ammonia must be stopped. Evaporate the solu- tion of ammonium nitrate thus obtained, until all the water is expelled ; when this is the case, withdraw the lamp and allow to cool ; a cake of fused ammonium nitrate is then obtained. If the salt does not solidify on cooling, heat must be again applied. 28. Preparation of nitrogen monoxide (nitrous oxide). When ammonium nitrate is heated, nitrogen mo- noxide and water are produced. Thus : (NH 4 ) NO 3 = N 2 + 2 H 2 O. 80 = 44 + 36. Break the cake of ammonium nitrate (27) into small pieces, and introduce them into a small dry flask C 2 ib PRACTICAL CHEMISTRY. provided with a conducting tube as in 6. Heat gently, and after allowing the air in the flask to be expelled, begin to collect the gas which is produced. As this gas is soluble to a considerable extent in cold water, it is better to fill the pneumatic trough with warm water, in which the gas is much less soluble. Collect four jars of the gas, and then withdraw the conducting tube from the trough. Do not heat till all the ammonium nitrate is decomposed, as towards the close of the evolution of gas the decomposition sometimes becomes complex, and other gases are generated so rapidly that an explosion may occur. 29. Properties of nitrous oxide. Combustible substances burn in nitrogen monoxide almost as brightly as in oxygen. They decompose the gas, uniting with its oxygen to form oxides, and leaving unaltered nitrogen. Burn a taper, phosphorus, and sulphur in this gas, in the same manner as they were burned in oxygen. The products of combustion are the same, namely, carbon dioxide, phosphorus pentoxide, and sulphur dioxide ; but there is in addition a residue of nitrogen, 30. Preparation of nitrogen dioxide (nitric oxide). When nitric acid acts upon metals, such as copper or mercury, nitrogen dioxide, water, and a nitrate of the metal used are produced. Thus : Copper and nitric acid yield nitrogen dioxide, copper nitrate, and water : 3 Cu + 8 HN0 3 '= 2 NO + 3 (Cu (NO 3 ) 2 ) + 4 H 2 O. 190-5 +504 =- 60 + s6r$ -f 72. PREPARATION OF NITRIC OXIDE. 21 Place some copper turnings in a flask provided with a funnel and conducting tube, as in the figure. Cover the copper with a layer of water, and add nitric acid by degrees until the gas comes off steadily. Fill two FIG. 8. jars completely, and a third, half full of the gas, and leave the last on the support in the pneumatic trough. 31. Properties of nitrogen dioxide (nitric oxide}. (a} Nitrogen dioxide readily unites with free oxygen, forming higher oxides of nitrogen which have a brown colour, 2 NO + O = N 2 O 3 and NO + O = NO 2 . Remove a jar of the gas from the pneumatic trough and expose it to the air. Observe at once the formation of red fumes, consisting chiefly of nitrogen tetroxide, (NO 2 ). Now dip the mouth of the jar under water, and observe that these fumes disappear owing to their solubility in water, and at the same time observe the rise of water in the jar. (b] Allow a few bubbles of oxygen to pass into the bottle half filled with nitrogen dioxide, and observe a.s 22 PRACTICAL CHEMISTRY. before the formation of red fumes, and the rise oi water in the jar as these dissolve. (c] Nitrogen dioxide does not support the combus- tion of a taper, but phosphorus, when burning very brightly, is not extinguished when placed in the gas. Place a small piece of dry phosphorus in a spoon, light it, and place it in a jar of nitric oxide : it will be extin- guished. Now heat it strongly in the gas flame, and again place it, whilst burning brightly, in the gas, and observe that it continues to burn. 32. Preparation of carbon dioxide (carbonic acid gas\ FIG. 9. When hydrochloric acid acts on calcium carbonate, Carbon dioxide, calcium chloride, and water are pro- duced. Thus : CaC0 3 + 2 HC1 = CO, + CaCU + H 2 O. 100 + 73 = 44 + in + 18. Place some pieces of marble (calcium carbonate) in a flask with funnel and conducting tube, pour some water over it, and then a little hydrochloric acid : PREPARATION OF CARBON DIOXIDE. 23 a rapid effervescence will begin, and the gas will be given off copiously. It may be collected over water, but as it is much heavier than air it is best collected by downward displacement. This is done by placing the conducting tube at the bottom of the gas-jars, and allowing the heavy gas to collect in them and displace the air (Fig. 9). As the gas extinguishes flame, the jars are ascertained to be full when a lighted taper placed in the mouth is at once extinguished. When this is the case, cover the jar and substitute another ; in this way collect four jars of the gas. 33. Properties of carbon dioxide. (a) Carbon dioxide does not support combustion. Place a lighted taper in the gas : observe that it is at once extinguished. (&) Carbon dioxide precipitates lime water. Add some clear lime water to a jar of the gas : it is at once rendered milky owing to the formation of cal- cium carbonate (see 8). Cover the jar closely with the palm of the hand and shake the bottle : the hand will adhere to the bottle owing to the partial vacuum caused by the absorption of the CO 2 by the lime-water. (<:) Carbon dioxide is heavier than air. Pour the carbon dioxide from one of the gas-jars into a jar of air. The gas is so heavy that it will displace the air from the jar, and that this is the case may be shown by placing a lighted taper in each jar : the one originally full of air will now be found filled with carbon dioxide, whilst the other will, if held mouth down- wards for a few minutes, be found to contain only air ; 24 PRACTICAL CHEMISTRY, in the former the taper will be extinguished, in the latter it will continue to burn. 34. Preparation of carbon monoxide (carbonic oxide gas). This gas is prepared pure by the action of sulphuric acid upon formic acid, which is decomposed into car- bon monoxide and water. Thus : CH,0 2 = CO + H 2 0. 46 = 28 -f- 18. It is, however, frequently prepared by the action of sulphuric acid on oxalic acid ; but in this case a mix- ture of equal volumes of carbon dioxide and carbon monoxide is obtained. Oxalic acid yields carbon dioxide, carbon monoxide, and water : 0211.0, = C0 2 -+CO-f-H 2 0. 90 = 44 + 28 + 1 8. Place some crystallized oxalic acid in a small flask provided with a conducting tube (see Fig. 2), cover it with strong sulphuric acid (oil of vitriol), heat gently on a piece of wire gauze, and after allowing the air to escape, collect the gas in the ordinary way at the pneu- matic trough. Fill one gas-jar, and a bottle with a neck narrow enough to allow it to be closed with the thumb. Be careful to remove the conducting tube from the pneumatic trough as soon as the bottles are filled with the gas ; and allow the contents of the flask to become quite cold before pouring the liquid away. 35. Properties of carbon monoxide. (a) The gas burns with a blue flame even when PROPERTIES OF CARBON MONOXIDE. 25 mixed with carbon dioxide, as it is when prepared as above. Apply a light to the wide-mouthed gas-jar containing the mixed gases, and observe the pale blue flame with which the monoxide burns. (b} Pour a small quantity of caustic soda solution into the narrow-necked bottle containing the mixed gases, close it tightly with the thumb, and shake up vigorously without removing the thumb. This will dissolve out the carbon dioxide, which is soluble in caustic soda, hydrogen sodium carbonate being formed, and leave the carbonic oxide. Thus : CO 2 + CO + NaHO = NaHCO 3 -f CO. 44 4- 28 + 40 = 84 + 28. Invert the bottle, mouth downwards, in the pneu- matic trough, taking care that the neck is quite covered with water, and withdraw the thumb : water will rush into the bottle and fill the space previously occupied by the carbon dioxide. Replace the thumb on the neck of the bottle, and shake again to dissolve the last traces of carbon dioxide, and again place the neck of the bottle under water in the trough, and observe that the space occupied by water is half the total capacity of the bottle. Pour a quantity of the gas thus freed from carbon dioxide'into a test tube, add a little lime water, and shake it up : no turbidity ought to be produced. Light the gas at the mouth of the test tube, observe the pale blue flame of the carbonic oxide, and after the gas is burned shake up again, and observe that the lime-water is now rendered turbid. This is because the monoxide in 26 PRACTICAL CHEMISTRY. burning takes up oxygen from the air and produces the dioxide (CO 2 ). 36. Preparation of chlorine * When sulphuric acid acts upon a mixture of common salt and manganese dioxide, sodium sulphate, manga- nese sulphate, water, and chlorine gas are produced. Thus : 2 NaCl + 2 H 2 S0 4 + Mn0 2 = Na 2 SO 4 + MnSO 4 -f 2 H 2 + C1 2 . 117 + 196 + 87 = 142 + 151 -f 36 + 71. Weigh out 30 grams of common salt and the same quantity of manganese dioxide, and mix them together FIG. 10. in a mortar. Then weigh 60 grams of water, place it in an evaporating basin, and add to it cautiously 60 * Experiments with chlorine must be made in a glass closet provided with a draught, as this gas when inhaled produces great irritation, front which serious results may ensue PREPARATION OF CHLORINE. 27 grams of strong sulphuric acid, stirring the liquid with a glass rod until it is thoroughly mixed. The liquid will become very hot, and must be allowed to cool completely; it is then poured into a flask (about one litre capacity), the mixture of salt and manganese dioxide is added, and the contents of the flask shaken. The flask is now provided with a conducting tube bent twice at right angles, heat is applied, and the gas col- lected by displacement like carbon dioxide (Fig. 10). Collect four jars of the gas, and then allow the gas to bubble through some distilled water until no more gas is dissolved. This solution, called chlorine water, has a yellow colour, bleaches vegetable colouring matters, and may be conveniently used instead of the gas in many reactions. 37. Properties of chlorine. (a) i. Chlorine unites readily with hydrogen, form- ing hydrochloric acid (HC1). Place a lighted taper in a jar of chlorine, and observe that it now burns with a smoky flame ; this is caused by the union of the chlorine with the hydrogen of the taper, while the carbon also present in the taper, is liberated, and causes the abundant smoke. 2. Moisten a strip of filter paper with turpentine (com- posed, like the taper, of carbon and hydrogen), and plunge it into a jar of chlorine : hydrochloric acid is at once formed, and the liberated carbon again appears as soot. The action is so energetic that the paper gene- rally takes fire, owing to the heat evolved during the reaction. 28 PRACTICAL CHEMISTRY. (b} Chlorine unites readily with finely-divided metals, forming metallic chlorides. Throw some finely-powdered antimony into a jar of chlorine : union will at once take place with evolution of light, and chloride of antimony will be formed. (c} Phosphorus and chlorine unite together, forming phosphorus chloride. Place a piece of dry phosphorus in a deflagrating spoon, and place it in a jar of chlorine. It will take fire spontaneously and continue to burn with a pale non- luminous flame, and chloride of phosphorus will be formed. (d) Moist chlorine bleaches vegetable colouring matters by uniting with the hydrogen of the water ; the oxygen thus liberated attacks the colouring matter and destroys it. Place a piece of red cloth, partly moistened and partly dry, in a jar of chlorine. The dry part will remain un- altered, but the moist portion will be rapidly bleached. 38. Preparation of hydrochloric acid. When sulphuric acid and sodium chloride react on each other, hydrogen-sodium sulphate (bisulphate of soda) and hydrochloric acid are formed. Sulphuric acid and sodium chloride yield hydro- chloric acid and hydrogen-sodium sulphate : H 2 S0 4 -f- NaCl = HC1 + HNaSO 4 . 98 -f 58-5 = 36-5 -f- 120. Place in a flask (about a litre in capacity), provided with a funnel and conducting tube, about 30 grams qf common salt Fit the end of the conducting tube HYDROCHLORIC ACID. 29 into a small bottle containing a little water, which serves to wash the gas, and the tube which passes out of this bottle conducts the purified gas into 100 c. c.* distilled water, as in the figure. When the apparatus is arranged, add by degrees about 50 grams of strong sulphuric acid. The gas will come off at first with- FlG. out heating ; when all the acid has been added and the evolution of gas slackens, heat the flask gently on a piece of wire-gauze or on a sand bath. Observe that the water in the wash-bottle has first to be saturated, and that then the gas begins to be absorbed by the distilled water. When this last fumes strongly, remove * c. c. = contraction for cubic centimetres 3 o PRACTICAL CHEMISTRY. the wash-bottle and fill a cylinder with the gas by downward displacement. (See Fig. 10.) 39. Properties of hydrochloric acid. (a) Hydrochloric acid is extremely soluble in water. Close the cylinder filled with hydrochloric acid gas with a glass plate, invert it, mouth downwards, in a vessel of water coloured with blue litmus, and withdraw the glass plate. Observe the rapid rise of the water inside the cylinder caused by the solution of the gas, and the change of colour from blue to red, showing the acid nature of the solution. () Hydrochloric acid is a powerful acid. Take some of the aqueous solution of the gas (pre- pared in 38) and add a drop or two of litmus solution to it, so as to colour the solution red. Now add a solution of caustic soda, constantly stirring the mix- ture, and observe that a point is reached when the red colour is changed to a faint blue colour. This point denotes that all the acid has been saturated. Evapo- rate the solution and obtain crystals of sodium chloride (common salt). These three substances, viz. the hydrochloric acid, the caustic soda, and the sodium chloride, are exam- ples of three quite distinct classes of chemical com- pounds, called respectively, acids, bases, and salts. (For a description of these three kinds of substances see Roscoe's " Chemistry," p. 55). (c) Ammonia and hydrochloric acid unite to form ammonium chloride. Hold a glass rod dipped in ammonia solution over BLEACHING ACTION OF CHLORINE. 31 the HC1 solution, and observe the white fumes of NH 4 C1 produced. (See also 25.) 40. Preparation of calcium hypochlorite (or bleaching powder}. Chlorine gas is rapidly absorbed by slaked lime, and calcium hypochlorite is formed. Lime and chlorine yield water and bleaching powder (chloride and hypochlorite of calcium) : 2 Ca(HO) 2 -f 4 Cl = 2 H 2 O + (CaCl 2 -f- Ca(OCl) 2 ). 148 -f 142 = 36 + in + 143. Place a small quantity of slaked lime in a beaker, fill up with water so as to have a milky liquid containing lime in suspension, and then allow chlorine to bubble through until the solution smells strongly of it : calcium hypochlorite will be found in solution. Place a piece of red cloth in some of this solution, and subsequently in a little dilute HC1 : chlorine will be liberated by the action of the acid, and the cloth will be bleached. 41. Preparation of hypo chlorous acid. When diluted nitric acid is added to calcium hypo- chlorite, hypochlorous acid and calcium nitrate are formed. Thus : Ca(OCl) 2 + 2 HNO 3 = 2 HC10 -f CaCNOg),. 143 -f- 126 = 105 -f- 164. Place a solution of the bleaching liquor, as prepared above, in a stoppered retort, the neck of which passes into a small flask kept cool by a stream of water. Add a few drops of dilute nitric acid to the con- tents of the retort, and boil the liquid : the distillate 32 PRACTICAL CHEMISTRY. contains a colourless solution of hypochlorous acid, which will rapidly bleach a piece of red cloth placed in it. Observe also the peculiar smell of hypo'chlorous acid, and remember that a solution of chlorine in water (which likewise bleaches) has a yellow colour. 4-2. Preparation of iodine. When sulphuric acid acts upon a mixture of potas- sium iodide and manganese dioxide ; iodine, potassium sulphate, manganese sulphate, and water are produced. Thus: 2KI+MnO 2 +2H 2 S0 4 = I 2 +K 2 SO 4 + MnSO 4 +2H 2 332-2+ 87 + 196 =254+174-2+ 151 + 36. Place in a retort (provided with a receiver kept cold by a stream of water) a few grams of potassium iodide arid a little manganese dioxide, add water and a little dilute sulphuric acid, and heat gently. Observe the violet vapours filling the retort as soon as the iodine is given off, and the grey deposit of iodine on the neck of the retort and in the receiver. If a large quantity forms in the retort neck, heat it gently, so as to obtain as much as possible in the receiver. 43. Properties of iodine. .(a) Iodine is very sparingly soluble in water, more so in alcohol, and very soluble in a solution of any alkaline iodide. Separate the iodine obtained above, from any liquid which has distilled over, and divide it into three small portions and one large portion. Add to one of the small portions some water, to another some alcohol, and to the third a solution of potassium iodide, and PROPERTIES OF IODINE. 33 observe the colours varying from pale to dark brown, which the solutions possess. (b) Free iodine forms with starch a compound called iodide of starch, which has a dark blue colour. Powder a piece of starch (about the size of a pea) in a mortar, stir it up with about 25 c. c. of cold water, and then heat the mixture (preferably in an evaporating basin) till it boils. A thin, clear solution of starch is thus obtained ; add a portion of it to about a quarter of a litre of water, and then a few drops of one of the solu- tions of iodine, and observe the deep blue colour which the liquid assumes. Heat a little of this blue liquid in a test tube, and observe that the colour disappears ; allow it to cool, and observe the reappearance of the colour. 44. Preparation of sodium iodide. Iodine is dissolved by a solution of caustic soda, and sodium iodide (Nal) and iodate (NaIO 3 ) are produced. 61+6 NaHO = 5 Nal -f- NaIO 3 -f 3 H 2 O. 762 -f- 240 = 750 + 198 + 54. Place the larger portion of iodine prepared as above in an evaporating basin, add water and caustic soda drop by drop until the solution becomes nearly colour- less, and evaporate carefully to dryness, and ignite to convert the sodium iodate into iodide. (NaIO 3 =* Nal -f- O 3 ). A white salt is obtained, which is sodium iodide. (a} Dissolve it in water and add to it about \ litre of water, then add a little of the starch solution, and observe that no blue colour is produced as in the case D 34 PRACTICAL CHEMISTRY. oifree iodine. Now add a drop of chlorine water : this will liberate the iodine and the blue colour will be pro- duced ; add more chlorine water, and the colour will disappear, owing to the formation of a chloride oi iodine which does not colour starch blue. (I)} Add a few drops of the sodium iodide solution to some distilled water, then solution of silver nitrate, and observe the pale yellow precipitate of insoluble silver iodide which is produced. Sodium iodide and silver nitrate yield sodium nitrate and silver iodide : Nal + AgNO 3 = NaNO 3 + Agl. 150 + i?o = 85 + 235. 4-5. Preparation of bromine. When sulphuric acid acts upon a mixture of potas- sium bromide and manganese dioxide, potassium sul- phate, manganese sulphate, bromine, and water are produced. Thus : 2 KBr -f MnO 2 -f- 2 H 2 SO 4 = K 2 SO 4 + MnSO 4 + Br 2 + 2H 2 O. 238-2 + 87 + 196 = 174*2 + 151 + 160 + 36. Proceed exactly as in the preparation of iodine (4-2), but substitute potassium bromide for potassium iodide, and keep the receiver very cold. A dark-coloured heavy liquid will be obtained in the receiver : this is bromine. Observe that it is heavier than the water, which always distils over with it, and that it possesses a powerful irritating smell. Observe also that it dis- solves in water, forming a red-coloured solution, which possesses bleaching properties. PREPARATION OF BROMINE. 35 4-6. Preparation of 'sodium bromide. Bromine is dissolved by a solution of caustic soda and sodium bromide (NaBr), sodium bromate (NaBrOs) and water are produced. Thus : 6 Br + 6 NaHO = 5 NaBr + NaBrO 3 + 3 H 2 O. 480 + 240 = 515 -h 151 -f 54. Proceed as in the preparation of sodium iodide (44-), substituting bromine for iodine. Dissolve the salt obtained in water, and observe that the solution is colourless ; add chlorine water and observe the yellow colour produced by the liberation of bromine. 4-7. Preparation and properties of hydrofluoric acid. When sulphuric acid acts upon calcium fluoride (fluor-spar), calcium sulphate and hydrofluoric acid are produced. Thus : CaF 2 + H 2 SO 4 = CaSO 4 -f- 2 HF. 78+98 = 136 -f 40. This gas cannot be collected in glass vessels as it combines with the silica of glass, forming silicon tetra- fluoride, SiF 4 . This property may be observed by covering a glass plate with a thin coating of bees wax, scratching away the wax at certain points, and then exposing the plate to the action of the gas. Place some powdered fluor-spar in a small lead or platinum dish, pour over it some strong sulphuric acid, and heat gently. Observe the fumes of hydrofluoric acid which come off, an-d then place the waxed glass plate across the dish so as to be exposed to the gas, taking care that the heat applied is not sufficient to melt the wax. Remove the plate after a few minutes, warm it to D 2 36 PRACTICAL CHEMISTRY. soften the wax, which may then be rubbed off, and observe that where the wax was scratched the glass is etched, while the part protected by wax has not been attacked. Avoid breathing the fumes, as the gas is very irri- tating to the lungs, and acts powerfully on the skin, producing painful wounds. 48 Preparation of sulphur dioxide. When sulphuric acid acts upon copper, copper sulphate, sulphur dioxide, and water are produced. Thus : Cu -f 2 H 2 SO 4 = CuSO 4 + SO 2 -f- 2 H 2 O. 63-5 -f- 196 = 159-5 + 64 + 36. Place some copper turnings in a flask of about half a litre in -capacity, provided with a funnel and con- ducting tube bent twice at right angles. Pour enough strong sulphuric acid down the funnel tube to cover the copper, and apply heat, collecting the gas, which is heavier than air, by downward displacement (see Fig. 10). Fill two jars with the gas, and then allow it to pass, first into a test tube containing nitric acid, and then into one containing a solution of potassium chro- inate. Observe in the first case the brown vapours of oxides of nitrogen which are given off, and in the second the change of colour from yellow to green. In both cases the sulphur dioxide has acted as a reducing agent, reducing the nitric acid to a lower stage of oxi- dation, and the potassium chromate to a chromium salt of a green colour ; while in each case the sulphur dioxide is oxidized to sulphuric acid, as may be seen PROPERTIES OF SULPHUR DIOXIDE. 37 by adding barium chloride, which will precipitate white barium sulphate, insoluble in hydrochloric acid. 49. Properties of sulphur dioxide, (a) Sulphur dioxide does not support combustion, and reddens litmus solution. Place a lighted taper in a jar of the gas and observe that it is at once extinguished. Then add a solu- tion of litmus, and observe the bright red colour produced. (b} Sulphur dioxide is extremely soluble in water. Invert a jar of the gas in the pneumatic trough, and observe the rapid rise of water inside the jar, produced by the absorption of the gas. SO. Preparation of sulphuretted hydrogen. When sulphuric acid acts upon ferrous sulphide, ferrous sulphate and sulphuretted hydrogen are pro- duced. Thus : FeS + H 2 SO 4 = FeSO 4 + H 2 S. 88 +98 = 152 + 34. Place a few pieces of ferrous sulphide in a small flask provided with a funnel and conducting tube, the latter of which passes air-tight into a small flask con- taining water (to collect any impurities which may pass over), and having a second tube fitted into it and bent so as to pass into a flask containing water (see figure). Now cover the ferrous sulphide with a layer of water, and add a few drops of strong sul- phuric acid: observe the effervescence which soon begins, and the disagreeable and characteristic smel* which the water in the flask soon possesses from 38 PRACTICAL CHEMISTRY. the solution of the gas in it. When the water smells strongly, remove the flask ;# then decant off the acid rrom the ferrous sulphide, wash it two or three times with water, retaining it in the flask so that it may be used again by simply adding fresh acid. FIG. 12, 51. Properties of sulphur retted hydrogen. Sulphurretted hydrogen precipitates the solutions of salts of certain metals in an acid solution, others in an alkaline solution, and does not, under any circum- stances, precipitate the remainder ; thus, solutions of copper salts are precipitated in an acid solution, solu- tions of iron salts in an alkaline one, and solutions of sodium salts are not precipitated at all. Thus : Copper sulphate and sulphuretted hydrogen yield copper sulphide and sulphuric acid : CuSO 4 -f- H 2 S = CuS + H 2 SO 4 . 159*5 + 34 = 95'5 + 98- * The solution decomposes on standing. It is best preserved in a corked bottle, kept inverted in a vessel of water. PREPARA TION OF CA USTIC SODA. 39 Iron sulphate and potash and sulphuretted hydrogen yield iron sulphide, potassium sulphate, and water. FeSO 4 -f 2 KHO -f H 2 S = FeS + K 2 SO 4 + 2 H 2 O. 152 + H2'2 + 34 = 88 -f- 174*2 + 36. Place in a test glass a solution of copper sulphate, in another a solution, of iron sulphate, and in a third a solution of sodium chloride ; to each add a few drops of hydrochloric acid, and then a little sulphuretted hydrogen water. Observe the black precipitate of copper sulphide in the first glass, and no precipitate in the other two glasses. To each of these add a little potash solution, and observe the black precipitate oi ferrous sulphide in the one case, and the absence of a precipitate in the other. 52. Preparation of sodium hydrate (caustic soda). When caustic lime is added to a solution of sodium bicarbonate, caustic soda and calcium carbonate are produced. Thus : CaO -f- NaHCO 3 = NaHO -f CaCO 3 . 56 + 84 = 40 -f loa Dissolve about 40 grams of sodium bicarbonate in about half a litre of hot water. Then weigh out 30 grams of quick lime, slake it with water, and when it is tho- roughly slaked, stir it up with more water so as to obtain a milky fluid having lime in suspension. Add this to the hot solution of sodium bicarbonate, and boil for a few minutes. Withdraw the lamp, allow the precipitate to subside, and observe if a small por- tion of the clear liquid effervesce when hydrochloric acid is added to it. If so, there is still some sodium 40 PRACTICAL CHEMISTRY. bicarbonate unconverted into caustic soda, and more lime must, therefore, be added. If, on the other hand, no effervescence occurs, the decomposition is complete, and the clear liquid is then evaporated in a clean iron or silver dish to dryness. The resulting white sub- stance is sodium hydrate (NaHO) or caustic soda. 53. Properties of sodium hydrate. Sodium hydrate is a powerful alkali, and turns red litmus solution blue. When hydrochloric acid is added to it, it is neutralized, and sodium chloride (common salt) formed. (See 39.) Sodium hydrate and hydrochloric acid yield sodium chloride and water : NaHO + HC1 = Nad + H 2 O. 40 + 36-5 = 58-5 + 1 8. Dissolve some of the caustic soda obtained in 52, in water, and add to it a solution of reddened litmus, and observe the change in colour from red to blue. Dis- solve a second portion in water, and add to it hydro- chloric acid by degrees, until a drop of the liquid taken out on a glass rod ceases to colour litmus paper blue. On evaporating the liquid thus obtained to a small bulk, sodium chloride will separate out. BLOW-PIPE ANALYSIS. PART II. BLOW-PIPE ANALYSIS.-PRELIMINARY EXAMINATION. 54. Blow-pi^e reactions. How to use the blow-pipe.* " Close the holes at the foot of the Bunsen lamp (Fig. 13) so as to exclude air, and thus obtain a luminous flame. Now place the nozzle of the blow-pipe in the centre of the flame, and blow gently through the tube : observe that the flame produced is blue and corre- sponds to the non-luminous flame of the Bunsen lamp. This is the oxidising or outer flame of the blow-pipe. Now hold the nozzle of the blow-pipe jtist outside the luminous gas flame, and blow gently : observe that the flame is partly yellow. This is the reducing or inner flame of the blow-pipe (Fig. 14). The oxidizing flame is used when a substance has to be oxidized, the reducing flame when a body has to be reduced, e.g. from a salt to the metallic state. * The student should be shown, once for all, the different uses of the blow-pipe, and then be allowed to practise on several different substance* Unless when otherwise expressed the substances used must be dry. PRACTICAL CHEMISTRY. EXAMPLES. I. Reduction. Mix together in a mortar equal small quantities of dry sodium bicarbonate and silver nitrate. Place a portion in a little hollow scooped out of a FIG. 13. FIG. 14. sound piece of charcoal, and heat in the reducing flame of the blow-pipe. Observe the bright metallic bead of silver obtained, dissolve it in nitric acid, and precipitate it as chloride with a few drops of HC1. 2. Oxidation (a) Make a small loop on the end ot a piece of platinum wire, heat it, and dip it while hot in some sodium bicarbonate, so as to cause a small quantity to adhere to the wire ; now heat it with the blow-pipe flame until it is fused.* Then place on it a minute quantity of any manganese compound, and heat again in the oxidizing flame of the blow-pipe ; by * In a similar way, beads are also made with bo-ax or microcosmic salt, instead of sodium bicarbonate. BLOW-PIPE REACTIONS. 43 this means sodium manganate is formed, which colours the bead bright green. (b) Heat a small portion of metallic lead on a piece of charcoal in the oxidizing blow-pipe flame. Observe the yellow incrustation on the charcoal produced by the oxidation of the lead to litharge (PbO). 55. Blow-pipe reactions for the commonly occurring metals.* (a] Compounds reduced to metal when heated with NaHCO 3 on charcoal in reducing flame : Silver. Lead. Bismuth. Antimony. Malleable beads. "Brittle beads. (ft) Compounds reduced to metal when heated with a mixture of KCN and NaHCO 3 in reducing flame : Tin. Copper. White. Red, (c] Compounds reduced to metal when heated on charcoal with reducing agents, but which at once volatilize : Mercury. Arsenic. (d] Compounds reduced to magnetic metallic pow- ders when heated on charcoal with reducing agents : Iron. Nickel. Cobalt. (e) Compounds reduced to metal when heated on charcoal with reducing agents, but which are at once converted into oxides : Cadmium. Zinc. Oxide is biown. Oxide is white. * These reactions must be carried out in the order indicated if they are applied to the examination of an unknown compound, since a metallic salt which is reduced by NaHCOa alone, is reduced a fortiori by a mixture of KCN and NaHCO 3 . 44 PRACTICAL CHEMISTRY. (/) Compounds which, after being heated on char- coal, then moistened with solution of CoCl 2 , and heated again, yield characteristic colours : Zinc. Aluminium. Magnesium. Green. Blue. Pink. (g) Compounds which when fused in a borax bead impart to it a characteristic colour : Iron. Cobalt. Nickel. Manganese. Chromium. Copper. Yellow. Blue. Reddish Amethyst. Green. Blue, yellow. (h} Compounds which impart a characteristic colour to any non-luminous flame : Barium. Strontium. Calcium. Potassium. Sodium. Green. Crimson. Red. Violet. Yellow. 56. Having found out approximately by these reac- tions what the substance is, proceed to apply the following Confirmatory Tests. (a*) Silver bead soluble in HNOg. Solution yields with HC1 white curdy precipitate of AgCL Lead bead soluble in HNO 3 . Solution yields with dilute H 2 SO 4 heavy white precipitate of PbSO 4 . Bismuth bead soluble in HNO 3 . Solution evapo- rated with HC1 yields with H 2 S black precipitate of Bi 2 S 3 . Antimony bead soluble in HNO 3 . Solution evapo- rated with HC1 yields with H 2 S orange precipitate of Sb 2 S 3 . (6) Tin bead soluble in HNO 3 . Solution on evapo- ration yields white precipitate of SnOg. * The letters correspond with those in the preceding paragraphs. BLOW-PIPE REACTIONS. 45 Coppei bead soluble in HNO 3 . Solution, on addition of (NH 4 )HO in excess, yields deep blue solution. (c] Mercury compounds, when heated in a small bulb- tube with NaHCO 3 , yield the metal in minute globules. Arsenic compounds, when heated in a small bulb- tube with KCN -f NaHCO 3 yield the metal as a shining mirror. (tl) Iron powder soluble in HNO 3 + HC1 yields yellow solution which gives a deep blue coloration with K 4 Fe(CN) 6 . Cobalt powder soluble in HNO 3 yields red solution, which gives a blue bead when fused with borax. Nickel powder soluble in HNO 3 yields green solu- tion, which gives a reddish yellow bead when fused with borax. (e) Cadmium oxide distinguished by its brown colour on the charcoal. Zinc oxide distinguished by its yellow colour while hot, turning white when cold. (See also under f.) (/) Zinc compounds, a green residue on charcoal when moistened with CoCl 2 and re-heated. (See also under e.} Aluminium" compounds, a blue residue on charcoal when moistened with CoCU tseen , and re-heated. ' when cold ' Magnesium compounds, a pink residue on charcoal when moistened with CoCl 9 and re-heated. (g) Iron borax bead, reddish yellow when hot, pate yellow on cooling in oxidizing flama 4 6 PRACTICAL CHEMISTRY. Iron borax bead, light green in reducing flame. Cobalt borax bead, deep blue in either oxidizing or reducing flame. Nickel borax bead, reddish yellow when hot, paler on cooling, and finally nearly colourless in oxidizing flame. Nickel borax bead, grey when heated in the reducing flame. Manganese borax bead, amethyst-coloured in oxidiz- ing flame, colour disappears in reducing flame. Chromium borax bead, green in either oxidizing or reducing flame. Copper borax bead, blue or greenish-blue in the oxidizing flame, becomes colourless in the reducing flame. (See also under b.} (h) Barium salts colour the non-lumi-> nous gas flame pale green. Strontium salts colour the non-luminous gas flame bright crimson. Calcium salts colour the non-luminous gas flame dull red. Potassium salts colour the non-luminous gas flame violet. Sodium salts colour the non-luminous gas flame yellow. BUNSEN'S FLAME REACTIONS. 57. The flame of an ordinary Bunsen lamp serves for nearly all the reactions which can be performed by the mouth blow-pipe. It is most convenient to use a Best seen after moistening the salt with HCL Br/NSEN*S FLAME REACTIONS. 47 lamp in which the admission of air can be regulated, and in which the flame is kept steady by a conical chimney supported from the tube of the lamp (see Fig. 15). Adjust the brass cap covering the holes d d t FIG. 15. Fig. 13, so as to obtain a small luminous point at 17. Fig. 1 5, and then notice the following zones of flame, and the purposes to which they are best suited. a. Temperature low. Suitable for observing flame colorations of volatile substances. /3. Highest temperature. Suitable for fusions at high temperatures. 48 PRACTICAL CHEMISTRY. y. The lower oxidizing flame. Suitable for oxida- tion of substances in borax or other beads. 8. The lower reducing flame. Suitable for reduc- tions on charcoal, and in fused borax or other beads. e. The upper oxidizing flame (obtained by admitting the maximum of air). Suitable for oxidation at lower temperatures than are found at /3 and y. ;. The upper reducing flame. Suitable for reduc- tions ; possesses greater reducing power than & Metallic films. The more volatile metals, such as arsenic, mercury, and zinc, are reduced from their compounds when these are heated on an asbestos thread in the upper reducing flame (;). If a small porcelain basin, filled with cold water, be held just above the substance to be examined, the volatilized metal condenses on the cold basin as a metallic film. Example. Place a minute quantity of any arsenic compound on a thread of asbestos. Hold this in one hand, and in the other a small porcelain basin filled with cold water. Now place the basin just above the upper reducing flame, and then the asbestos thread im- mediately below. In a few seconds the reduction will be complete ; remove the basin and observe the brown film of metallic arsenic. Moisten the film with cold dilute nitric acid, and observe that it is scarcely soluble ; moisten now with solution of sodium hypochlorite, and observe its instant solubility. Metallic beads. The less volatile metals may be FLAME REACTIONS. 49 obtained as beads, when their compounds are heated with sodium carbonate on a small charcoal rod held in the lower reducing flame (5). Example. Hold a crystal of sodium carbonate in the lamp flame for a few seconds till it begins to fuse ; then rub the fused salt thus obtained over a common wooden lucifer match, and burn this until it is converted into a rod of charcoal. Then allow a single drop of the fused sodium carbonate to fall on the palm of the hand, mix this intimately by means of a pen-knife with a small quantity of silver nitrate, and then place a very small portion of the mixture on the point of the rod of charcoal. Allow the mixture to melt in the lower oxidizing flame (y), and then push the charcoal splinter into the lower reducing flame (). When the reduc- tion is completed, remove the splinter and examine the point with a lens. Minute beads of fused silver will be seen, which may be further examined by break- ing off the end of the splinter, and crushing it, along with a few drops of water in a small agate mortar. Pour off the water (which will carry che charcoal with it), and examine the metal thus obtained in the same way as the bead obtained by the mouth blow-pipe (56, a). 58. Compounds of the following metals form metallic films : ANTIMONY, ARSENIC, BISMUTH, MERCURY, THALLIUM, CADMIUM, ZINC, INDIUM, and LEAD. They may be further distinguished by the following tests : E 5 o PRACTICAL CHEMISTRY. Antimony. Black film, thin part brownj J dilute HNO Black film, thin part brownl With diffi- Grey non-coherent thin film Black film, thin part brown Black film, thin part brown Bismuth. Mercury. Thallium. Cadmium. Zinc. Indium. lead. Compounds of the following metals give no film, but are reduced to metal on charcoal splinter : COPPER, TIN, SILVER, GOLD, PLATINUM, IRON, NICKEL, and COBALT. They may be further distin- guished as follows : HNO 3 . Instantly so- luble in cold dilute HNO 3 . Red bead, soluble in HNO 3 White bead Copper. Tin. Silver. Gold. Yellow bead, insoluble Platinum. Non-magnetic powder Iron. Magnetic powder Nickel. Cobalt. I Fusible to } metallic beads. Not fusible to beads, but ob- (tained as me- tallic powders. 69. Additional tests for the compounds of the following metals : Antimony. On asbestos thread in upper reducing flame, pale green coloration, unaccompanied by smell. Reduced on charcoal splinter, yields white brittle metallic beads. FLAME REACTIONS. 51 Arsenic. On asbestos thread in upper reducing 6ame, pale blue coloration, and characteristic smell (garlic). Reduced on charcoal splinter, yields no metallic bead. Bismuth. Reduced on charcoal splinter, yields shining yellowish brittle splinters of metal. Dissolve in HNO 3 , add SnCJ 2 , and NaHO, yields black preci- pitate of Bi 2 O 2 . Merctiry. Mixed with dry sodium carbonate and potassium nitrate, in a small thin test-tube (5 millimetres wide and 15 millimetres long), and held in the flame (by a platinum wire coiled round it) just below a small porcelain basin filled with water, yields grey film, which, on rubbing with a piece of filter paper, is col- lected into small globules. If the quantity of mercury be large, globules form at once. Thallium. Flame coloration bright green. Re- duced on charcoal splinter, yields white ductile bead, which quickly oxidizes, and is acted on by HQ with difficulty. Cadmium. Reduced on charcoal splinter imper- fectly to a white ductile bead. Zinc. On asbestos thread in upper oxidizing flame, yields white film of ZnO on the porcelain basin. Moisten a square centimetre of filter paper with HNO 3 , and rub it over the surface of the basin so as to dissolve the oxide film ; roll this up and place it in a coil of thin platinum wire. Now burn the paper in the upper oxidizing flame at as low a temperature as possible, and observe that the colour of the ash is E 2 S 2 PRACTICAL CHEMISTRY. yellow, while hot, and white on cooling. Moisten this ash with a drop of a very dilute solution of cobalt, heat in the lamp flame, and observe the green colour pro- duced. Indium. Flame coloration, intense indigo colour. Reduced on charcoal splinter with difficulty to silver- white ductile globules, slowly soluble in HC1. Lead. Reduced on charcoal splinter, yields soft ductile metallic beads, soluble in HNO 3 . Add H 2 SO 4 , yields white precipitate of PbSO 4 . Copper. Reduced on charcoal splinter, yields ductile metallic beads of a red colour. Dissolve in HNO 3 and add K 4 Fe (CN) fl , yields chocolate coloured preci- pitate of Cu 2 Fe (CN) C . Fuse a small quantity of borax on a straight piece of platinum wire (of the thickness of a horsehair), and, having obtained a clear bead, add a trace of any copper compound, and observe the blue bead obtained, which does not alter in the lower reducing flame. Add a trace of tin or any tin salt, and heat in the lower reducing flame ; observe the change of colour to red. owing to the formation of Cu 2 O. Tin. Obtain a borax bead coloured faintly blue by copper, and add the tin compound ; proceed just as described under copper. The change of colour from blue to red indicates presence of tin. Silver. Reduced on charcoal splinter, yields white ductile beads. Dissolve in HNO 3 and add HC1, yields white curdy p recipitate of Ag Cl. Gold-. Reduced on charcoal splinter, yields yellow FLAME REACTIONS. 53 very malleable beads. Dissolve in aqua regia, absorb the yellow solution on filter paper, and moisten with SnCl 2 , yields purple of Cassius. Platinum. Reduced on charcoal splinter yields a grey spongy mass, which becomes lustrous when rubbed in the mortar. Observe insolubility in HC1 and in HNO 3 , and solubility in aqua regia, forming a light yellow solution. Iron. Reduced on charcoal splinter, yields no bead, but minute metallic particles. Crush the end of the splinter in an agate mortar with a little water, and stir gently with a magnetized knife-blade. The finely divided metal will adhere to the knife. Rub this off on filter-paper, dissolve in aqua regia and add K 4 Fe(CN) 6 : observe the blue coloration from forma- tion of Prussian blue. Borax bead : In the oxidizing flame when hot, yellow to brownish red. In the oxidizing flame when cold, yellow to brownish yellow. In the reducing flame, bottle green. Nickel. Reduced on charcoal splinter, yields white lustrous ductile particles, which form a brush on the point of the magnetized knife-blade. Rub these on paper and dissolve in HNO 3 , and observe the green colour produced. Borax bead : In the oxidizing flame, greyish brown or dirty violet. Upper reducing flame, grey from reduced nickel 54 PRACTICAL CHEMISTRY. which often collects to a spongy mass, leaving the bead colourless. Cobalt. Reduced on charcoal 'splinter, yields white lustrous ductile particles, which adhere to the mag- netized knife-blade. Rub these on paper and moisten with HNO 3 3 observe the red colour, which changes to green on addition of HC1 and drying ; moisten with water, and observe the disappearance of the colour. Borax bead : In the oxidizing flame, bright blue, unaltered in the lower reducing flame. 6O. The following metals are most easily re- cognized as compounds : CHROMIUM, MANGANESE, URANIUM. Chromium. Heated on platinum foil with Na 2 CO 3 and with repeated additions of KNO 3 , yields a yellow mass soluble in water to a yellow solution. Manganese. Borax bead amethyst-coloured in the oxidizing flame, colourless bead in the reducing flame. Heated with Na 2 CO 3 and KNO 3 , yields a green bead, soluble in water to a green solution, which turns red on addition of acetic acid. Uranium. Borax bead yellow, in the oxidizing flame, which becomes green in the reducing flame, especially on addition of SnCl 2 . The following tests for phosphorus and sulphur compounds are exceedingly delicate : Phosphorus. Ignite the sample, and then powder finely and place in a small test tube about the thick- ness of a straw, alon^ with a piece of magnesium wire PRELIMINAR Y EX AM IN A TION. 5 5 about five millimetres long, which should be covered by the powder. Now heat, and observe the incan- descence caused by the formation of magnesium phos- phide. Moisten the residue, and observe the highly characteristic smell of phosphuretted hydrogen. Sulphur. Reduced on charcoal splinter with Na 2 CO 3 , yields Na 2 S. Break off the charcoal point, place it on a silver coin, and moisten with a drop of water. The silver is at once blackened, owing to the formation of silver sulphide. 61. PRELIMINARY EXAMINATION OF SINGLE SALTS. Before proceeding to the systematic analysis of single salts or mixtures, it is always advantageous to subject substances to a preliminary examination, in order to ascertain the probable nature of the substance. The tables A and B, which follow, are adapted for the detection of single salts, either soluble or insoluble. 5 r > PRACTICAL CHEMISTRY. TABLE A. PRELIMINARY EXAMINATION OF SINGLE SALTS CONTAIN- ING ONE ACID AND ONE BASE. a. EXAMINATION FOR ACID.* Heat the substance in a dry tube, and observe whether (a) Water is given off. If so, test its reaction with litmus paper. Acid reaction indicates Sulphites (105), t Chlorides (m), &c. Alkaline reaction indicates Ammonium Salts (91). (3) A sublimate forms Yellow (or in red globules) indicates Sulphur (60). | Ammonium Salts (91), White ,, -v Mercury (64 or 67), Anti- 1 xnony (74), Arsenic (75). (c) A gas is evolved Oxygen indicates { C e *f e I9) ' Nitrates ^' Carbon monoxide Oxalates (100). Nitrogen tetroxide ,, Nitrates (23). Ammonia ,, Ammonium Salts (91). Carbon dioxide Carbonates (102). (if) The substance alters in colour. To black, indicates Organic matter, to yellow (while hot), indicates Zinc Oxide or Carbonate, to brown, indicates Cadmium Carbonate. Take a fresh portion of the substance, add HC1, and observe whether (e) A gas is evolved with effervescence. Smelling like burning sulphur indicates { !%**&$&$ ,, ,, rotten eggs ,, Sulphides (117). bitter almonds Cyanides (114). ( Peroxides, Chro- of chlorine (on heating) ,, < mates (no), Hypo- ( chlorite s (115) Rendering lime water turbid Carbonates (102). (f) Try if the substance is soluble in water ; if so, add BaCl 2 solution to a portion of the solution, and observe if a precipitate form. White and insoluble in HC1 indicates Sulphates (96). * Certain substances, not acids (e.g. mercury, sulphur, ammonia), are for convenience included here. t These numbers refer to the paragraphs where the reactions of the acids and bases are to be found. TESTS FOR SINGLE SALTS. 57 White and soluble in HC1 indicates Phosphates (98), Silicates (103), Oxalates (100), Borates (99), and Fluorides (101). Also Carbonates and Sulphites (see e). If BaCl 2 has given no precipitate, add AgNO 3 to another portion of the solution, and observe if a precipitate form. White and insoluble in HNO 3 indicates Chlorides (in), or Cyanides (114) (see e). Yellowish-white and insoluble in HNO 3 indicates Bromides (112) and Iodides (113). Yellow and soluble in HNO 3 indicates Phosphates (98). Black indicates Sulphides (117) (see e). () If neither water nor HC1 has dissolved the substance, try nitric acid ; and if this does not dissolve it, try aqua regia; and if this does not dissolve it, examine the substance according te the methods described in Table B, pag 58. NOTE. Use as little acid as possible to dissolve the substance. If nitric acid or aqua regia has been used, evaporate the solution to dryness with HCI before proceeding to examine for the base. /?. EXAMINATION FOR BASE. A solution having been obtained, observe (It) If HCI produce a precipitate, it indicates Silver (63), Mercurous Salts (64), or Lead (65). (0 If HCI + H 2 S produce a precipitate, it indicates If black, Mercuric Salts (67), Lead (65), Bismuth (69), or Copper (70). If yellow, Cadmium (71), Arsenic (75), or Stannic Salts (73)- If orange, Antimony (74). If brown, Stannous Salts (73). (/) If (NH 4 ) HO + NH 4 C1 + (NH 4 ) 2 S produce a precipitate, it indicates If black, Iron ( 77 ), Nickel (78), or Cobalt (79). If white, Zinc (80), or Aluminium (81). If flesh-coloured, Manganese (82). If green, Chromium (83). (All Chromium compounds are coloured.) (#) If (NH 4 )HO + NH 4 C1 + (NH 4 ) 2 C0 3 produce a precipitate, it indicates /Barium (85) (tinges flame green). The colours are best { Strontium (86) ( ,, crimson), seen after moistening lOr Calcium (87) ( dulired). the salts with HCI. (/) If not precipitated by the above reagents, it indicates ^-Magnesium (89), precipitated by Na 2 HPO 4 + (NH4)HO \ (white). J Potassium fgo) (tinges flame violet). ) Sodium (92) ( yellow). I Or Ammonium Salts (91) heated with NaHO give smeH > of NH 3 . 58 PRACTICAL CHEMISTRY. TABLE B. EXAMINATION OF INSOLUBLE SUBSTANCES. The following substances are, under certain circumstances, insoluble In acids, and must be examined specially : Silica, Silicates, Alu- minia, Aluminates, Oxides of Antimony, Chromium, and Tin, Chrome Iron Ore, Sulphates of Barium, Strontium, and Lead, certain Fluorides (e.g. of Calcium), certain Sulphides (e.g. of Lead), the Chloride, Bromide, and Iodide of Silver, Carbon, and Sulphur. Heat the substance in a dry tube, and observe whether (a) It fuses and volatilizes completely. Sulphur (smells of SO 2 ). (3) It fuses, but does not volatilize. Chloride, Bromide, or Iodide Of Silver. (Yields metallic silver on fusing on charcoal with Na 2 CO 3 ). (f) It is infusible, but disappears on heating. Carbon (Deflagrates when heated with K N0 3 ). ( whilst the sulphides of the remaining metals are insoluble in that reagent. GROUP III. (IRON GROUP.) Group reagent, (NH 4 ) 2 S, in presence of NH 4 C1 and (NH 4 )HO. Iron, Nickel, Cobalt, Zinc, Aluminium, Manganese, and Chromium. The sulphides and hydrated oxides of the metals of this group are insoluble in water, and are therefore precipitated on addition of the group reagent. Alu- minium and chromium are precipitated as hydrated oxides, the others as sulphides. 62 PRACTICAL CHEMISTRY. GROUP IV. (BARIUM GROUP). Group reagent, (NH 4 ) 2 CO 3 , in presence of (NH 4 )HO and NH 4 C1. Barinm ; Strontium, Calcium; The carbonates of the metals of this group are in- soluble in water, and are precipitated on the addition of (NH 4 ) 2 CO 3 ; as, however, they are soluble in acids, (NH 4 )HO must be added when the solution is acid. GROUP V. (POTASSIUM GROUP.) Metals unprecipitated by the above group reagents. Magnesium, Potassium, Sodium, Ammonium. These metals have no common precipitant, and are therefore distinguished by individual tests. The student should at first have several metallic salts given to him, and be asked merely to determine to which of the above groups each salt belongs ; he ought next to make himself familiar with the individual tests for each metal which follow, and then proceed to the separations of the different metals. It will also be well for him to attempt to frame a table of separations for each 'group before consulting those given in the book. METALS OF THE SILVER GROUP. 63 Reactions of the Metals of the Silver Group. f 63. Metals whose chlorides are insoluble in water, and which are precipitated on addition of the group reagent, HC1 : Silver, Mercury, Lead. SILVER. Ag, combining weight 108. 1. HC1 produces a white curdy precipitate of AgCl, insoluble in hot water and in HNO 3 , but readily soluble in (NH 4 )HO. 2. H 2 S, or (NH 4 ) 2 S, produces a black precipitate of Ag 2 S, soluble in boiling HNO 3 , with separation of sulphur. 3. NaHO produces a light brown precipitate of Ag 2 O, insoluble in excess of NaHO, but soluble in (NH 4 )HO. *4. K 2 CrO 4 produces a dark red precipitate of Ag 2 CrO 4 , soluble in hot HNO 3 ; this solution deposits on cooling an acid chromate in needle-shaped crystals. 5. KI produces a pale yellow precipitate of Agl insoluble in HNO 3 . 6. Heated on charcoal with Na 2 CO 3 , in the reducing flame of the blow-pipe, yields bright, malleable metallic beads, soluble in HNO 3 (56, a). 64. MERCURY. Hg, C.w. 200. Mercurous Salts. ; I. HC1 produces a white precipitate of Hg 2 Cl 2 (calomel), insoluble in cold HNO 3 , and blackened by (NH 4 )HO, from formation of Hg 2 Cl(NH 2 ). t The best confirmatory tests are indicated in the following pages by an asterisk. 64 PRACTICAL CHEMISTRY. 2. H 2 S, or (NH 4 ) 2 S, produces a black precipitate of Hg 2 S, not dissolved by boiling HNO^ 3. NaHO produces a black precipitate of Hg 2 O, insoluble in excess of NaHO or (NH 4 )HO. *4. SnCl 2 produces a grey precipitate of Hg. If the fluid be poured off and the residue boiled with HC1, distinct globules are obtained. 5. KI produces a dark green precipitate of Hg 2 I 2 . 6. K 2 CrO 4 produces an orange precipitate of mer- curous chromate. 7. If a drop of neutral or only slightly acid solution of a mercurous salt be placed on a bright piece of copper, metallic mercury is deposited and ihe stair, becomes bright on rubbing : it disappears on heating, owing to the volatility of the mercury. 8. Heated in a small tube with NaHCO 3 , yields grey deposit of Hg, which on rubbing appears in distinct globules (59). 65. LEAD. Pb, c.w. 207. - i. HC1 produces a white precipitate of PbClg, which is converted into a basic salt on adding ammonia, but without change of appearance. PbCl 2 is soluble in a small quantity of hot water, or in a large quantity of cold water. -r 2. H 2 SO 4 produces a heavy white precipitate of PbSO 4 , soluble in NaHO. In dilute solutions this precipitate appears only on standing; if therefore there is no immediate precipitation, the solution should be concentrated by evaporation. PbSO 4 is soluble in boiling HC1, and the solution on cooling deposits needle-shaped crystals of PbCl 2 . SEPARATION OF SILVER GROUP. 65 - 3. K 2 CrO 4 produces a bright yellow precipitate of PbCrO 4 , readily soluble in NaHO, but with difficulty in HNO 3 . * 4. KI produces a bright yellow precipitate of PbI 2 , soluble in boiling water ; the solution on cooling deposits the salt in brilliant golden hexagonal scales. 5. Heated on charcoal with NaHCO 3 , yields malleable beads, and at the same time a yellow incrustation of PbO on the charcoal (56, a). TABLE C. SILVER GROUP (I.). 66. Separation of Silver, Mercury, and Lead. Add HC1 and filter from the precipitated chlorides. PRECIPITATE. FILTRATE. AgCl, Hg 2 Cl 2 , PbOi,. Groups II., III., IV. & V. Wash precipitate twice with cold water, and add washings to filtrate, then twice with hot wateiy,and test part of this for lead with dilute tiaSCV White pre. ' ' ^' " f2P wa?(h rcJ-atedly White precipitate indicates Lead. Boil the remaining part ain the needle-shaped crystals of PbCL,. If lead be found, cipitate free from it with hot water, and treat the residue If the residue is black, this indicates Mercury. Dissolve in HC1 + HNO 3 and test with Sn C1 2 . Confirm by reduction fest . ' A white preci- ' Add HN0 3 . A white preci- pitate indicates the presence ol Silver. CDrrfirm with bluw pipu 66 PRACTICAL CHEMISTRY. Reactions of the Metals of the Copper Group. 67. Metals whose sulphides are insoluble in HC1 and are precipitated in presence of that acid by the gro^ip reagent H 2 S. Mercury j Lead, Bismuth, Copper, Cadmium, Arsenic, Antimony, and Tin. SUB-GROUP A. Sulphides of the above metals in- soluble in (NH 4 ) 2 S a , viz., Mercury, Lead, Copper, Bis- muth, and Cadmium. MERCURY. Hg, C.W. 200. Mercuric Salts. - I. H 2 S produces, when added by degrees, first a white precipitate, which changes to orange, then to brownish red, and finally to a black precipitate of HgS. These successive changes of colour on the ad- dition of H 2 S are exceedingly characteristic. This precipitate is insoluble in HC1 and in HNO 3 , even on boiling ; it is soluble, however, in KHS and in aqua regia. 2. KHO produces a yellow precipitate of HgO, which is insoluble in excess of the precipitant, except when added to very acid solutions. 3. (NH 4 )HO produces in solutions of HgCl 2 a white precipitate of HgCl(NH 2 ) (" white precipi- tate"). METALS OF THE -* 4. SnCl, produces, when added in small quantities, a white precipitate of Hg 2 Cl 2 , but on adding an excess of the reagent, metallic mercury precipitates as a grey powder, and may be united into a coherent globule by boiling with HC1. 5. KI produces a bright red precipitate of HgI 2 soluble in excess either of KI or of HgCl 2 . 6. Reactions 6 and 7 for mercurous salts (64) are also produced with mercuric salts. 68. LEAD. Pb, c.w. 207. 1. H 2 S produces a black precipitate of PbS, even in solutions of PbCl 2 , so that a weak solution of a lead salt which has not been precipitated with HC1 will be precipitated with H 2 S. Hence lead occurs both in the silver and copper groups. 2. Reactions 2, 3, 4, and 5, for lead, in Group I. (65), are also applicable in this group. 69. BISMUTH. Bi, c.w. 210. i. H 2 S produces a black precipitate of Bi 2 S 3 , insolu- ble in KHS and KHO, but soluble in HNO 3 . 2. KHO or (NH 4 )HO produces a white precipitate, which on boiling becomes yellow (Bi 2 O 3 ); the precipi- tate is insoluble in excess of either reagent. * 3. H 2 O, when added in considerable quantity to normal salts of bismuth, produces an immediate white precipitate of a basic salt of bismuth. Bismuth trichloride is most easily precipitated by H 2 O. If another salt of this metal is being examined, it is best to precipitate the oxide first by ammonia ; F 2 68 PRACTICAL CHEMISTRY. dissolve it in as little HC1 as possible, and evaporate it almost to dryness. On adding water to this solu- tion, a precipitate of BiOCl at once forms, which is insoluble in tartaric acid (compare corresponding re- action with antimony, 74, 4). Solutions of bismuth salts containing much free acid do not give this reaction with H 2 O until the excess of acid has been expelled by evaporation. 4. K 2 Cr O 4 produces a yellow precipitate of Bi 2 (CrO 4 ) 3 , soluble in HNO 3 , and insoluble in NaHO. (Compare reaction for lead, 65, 3.) 5. Heated on charcoal with NaHCO 3 in the re- ducing flame of the blow-pipe, yields brittle metallic globules ; also a slight yellow incrustation of oxide on the charcoal. I 7O. COPPER. Cu, c.w. 63-5. " i. H 2 S produces a black precipitate of CuS, soluble in HNO 3 , but insoluble in KHS, and only slightly soluble in (NH 4 ) 2 S 2 . CuS is also dissolved by KCN, but is insoluble in hot dilute H 2 SO 4 . 2. KH O produces a pale blue precipitate of Cu(H O) 2 , insoluble in excess of the precipitant. If the KHO ba added in excess and the mixture boiled, the precipitate becomes black and loses water. * 3. (NH 4 )HO produces, when added in small quanti- ties, a greenish blue precipitate of a basic salt, soluble in excess to a dark blue solution, which consists of a double basic salt of copper and ammonium. 4. K 4 Fe(CN) 6 produces a brown precipitate of TESTS FOR COPPER AND CADMIUM. 69 Cu 2 Fe(CN) 6 , insoluble in dilute acids, but decomposed by KHO. 5. Fe precipitates copper in the metallic state, from its solutions, especially in presence of a little free acid. The iron ought to be bright and clean. 6. Zn also precipitates copper solutions. If a solu- tion of copper containing a few drops of HC1 be placed in a platinum capsule, and a fragment of zinc added, the copper will be precipitated on the platinum as a red coating. 7. Compounds of copper, when heated in the Bunsen lamp flame, impart a green colour to it, especially after addition of AgCl. 8. Mixed with NaHCO 3 + KCN and heated on charcoal before the reducing blow-pipe flame, yields bright red metallic particles, soluble in HNO 3 , and giving a deep blue solution on adding (NH 4 )HO. 71. CADMIUM. Cd, c.w. 112. * i. H 2 S produces a yellow precipitate of CdS, soluble in HNO 3 , but insoluble in KHS, in (NH 4 ) 2 S, and in KCN. CdS is dissolved by hot dilute H 2 SO 4 (Compare reaction for copper, 7O, i.) 2. KHO produces a white precipitate of Cd(HO) 2 > insoluble in an excess of the precipitant. 3. (NH 4 )HO also precipitates Cd(HO) 2 ,but the pre- cipitate is soluble in excess. * 4. Heated on charcoal with NaHCO 3 in the reducing blow-pipe flame, yields no metallic bead, but a brown incrustation of CdO. &1 o 2 2 rt r* m , 9 . l> * 72 PRACTICAL CHEMISTRY. 73. SUB-GROUP B. Sulphides soluble in viz., those of Tin, Antimony, and Arsenic. TIN, Sn, c.w.i 18. Stannous Salts. 1. H 2 S produces a dark brown precipitate of SnS, soluble in (NH 4 ) 2 S 2 (yellow), but nearly insoluble in (NH 4 ) 2 S (colourless). From its solution in (NH 4 ) 2 S 2 it is re-precipitated as SnS 2 (yellow) on adding HC1. SnS is also soluble in KHO, and, on the addition of acids, is re-precipitated as SnS (brown). 2. KHO produces a white precipitate of Sn (HO) 2 soluble in excess of the reagent. 3. (NH^HO also produces i\ precipitate of SiXHO)^ but not soluble in excess. .* 4. HgCl 2 produces at first a white precipitate ol Hg 2 Cl 2 , and on boiling with excess of the reagent, a grey precipitate of Hg. 5. AuCl 3 produces a purple precipitate (purple ot Cassius) on addition of a little HNO 3 . 6. Zn produces a precipitate of metallic tin, in shining laminae or as a spongy mass. 7. Mixed with NaHCO 3 -f KCN and heated on char- coal in the reducing blow-pipe flame, yields small globules of Sn and a white incrustation of SnO 2 . stannic Salts. i. H 2 S produces a yellow preci- pitate of SnS 2 , soluble in (NH 4 ) 2 S, in KHO, and in boiling concentrated HC1. It is with difficulty soluble in (NH 4 )HO, and insoluble in (NH^COg. 2. KHO or (NH 4 )HO produces a -white precipitate of SnO(HO) 2 , soluble in an excess of the precipitant. TESTS FOR TIN AND ANTIMONY. 73 * 3. Zn produces the same reaction as with stannous salts. (See above.) 4. The blow-pipe reaction for Stannic is the same as for Stannous Salts. (See above.) 74. ANTIMONY. Sb, c.w. 120. 1. H 2 S produces in acid solutions an orange pre- cipitate of Sb 2 S 3 , soluble in alkaline sulphides, in KHO, in boiling concentrated HC1, but insoluble in (NH 4 ) 2 C0 3 . 2. KHO produces a precipitate of Sb 2 O 3 , soluble in excess of the reagent. 3. (NH 4 )HO produces a precipitate of Sb 2 O 3 , in- soluble in excess of the reagent. 4. H 2 O produces in solutions of SbCl 3 a white pre- cipitate of SbOCl, soluble in tartaric acid : compare corresponding reaction with bismuth (69, 3). 5. Zn, in presence of HC1 and platinum, precipitates Sb as a black powder, which adheres to the platinum This is best done either by placing a strip of Zn and Pt (in contact) in a solution of Sb made acifl with HC1, or by placing a similar solution in a platinum capsule and dropping in a piece of Zn. The black stain on the platinum is not dissolved by cold HC1, but is immediately dissolved by warm HNO 3 . * 6 (Marsh's Test). If a solution of Sb be placed in , a flask along with Zn and dilute H 2 SO 4 , SbH 3 is given off as a gas, which is decomposed by heat, Sb being deposited. The apparatus is arranged as in the figure, a is the evolution flask ;. , a tube containing 74 PRACTICAL CHEMISTRY. CaCl 2 to absorb moisture ; and c, a tube of hard glass attached to the drying tube, which is drawn out to a point at the extreme end, so as to form a jet. When the dilute acid and zinc have been in contact some time, and when the air has been expelled from the flask, the hydrogen evolved is lighted at the jet and a cold porcelain crucible lid is held against the flame ; if a black stain be produced on it, the materials employed FIG. 16 are not pure, and must be rejected. Having ascer tained that the Zn and H 2 SO 4 are pure, add now by the funnel tube a few drops of a solution of antimony, and observe that the flame now burns with a bluish green colour, and gives off white fumes (Sb 2 O 3 ), and that on placing a cold porcelain lid against the flame a dull black stain of metallic antimony is deposited on it. Obtain several of these stains in order to com- pare them with the corresponding arsenic ones, and then heat the tube c with the lamp flame. Observe TESTS FOR ANTIMONY. 75 the deposition of Sb close to the flame, and the simul- taneous decrease of the green tint of the flame at the jet. Obtain several of these mirrors, observe their silvery lustre, and keep some for comparison with arsenic mirrors. Further reactions. (a) Add to the stain on porcelain a drop of NaCIO : the stain will remain undissolved. () Cut off with a file the portion of tube containing the metallic mirror, and heat it in a dry test tube. The mirror will be oxidized to Sb 2 O 3 , which will deposit as a sublimate on the test tube. Examine this with a tens, and ascertain that it is amorphous. (c) Attach another tube containing an antimony mirror to an apparatus evolving dry H 2 S, and warm the mirror gently (heating first the part of the mirror farthest from the evolution flask) : observe the change in colour from the formation of orange Sb 2 S 3 . Now detach the tube and pass through it (without heating) a current of dry HC1 gas ; the Sb 2 S 3 will be converted into SbCl 3 , which is volatile, and may be collected by dipping the end of the tube under water. On adding H 2 S to this liquid, orange Sb 2 S 3 will be re-precipitated. 7. Allow a current of SbH 3 to pass through a solution of silver nitrate : SbAg 3 will be precipitated (black) and nitric acid left in solution. Filter and dissolve the residue in a hot solution of tartaric acid, add a few drops of HC1, and pass H 2 S through the solution : an orange precipitate of Sb 2 S 3 will be obtained. 8. Heated with NaHCO 3 on charcoal in the reducing blow-pipe flame, yields brittle globules of the metal 76 PRACTICAL CHEMISTRY. and a white incrustation of Sb 2 O 3 on the charcoal. Fumes of the oxide are also given off after the metal has been removed from the flame, and they occasion- ally condense round the bead in a crystalline mass. 75. ARSENIC. As, c.w. 75-2. 1. H 2 S produces in acid solutions a yellow precipi- tate of As 2 S 3 , soluble in alkaline sulphides, in KHO, in HNO 3 , and in (NH 4 ) 2 CO 3 , but nearly insoluble in boiling concentrated HC1. (Compare reactions for Sb r 74, i.) 2. AgNO 3 produces in neutral solutions of the arse- nites a pale yellow precipitate of Ag 3 AsO 3 . This is best obtained by adding AgNO 3 to an aqueous solution of As 2 O 3 , and then drop by drop a very dilute solution of (NH 4 ) HO prepared by adding one or two drops of ordinary (NH 4 ) HO to a test-tube full of H 2 O. The pre- cipitate is readily soluble in excess of (NH 4 )HO, hence the necessity for using a very dilute solution of that reagent. 3. CuSO 4 added under the same conditions as the AgNO 3 , produces a pale green precipitate of CuHAsO 3 (Scheele's green), soluble in (NH 4 )HO. 4. Acetic acid, added to solutions ot As 2 O 3 and then KHO in slight excess, yields (after evaporation to dryness), on ignition in a small tube, oxide of cacodyl 2 (As(CH 3 ) 2 )O, readily recognized by its power- ful and characteristic odour. If SnCl 2 be added to the contents of the tube after ignition, the equally charac- teristic smell of cacodyl chloride, As (CH^Cl, is ob- TESTS FOR ARSENIC. 77 served. These experiments (and also Marsh's test (5) ) must be done with an exceedingly small quantity of sub- stance, owingto the poisonous properties of the products. *5. Proceed exactly as in Marsh's test for Sb (74, 6), substituting a solution of As for one of Sb, and observe the bluish flame with which the mixture of H and AsH 3 burns, and also the production of white fumes of As 2 O 3 . Obtain, as in the case of Sb, stains on porcelain lids, and mirrors by heating the hard glass tube. Compare these with the antimony stains and mirrors, and observe that the latter in the case of As are deposited at a greater distance from the heated part of the tube, owing to the greater volatility of As. Observe also the distinction in colour of the stains : dark brown or almost black in the case of Sb, and, when seen in thin films, pale brown and lustrous in the case of As. Further reactions. (a] Add to one of the stains on porcelain a drop of NaCIO : it will be rapidly dis- solved. (&) Cut off the portion of tube containing a metallic mirror, and heat it in a dry test tube. The mirror of As will be oxidized to As 2 O 3 , which will be deposited in crystals on the cool part of the tube. Examine these with a lens, and observe the octahedral form of the crystals. Take out the piece of tubing which con- tained the mirror, and dissolve the crystals left in the test-tube in warm H 2 O ; add to this solution AgNO 3 and very dilute (NH^HO, and observe the yellow pre- cipitate of AggAsOg. Or, to the aqueous solution of ?a PRACTICAL CHEMISTRY. As. 2 O 3 add a drop of HC1, and pass H 2 S through the solution, and observe the yellow precipitate ol As 2 S 3 . (c) Attach another tube containing an arsenic mirror to an apparatus evolving dry H 2 S, and warm the mirror gently : it will be converted into yellow As 2 S r Now pass dry HC1 through the tube (without warm- ing), and observe that the sulphide remains unaltered. (Compare corresponding Sb reactions, 74, 6, c}. (d) Allow a current of AsH 3 to pass through a solu- tion of silver nitrate : a black precipitate of Ag will be produced, and As 2 O 3 will be found in solution along with HNO 3 liberated from the AgNO 3 . Filter from the Ag, and very cautiously neutralize the free acid with highly diluted (NH 4 )HO, when a yellow precipitate ol Ag 3 AsO 3 will be formed. (Compare corresponding Sb reaction, 74, 7.) 6. Compounds of arsenic, when treated with Zn and strong solution of KHO, are converted into AsH 3 by the action of the nascent hydrogen. If this reaction be performed in a test tube, and the gas escaping be held near a piece of paper moistened with AgNO 3 , a bluish black coloration is produced by the formation of AsAg 3 . (Sb compounds give no similar reaction.) - 7 (Reinsch's Test). Add to the solution of arsenic, HC1 and a few strips of bright copper wire or foil : As is deposited on the copper, which may be removed from the solution, dried by filter paper, and heated in a dry test-tube to obtain the octahedral crystals of ARSENIC REACTIONS. 79 8. Dry reactions. Place the dry arsenic compound at a (Fig. 17), in a drawn-out hard glass tube. Then place above it at b a small rod of well-ignited charcoal, and heat the portion containing the charcoal until it is red hot. This will cause the glass to soften, and the tube will bend so as to bring the portion a into the flame. The arsenic compound will volatilize and be decomposed by the red-hot charcoal, and a metallic mirror will form at c* 9. Place the dry arsenic compound in a bulb tube as at a (Fig. 18), along with a mixture of equal parts of dry Na 3 CO 3 and KCN, and heat the bulb. A mirror FIG. 18. of As will form at 6, which may be further tested by the reactions mentioned for the mirror obtained in Marsh's test (75, 5, b\ If any moisture be deposited on first heating the tube, remove it by inserting a small coil of filter paper. 10. The above reaction is more delicate when the mixture is heated in a current of dry CO 2 . For this * Non-volatile compounds of As must be mixed with dry charcoal powder, and heated in a similar tube, having in addition a small biilb at the lower end to contain the mixture, 8o PRACTICAL CHEMISTRY. purpose, the arsenic compound is pounded in a mortar with a perfectly dry mixture of three parts Na 2 CO 3 with one part of KCN, and placed at a in the tube ab (Fig. 19), through which a slow current of dry CO 2 is FIG. 19. led, and the whole tube gently heated until every trace of moisture is expelled. When this is the case the tube is more strongly heated at a, and the mirror is obtained at b; traces of arsenic escape condensation, and there- fore a slight garlic odour is observed at the extremity of the tube. Antimony compounds treated in this way yield no metallic mirror. ii. Arsenic compounds, when mixed with Na 2 CO 3 and heated on charcoal by the blow-pipe flame, are reduced to metallic arsenic, which at once volatilizes, and may be recognized by the characteristic odour resembling garlic. TABLE E. GROUP II. 76. Separation of Arsenic, Antimony, and Tin (Sab-Group B). Solution in (NH 4 ) 2 S e contains sulphides of As, Sb, and Sn. Add HC1 until acid : the metals arc re-precipitated as sulphides^ Filter, wash precipitate with not water till free 'from HC1 ; 'digest precipitate with one or two pieces of .solid (NH 4 )X0h and H 2 O. Filter. TESTS FOR IRON. 81 RESIDUE. FILTRATE. SnS 2 , Sb 2 S 3 . Wash and dissolve in strong boil- ing HC1, dilute with water, filter, and divide the filtrate into two parts. In one, place a piece of platinum foil and a fragment of zinc touching it. Sb forms a black stain on the platinum. Dissolve by warming with a few drops of H N 0$, dilute with water and add H 2 S. An orange precipitate in- dicates Antimony. Boil the other portion of the filtrate for at least five minutes with some metallic copper. Pour off the liquid and add HgCl White or grey precipitate indicates Tint. Add HC1 until acid; wash pre- cipitated sulphide, and dissolve in HC1 and a little KC1O 3 , boil down to a small bulk, and apply Marsh's test (75, 5). Metallic mrirror, yielding octahedral crystals on heating, indicates Arsenic. Dissolve in H 2 O, and confirm by adding AgNO 3 and dilute (NH 4 ) HO, to obtain yellow precipitate of Ag 3 As0 3 (75, 5 J). G Reactions of the Metals of the Iron Group. 77. Metals whose sulphides and hydrated oxides are insoluble in water, and are precipitated on addition of the group reagent (NHJ 2 S in presence of (NH 4 )HO and NH 4 Q. Iron, Nickel, Cobalt, Zinc, Aluminium, Manganese, and Chromium. IRON. Fe, C.W. 56. Ferrous Salts. \/ i. (NH 4 ) 2 S produces a black precipitate of FeS, insoluble in alkalies, but soluble in HC1. In dilute solutions of ferrous salts (NH 4 ) 2 S produces at first a green colour ; on standing, however, FeS separates as a black precipitate. 2. KHO or (NH 4 )HO produces a white precipitate of ferrous hydrate Fe(HO) 2 , which rapidly acquires a dirty green, and ultimately a reddish brown colour, G 82 PRACTICAL CHEMISTRY. owing to absorption of oxygen and conversion into ferric hydrate, Fe 2 (HO) 6 . 3. Carbonates of the alkalies precipitate FeCO 3 (white), which rapidly darkens in colour owing to absorption of oxygen. "\, 4. K 4 Fe(CN) 6 produces a white precipitate ot K 2 Fe 2 (CN) 6 , which rapidly becomes blue by oxida- tion to Fe 5 (CN) 12 (Prussian blue). It is insoluble in acids, but is decomposed by alkalies. SSs N*5/ K 3 Fe(CN) 6 produces a blue precipitate of Fe 3 Fe 2 (CN) 12 (Turnbull's blue), also insoluble in acids, but decomposed by alkalies. "--^ 6. KCNS produces no coloration. 7. Ba CO 3 produces no precipitate in cold solutions of ferrous salts. 8. Fused with borax in the oxidizing flame, yellowish red beads are produced ; in the reducing flame the beads become green. (See also 59.) Ferric Salts. I. H 2 S in acid solutions produces a precipitate of sulphur, and the salt is reduced to proto- salt, thus : Fe 2 Q 6 -f- H 2 S = 2 FeCl 2 + 2 HC1 + S. 2. (N H 4 ) 2 S produces a black precipitate of FeS mixed with sulphur, insoluble in excess of the reagent and in alkalies, but soluble in HC1 and in HNO 3 . In dilute solutions of iron only a greenish coloration is produced. 3. KHO or (NH 4 )HO produces a brownish red precipitate of Fe 2 (HO) 6 , insoluble in excess of either reagent. *4. K 4 Fe(CN) 6 produces a precipitate of Fe 5 (CN) ]2 (Prussian blue), insoluble in HC1, soluble in C 2 H 2 O 4 , TESTS FOR NICKEL. 83 and decomposed by KHO or NaHO with separation of Fe 2 (HO) G . ^ 5. K 3 Fe(CN) G changes. the colour of the solution to reddish brown, but does not produce a precipitate. 6,-6. KCNS produces even in dilute solutions a blood- red coloration, due to the formation of a soluble iron sul- phocyanide. HC1 does not destroy the coloration, but it is destroyed by C 2 H 3 O 2 Na, HgCl 2 , H 3 PO 4 ,andbyC 4 H O fi . 7. BaCO 3 precipitates ferric solutions completely as Fe 2 (HO) mixed with basic salt. 8. The blow-pipe reactions are the same as for ferrous compounds. 78. NICKEL. Ni, c.w. 587. -J i. (NH 4 ) 2 S produces a black precipitate of NiS, slightly soluble in excess, forming a brown solution from which NiS is precipitated on boiling. The pre- cipitate is very difficultly soluble in HC1, but dissolves in HNO 3 and in aqua regia. ~-^ 2. NaHO or KHO produces a light green precipitate of Ni(HO) 2 insoluble in excess of the reagent, and unalterable in air. 3. (NH 4 ) HO produces also a precipitate of Ni(HO) 2 , readily soluble in excess, yielding a blue fluid, which is re-precipitated by KHO or NaHO. Acid solutions, or those containing salts of ammonia, yield no precipi- tate with (NH 4 )HO. *4. KCN produces a yellowish green precipitate of Ni(CN) 2 , soluble in excess to a brownish yellow solu- tion of 2 KCN, Ni(CN) 2 . This solution is re-precipi- tated by addition of dilute HC1 or HgSO^ and, if boiled G 2 84 PRACTICAL CHEMISTRY. with a strong solution of NaCIO, yields a black preci- pitate of Ni 2 (HO) 6 . 5. KNO 2 , in presence of C 2 H 4 O 2 , produces no preci- pitate. 6. Fused with borax in the oxidizing blow-pipe flame, yields reddish yellow beads when hot, which become paler on cooling. In the reducing flame the bead becomes grey by the presence of metallic nickel. (See also 59.) 79. COBALT. Co, c.w. 587. 1. (NH 4 ) 2 S produces a black precipitate of CoS, insoluble in excess of the reagent and in HC1, but soluble in aqua regia. 2. KHO or NaHO precipitates blue basic salts, which turn green on exposure to air by oxidation. On heating the precipitate is converted into red hydrate Co(HO) 2 , which is soluble in(NH 4 ) 2 CO 3 to a reddish violet solution. 3. (NH 4 )HO produces the same precipitate as KHO, soluble in excess, yielding a reddish brown fluid, which is re-precipitated by KHO or NaHO. Acid solutions, or those containing salts of ammonia, are not preci- pitated. 4. KCN produces a light brown precipitate of Co (CN) 2 , soluble in excess of the reagent by formation of 2 KCN, Co(CN) 2 . This solution is re-precipitated by the addition of HC1 or H 2 SO 4 . (If the cobalt solution to which KCN is added be acid, a precipitate is pro- duced, soluble in excess of the reagent. When this solutionls boiled, potassium cobalti-cyanide K 3 Co(CN) 6 is formed which is not re-precipitated by HC1 or H 2 SO 4 , nor by N'dClO.) \ TESTS FOR ZINC. 85 5. KNO 2 , added to cobalt solutions with addition of acetic acid, produces on standing in a warm place a yellow crystalline precipitate (double nitrite of potas- sium and cobalt). * 6. Fused with borax in either blow-pipe flame, yields deep blue beads, which are almost black if the quantity of Co be large. (See also 59.) so. ZINC. Zn, c.w. 65-2. 4 i. (NH 4 ) 2 S produces a white precipitate of ZnS, insoluble in. excess of the ~eagent and in KHO, but soluble in the mineral acids. cfet*-ir&*** i ' Cff+fy \^ 2. KHO or NaHO produces a white precipitate of Zn(HO) 2 , soluble in excess of either reagent and in (NH 4 )HO. This solution is re-precipitated by diluting and boiling, but is not precipitated by addition of NH 4 C1. 3. Na 2 CO 3 produces a white precipitate of basic carbonate, insoluble in excess of the reagent. 4. (NH 4 ) 2 CO 3 also precipitates the basic carbonate, but it is soluble in excess of the reagent. 5. Heated on charcoal with Na 2 CO 3 in the reducing blow-pipe flame, a yellow incrustation of ZnO is obtained, which becomes white when cold. * 6. Heated on charcoal by the blow- pipe flame after moistening with CoCl 2 solution, an infusible green mass is obtained. (See also 59.) 7 ^ // // 7 KH H HO^j^^~^'J Sls"^ 81. ALUMINIUM. Al, c.w. 27*3. i. (NH 4 ) 2 S produces a white flocculent precipitate ofA! 2 (HO) 6 . 86 PRACTICAL CHEMISTRY. 2, KHO or NaHO produces also a precipitate of A1 2 (HO) 6 , soluble in acids, even in hot acetic acid and in excess of the reagent. This solution is not precipi- tated by H 2 S, but is re-precipitated by NH 4 C1, or by adding (NH 4 )HO after acidifying with HC1. 3. (NH 4 )HO also precipitates A1 2 (HO) 6 , soluble in a very large excess of the reagent, more difficultly soluble in presence of salts of ammonia. 4. BaCO 3 produces a precipitate ot A1 2 (HO) 6 mixed with basic salt. 5. Na 2 HPO 4 precipitates aluminium phosphate, in- soluble in (NH 4 )HO and in NH 4 C1, but soluble in KHO or NaHO, and in acids. It does not, however, dis- solve in hot acetic acid like aluminium hydrate. * 6. Heated on charcoal in the blow-pipe flame, then moistened with CoCl 2 , and re-heated, an infusible blue mass is obtained. 82. MANGANESE. Mn, c.w. 55. x^ i. (NH 4 ) 2 S produces a flesh-coloured precipitate of MnS, soluble in acids, even in acetic acid. ^ 2. KHO or NaHO produces a dirty- white precipi- tate of Mn(HO) 2 , insoluble in excess of the reagent ; the precipitate rapidly darkens in colour by absorption of oxygen. The freshly-precipitated hydrate is dis- solved by NH 4 C1, but the higher oxide is insoluble. 3. (NH 4 )HO produces the same precipitate of Mn(HO) 2 , insoluble in excess of the reagent ; but it gives no precipitate if the manganese solution contain NH 4 C1. Such a solution on standing precipitates the dark brown hydrate. TESTS FOR CHROMIUM. 87 4. Na 2 CO 3 produces a white precipitate of MnCO 3 , which darkens in colour by absorption of oxygen. 5. If any manganese solution (free from chlorine) be treated with PbO 2 and then boiled with HNO 3 , it is converted into permanganate, which is recognized by its pink colour as soon as the mixture has settled. * 6. If any manganese compound be fused on platinum foil with Na 2 CO 3 and a trace of KNO 3 , it is converted into Na 2 MnO 4 , recognized by its bright green colour. 7. Fused with borax in the oxidizing flame, an ame- thyst-coloured bead is obtained, which becomes colour- less in the reducing flame. 83. CHROMIUM. Cr, c.w. 52-1. \ i. (NH 4 ) 2 S produces a bluish green precipitate of Cr 2 (HO) , insoluble in excess of the reagent. 2. (NH^HO also precipitates the hydrate, soluble to some extent in excess, yielding a pink fluid, but on heating the precipitation is complete. 3. KHO or NaHO precipitates also Cr 2 (HO) 6 , soluble however in excess, yielding a green or bluish violet solution. On continued boiling or addition of NH 4 C1 and heating, the hydrate is re-precipitated. 4. BaCO 3 produces a precipitate of Cr 2 (HO) 6 along with basic salt ; the precipitation is not complete till the mixture has stood some time. 5. Fused with Nq^COg and KNO 3 on platinum f yellow Na 2 CrO 4 is obi * 6. Fused_with_^Di^x-Trr^rther flame (but best in tl lame), green beads are obtained. iQ^a^+b^s; HL ' ^^: T^^sk TABLE 84. IRON GROUP*(III.).-Separationof iron, Nickel, To the filtrate from the sulphides of the Cu and As groups add (NH 4 ) HO and shake for some time. Filter (preferably by the Bunscn pump). Wash filtrate is of:en brown in presence of Ni. Treat RESIDUE. NiS and CoS. Test for Co by borax bead. Dis- solve the black residue in H Cl METHOD I. Cr is ABSENT. and K Cl O 3 . Boil down with a little KC1 Os till it smells of Cl. Add Boil down just pure NaHO till strongly alkaline. Filter. to dryness, di- RESIDUE. FILTRATE. lute with H 2 O, add KCN in ex- cess, then a drop Fe 2 (HO) 6 ,M'n (H0) 2 . Wash with hot H 2 O, dissolve in HC1, add (NH 4 )HO, and filter. Al, Znt Divide into two parts. boil for a few RESIDUE. 1 FILTRATE. i. Add H 2 S minutes, add NaCIO t in ex- cess, and boil again. A black precipitate indi- cates Nickel. Fe 2 (H0) 6 . Dissolve in HC1. Test with K 4 Fe(CN) 6 . Blue precipitate indi- Mn. Boil down and ignite, to expel salts of ammo- nium. Fuse or (NH4) 2 S. A white precipitate indicates Zinc. Confirm by flame-reaction. 2. Add HC1 The filtrate from cates Iron. To with NaHO and till acid, then this precipitate may be tested for ascertainwhether the iron is pre- KNO 3 . A green residue indicates (NH 4 )HO till al- kaline. A white Co by evaporat- ing to dryness, sent as ferrous or ferric salt, the Manganese. Traces of-Ni and precipitate indi- cates Alumi- and fusing in a original solution Co are found nium. Confirm borax bead. Blue must be tested along with the by flame re- colour indicates withK 4 Fe(CN) 6) Mn. action. Cobalt. and K 3 Fe(CN) 6 . For the separation of the metals of this group in presence of Phosphoric Acid, &c., see Table L in the Appendix. t NaHO and Bromine water may be used instead of NaCIO. F. Cobalt, Aluminium, Zinc, Manganese, and Chromium. (till alkaline) + (NH 4 C1 + NH 4 ) 2 S. Warm the mixture gently in a small flask well with H 2 O, containing (NH 4 ) 2 S, and finally once with H 2 O alone. Th the precipitate with cold dilute HC1, and filter. Cr, Al, Fe, Zn, and Mn. (Green or violet if Cr be present. Boil down a portion and test for Cr by borax bead. Adopt Method I. if absent, Method II. if present.) METHOD II. Cr is PRESENT. Boil down with a little KC10 3 till it smells of Cl. Add Na 2 CO 3 or NaHO till just neutral or slightly ac d ; allow to become perfectly cold. Add excess ol BaCOj, place in a flask, cork up and shake well, allow to stand till clear. Filter. RESIDUE. | FILTRATE. Fe 2 (H0) 6 , Cr 2 (H0} 6 , Al^HOs (also excess of BaCO 3 l Wash well, boil with pure NaHO, and filter; add HC1 to the filtrate, and then (NH 4 ) HO till alkaline. A white precipitate indicates Aluminium. Confirm by flame-reaction. Fuse the residue insoluble in NaHO with a mixture of Na 2 CO3and KNO 3 , extract with water, and filter. RESIDUE. FILTRATE. Zn, Mn. Precip.tate the Ba present with H 2 SO4 ' n t ne hot solution. Boil well, and filter ; add NaHO. Pre- cipitate indicates Manganese. Confirm by fus- ing with Na 2 COt> and KNO 3 m platinum foil. To the filtrate from the Mn (HO) 2 add(NH 4 ) 2 S. A white precipitate indicates Zinc. Confirm by flame-reaction. Fe 2 (HO) 6 . Dissolve in H 'Cl, and test with K 4 Fe(CN) 6 . A blue precipitate indicates Iron. "1 Cr. Yellow in colour. Acidify with acetic acid ; add lead acetate. A bright vellow precipitate indicates Chro- mium. ANOTHER METHOD* Boil filtrate from Cu group to expel H 2 S, then boil with KC1O 3 , add N H 4 C1, and then(NH 4 )HO till alkaline. Filter. RESIDUE. | FILTRATE. Fe, Al, Cr. Test washed ppt. for Cr by borax bead, and for Fe with K 4 Fe(CN) 6 . Treat residue with warm NaHO, and test for Al in filtrate as in Method I. 2. Co, Ni, Mn, Zn, Add (NH 4 ) 2 S and filter. Separate Co and Ni with HC1 as above. Boil filtrate from NiS and CoS, add NaHO till alkal.ne, and separate Mn and Zn as in Method II. * When this method is used the filtrate cannot be afterwards tested tor potassium. 90 PRACTICAL CHEMISTRY, Reactions of the Metals of the Barium Group. 85. Metals whose carbonates are insoluble in water, and whose solutions are precipitated on the addition of (NH 4 ) 2 CO 3 : as, however, the carbonates are soluble in acids, the solution, if acid, must be neutralized by addition of (NH 4 )HO. Barium, Strontium, Calcium. BARIUM. Ba, c.w. 137. [ i. (NH 4 ) 2 CO 3 produces a white precipitate of BaCO^, soluble in acids, and to a slight extent in NH 4 C1. Qj(^ 2. K 2 CO 3 or Na 2 CO 3 produces also a precipitate of BaCO 3 , insoluble in excess of either reagent. 3. H 2 SO 4 or any soluble sulphate produces, even in dilute solutions, a heavy white precipitate of BaSO 4 , insoluble in acids, alkalies, or salts of ammonium. 4. CaSO 4 or SrSO 4 produces an immediate white precipitate of BaSO 4 . 5. H 2 SiF 6 produces a white precipitate of BaSiF 6 . 6. C 2 (NH 4 ) 2 O 4 produces a white precipitate- of C 2 BaO 4 , soluble in HC1 and in HNO 3 . * 7. K 2 CrO 4 produces a yellow precipitate of BaCrO 4 , insoluble in C 2 H 4 O 2 , but soluble in HC1 and HNO 3 . j 8. Heated in the lamp flame a green coloration is produced, especially on moistening the salt with HC1. 86. STRONTIUM. Sr, c.w. 87-5. i. (NH 4 ) 2 CO 3 or K 2 CO 3 precipitates white SrCOg, soluble in acids, but less soluble in NH 4 C1 than BaCO <3 . STRONTIUM AND CALCIUM. 91 2. H 2 SO 4 produces a white precipitate of SrSO 4 , which is much less insoluble in H 2 O than BaSO 4 , and it therefore precipitates from dilute solutions only on standing or warming. SrSO 4 is slightly soluble in HC1. * 3. CaSO 4 produces, after standing some time, a white precipitate of SrSO 4 . * 4. H 2 SiF 6 does not precipitate strontium solutions. 5. C 2 (NH 4 ) 2 O 4 produces a white precipitate of C 2 SrO 4 , soluble in HC1 and in HNO 3 , also to a slight extent in NH 4 C1, but very sparingly in C 2 H 4 O 2 . ^ 6. K 2 CrO 4 produces, only in concentrated solutions, a yellow precipitate of SrCrO 4 , soluble in C 2 H 4 O 2 . \ ]/T- Heated in the lamp flame a crimson coloration is produced, especially on moistening the salt with HC1. .^67. CALCIUM. Ca, c.w. 40. jr i. (N H 4 ) 2 CO 3 or K 2 CO 3 produces a white precipitate of CaCO 3 , which becomes crystalline on heating. i/*' 2. H 2 SO 4 precipitates from strong solutions of cal- cium salts CaSO 4 , as a white precipitate, which dis- solves in a large excess of water, and also in acids. 3. CaSO 4 produces no precipitate. 4. H 2 SiF 6 produces no precipitate. * 5. C 2 (NH 4 ) 2 O 4 produces, even in dilute solutions, a white precipitate of C 2 CaO 4 , soluble in HC1 or HNO 3 , but insoluble in C 2 H 2 O 4 or in C 2 H 4 O 2 . V 6. Heated in the lamp flame, a dull red coloration is produced, especially on moistening the salt with HC1. This reaction is imperceptible in presence of Ba or Si salts. - J < > TABLE 88. CALCIUM GROUP (IV.). Separation of Barium, Heat filtrate from iron group, add to the hot solution NH^Cl and and add to a portion CaSO4 solution. An immediate precipitate indicates dilute Barium solution. (Test another portion with SrSC>4 for Ba ) To and Sr. Filter. Neutralize filtrate with (NH^HO, and add C^N H^O* I. Ba PRESENT. Ca AHSENT. Dissolve the carbonate in HC1, and evaporate to dryness. Treat the residue with strong alcohol. Filter. RESIDUE. i FILTRATE. Confiini by flame test. Confirm by lighting Green coloration indi- the alcoholic solution. cates BaJ'iom. Crimson coloration in- dicates Strontium. TO TEST FOR II. Ba ABSENT. Dissolve as before, and Filter, and RESIDUE. SrS0 4 . If small, burn the fil- ter in the reducing gas flame to convert SrSC>4 into SrS : moisten with HC1, and test in the lamp-flame. Crimson coloration indicates Strontium. ANOTHER Dissolve in HNO 3 , ness. Treat with strong RESIDUE. Sr(N0 3 i 2 . Confirm as above. G. Strontium, and Calcium. (NH 4 )oCO 3) and filter. Wash precipitate with hot H 2 O, dissolve in HC1, Barium; a precipitate after some time indicates Strontium, or a another portion of the solution in HC1 add H^SO^ and boil to remove Ba An immediate precipitate indicates Calcium. STRONTIUM. Ca PRESENT. precipitate with H 2 SC>4. wash well. FILTRATE. I Ca. Neutralize the solution with (NH 4 )HO, and test with C 2 (NH 4 ) 2 O 4 . White precipitate indi- cates Calcium. METHOD. and evaporate to dry- alcohol. Filter. FILTRATE. Ca. Confirm as above. III. Ba AND Ca PRESENT. Add H 2 S0 4 to the HCl solution (diluted to prevent precipitation of Calcium), and filter. RESIDUE. j FILTRATE, Ca. Neutralize the solution with (NH 4 )HO, and test with C 2 (NH 4 ) 2 4 White precipitate indi- cates Calcium. BaS0 4 , SrS0 4 . Boil in a beaker with a little water, together with a mix- ture of three parts K 2 SO4 an< l one P art K 2 CO 3 . Filter, and treat residue with HNO 3 . The SrSO, is dissolved, and the BaSO 4 left undissolved. (Traces of Ca may be found with the Sr.) ANOTHER METHOD. Dissolve the carbonates in C 2 H 4 O 2 , and pre- cipitate the Ba with K 2 CrO 4 ., Filter. Precipi- tate the Sr and Ca by (NH 4 ) 2 CO 3 , and proceed as in Method II. (Ba Absent, Ca Present). 94 PRACTICAL CHEMISTRY. Reactions of the Metals of the Potassium Group. 89. Metals whose solutions are unprecipitated by the preceding group reagents, but which have no common precipitant, and are therefore recognized by individual tests. Magnesium, Potassium, Ammonium, Sodium. MAGNESIUM. Mg, c.w. 24. Y i. Na 2 HPO 4 produces, in presence of (NH 4 )HO and NH 4 C1, a crystalline white precipitate of MgNH 4 PO 4 . From dilute solutions the precipitation is slow, but may be hastened by stirring with a glass rod. The precipi- tate is soluble in dilute mineral acids and in C 2 H 4 O 2 , but is almost insoluble in dilute solution of (NH 4 )HO. 2. (NH 4 )HO in neutral solutions produces a partial precipitation of the hydrate Mg(HO) 2 , but gives no precipitate in presence of NH 4 C1, in which the hydrate is readily soluble. 3. H 2 S0 4 , H 2 SiF 6 , and C 2 (NH 4 ) 2 O 4 yield no preci- pitates. * 4. Heated on charcoal in the blow-pipe flame, then moistened with CoCl 2 and re-heated, a pink mass is obtained. so. POTASSIUM. K, c.w. 39-1. i. PtCl 4 produces a crystalline yellow precipitate of 2 KC1 + PtCl 4 , except in dilute solutions, which are not precipitated. The precipitation is promoted by stirring, or by addition of alcohol. 2 Tartaric acid precipitates white crystalline hydro- V \ TESTS FOR AMMONIUM. 95 gen potassium tartrate from strong and neutral solu- tions. The precipitation is promoted by stirring. 3. HgSiFg produces a white gelatinous precipitate or K 2 SiF 6 . 4. Heated on platinum wire in the non-luminous flame a violet coloration is produced, which when viewed through a piece of blue glass appears reddish violet. 91. AMMONIUM. NH 4 , c.w. 18. 1. PtCl 4 produces a crystalline yellow precipitate of 2 NH 4 C1 -f- PtCl 4 , except in dilute solutions, which are not precipitated. The precipitate is insoluble in alcohol and ether, and when ignited leaves a residue of spongy platinum. 2. Tartaric acid produces, in strong solutions, a white precipitate of hydrogen ammonium tartrate, similar in appearance to the corresponding potassium salt. 3. Nessler's solution, added to ammonia" or its salts, produces a yellow coloration, or, if the ammonia com- pound be present in large quantity, a brown precipitate. * 4. NaHo or KHO solution when warmed with am- monia salts decomposes them, and NH 3 is evolved, which is recognized by its pungent odour, by its turning red litmus paper blue, and by its forming white fumes with a strong solution of any volatile acid, e.g. HC1. 5. Heated on platinum foil, all compounds of am- monia volatilize completely. 92. SODIUM. Na, c.w. 23, The soda salts are almost without exception soluble in water, so that the flame test alone serves to distin- guish the salts of this metal 9 6 PRA CT1CAL CHEMISTR V. * i. Heated on platinum foil or wire in the non- luminous lamp flame, an intense yellow colour is produced, which, however, is not seen when viewed through blue glass. It is thus possible to distin- guish potassium salts when mixed with sodium salts. TABLE H. GROUP V. 93 Separation of Magnesium, Potassium, Sodium, and Ammonium. The filtrate from the Barium Group is concentrated by evaporation, and a portion ignited on platinum foil. If no residue is left on igni- tion, Mg, K, and Na are absent. Detection of Detection Detection of K and Na. NH 4 . The original of Mg. To a portion (i.)Mg. being absent. ( 2 .)Mg.being present. substance or so- lution is heated with Na HO in a te* tube. Pre-. sence of Am- moniumshown by smell, by the white fumes with HC1, and by its of the concentra- ted cold solution add (NH 4 ) HO and Na 2 HP9 4 . White crystalline precipitate de- notes Magne- sium. Evaporate an- other portion of the solution to dryness, ignite residue, dissolve in a smail quan- tity of water, fil- ter if required, and add to the Evaporate the solution to dry- ness, ignite resi- due, dissolve in water, and add baryta water un- til the solution has an alkaline reaction; boil; fil- litmus paper. clear HquidPtCl 4 , ter. To filtrate, evaporate nearly add (NH 4 ) 2 CO3, To detect Na. Evaporate alcoholic solution (which must have a yellow colour, showing that excess of Pt C1 4 has been added) nearly to dryness, add adda'Sot. Yel- low precipitate indicates Po- tassium. porate to dryness, and test the resi- due for K and Na, As Sub. (I.) a grain or two of sugar, and ignite residue. Exhaust with water, fil- ter, evaporate to dryness ; and if a residue be left, test it by flame re- action for Na. Yellow coloration indicates Sodium. PART IV* REACTIONS OF THE ACIDS. 94. Grouping of the Adds. The acids do not admit of being grouped with the same precision as the bases, but they can be approxi- mately, classified by means of certain group reagents. They are divided into two great classes, Inorganic and Organic Acids. These are readily distinguished 'f-y the action of heat. Salts of Inorganic Acids when heated to redness aie not charred; salts of Organic Acids are at once r.harred, owing to decomposition and consequent sepa- ration of carbon.* >5. Grouping of the Inorganic Acids. GROUP I. (SULPHURIC ACID GROUP). Group reagent, BaCl 2 in presence of HC1. Sulphuric Acid, Hydrofiuo-silicic Acid. The acids of this group are precipitated by Bad,, and the precipitate is not dissolved on addition of HC1, * With the exception of acetic and formic acids. (Sec 128 and 129 H 98 PRACTICAL CHEMISTRY. GROUP II. (PHOSPHORIC ACID GROUP). Group reagent, BaCl 2 . X /X K /"" Phosphoric, Boric, Oxalic, Hydrofluoric, Carbonic, Silicic, Sulphurotui, Hyposulphurous, Arseziions, Arsenic, lodic, and Chromic Acidst The acids of this group are precipitated in neutral solutions by GROUP III. (HYDROCHLORIC ACID GROUP) Group reagent, AgNO 3 . v/* Hydrochloric, Hydrobroxnic, Hydriodic, JHydro- cyanic, Hypochlorous, Nitrous, and Hydro Sulphuric Acids. The acids of this group are precipitated by AgNO 3 , and not by GROUP IV. (NITRIC ACID GROUP). Nitric, Chloric, and Perchloric Acids. * These acids are not precipitated by any reagent, as all their salts are soluble in water. TESTS FOR SULPHURIC ACID. 99 Reactions of the Inorganic Acids belonging to Group I, 96. Acids precipitated by BaCl 2 in presence of HCL Sulphuric Acid, Hydrofluo-siiicic Acid. SULPHURIC ACID. H 2 SO 4 , c.w. 98. i. BaCl 2 produces a white precipitate of BaSO 4 , 'nsoluble in HC1 or HNO 3 . In very dilute solutions .he precipitation is not immediate, but on standing :he solution becomes clouded, and ultimately the pre- .ipitate subsides. *2. Pb(NO 3 ) 2 produces a heavy white precipitate of PbSO 4 , soluble in- NaHO, and in boiling HC1 (on allowing this solution to cool, PbCl 2 crystallizes out). 3. Fused on charcoal with Na 2 CO 3 in the reducing flame of the blow-pipe, a sulphide is produced. If the fused mass be moistened with HC1, the odour of H 2 S is at once perceptible ; or if it be placed on a bright piece of silver and moistened with water, a black stain of Ag 2 S is produced. As the latter reaction is very delicate, care must be taken to use Na 2 CO 3 per- fectly free from Na 2 SO 4 , and it must be borne in mind that other sulphur acids give the same reaction when heated with Na 2 CO 3 . 97. HYDROFLUO-SILICIC ACID. H 2 SiF 6 , c.,7. i^. 1. BaCl 2 produces a crystalline precipitate of BaSiF c , insoluble in HC1. ioo PRACTICAL CHEMISTRY. 2. KC1 produces a gelatinous precipitate of K SiF 6 . 3. Heated with sulphuric acid in a plantium or leaden crucible covered with a watch-glass, the latter is etched owing to the disengagement of HF. Reactions of the Acids belonging to Group II. 98- Acids precipitated by BaCl 2 in neutral solutions. Phosphoric, Boric, Oxalic, Hydrofluoric, Carbonic, Silicic, Sulphurous, Hyposulphurous, Arsenious, Arsenic, lodic, and Chromic Acids. PHOSPHORIC ACID. H 3 PO 4 , c.w. 98. (Ortho-phos- phoric Acid.) 1 i. BaCl 2 produces a white precipitate of BaHPO 4 , readily soluble in HNO 3 or HC1, but with difficulty in NH 4 C1. -\ 2. MgSO 4 along with (NH 4 )HO and NH 4 C1, pro- duces a white crystalline precipitate of Mg(NH 4 )PO 4 -f 6H 2 0, insoluble in (NH 4 )HO, but soluble in HC1,HNO 3 , and acetic acid. In dilute solutions the precipitation does not take place till after the lapse of some time, but is promoted by stirring and gentle warming. 3. AgNO 3 produces a yellow precipitate of Ag 3 PO 4 , soluble in HNO 3 , and also in (NH 4 )HO. 4. Lead acetate produces a white precipitate of Pb 3 (PO 4 ) 2 , soluble in HNO 3 , but almost insoluble in acetic acid. 5. Fe 2 Cl 6 , in presence of excess of sodium acetate, TESTS FOR PHOSPHORIC ACID. 101 produces a yellowish precipitate of FePp^ soluble in HC1, and in excess of Fe 2 Cl G , which eugh't .ftiibrMore to be added drop by drop.* v * , * 6. Ammonium molybdate produce's ^in ,'slolutioijs. made acid with HNO 3 a yellow coloration and ult'i-' mately a precipitate. The reaction is hastened by warming the mixture. The following acids of phosphorus are distinguished from each other and from ortho-phosphoric acid by the following reactions : Pyrophosphoric Acid. H 4 P 2 O r , C.w. 178. 1. AgNO 3 produces a white precipitate of Ag 4 P 2 O 7 , soluble in HNO 3 , and in (NH 4 )HO. 2. Albumen gives no precipitate. Metaphosphoric Acid. HPO 3 , C.w. 80. 1. AgNO 3 produces a white gelatinous precipitate of AgP0 3 . 2. Albumen produces a flocculent white precipitate when added to metaphosphoric acid, and the same precipitate when added to a solution of a metaphos- phate acidified with acetic acid. 3. MgSO 4 -f NH 4 C1 + (NH 4 )HO produces no pre- cipitate. * If the test be applied to an acid solution of a phosphate insoluble in water (e.g. Cas (PO^a in HC1), the free acid is first nearly neutralised with (Nt)4)HO, sodium acetate next added, and then FeoCIg; after this the mixture is boiled. The precipitate, which is of a reddish brown colour, contains all the iron and phosphoric acid : the filtrate contains the base. The phosphoric acid can easily be separated from the iron, and obtained as a soluble ammonium salt by treating the ferric phosphate with (XH 4 ) 2 S. 102 PRACTICAL CHEMISTRY. 99. BORIC AdiD. B(HO) 3 , c.w. 62. i,. IJaQl^'produces a white precipitate of Ba(BO 2 ) 2 , oluMe, in aqds -^ r //^i^L >',. ^gNO^Rixydtices in strong solutions a yellowish white precipitate. In dilute solutions Ag 2 O is pre- cipitated. 3. H 2 SO 4 or HC1, added to hot concentrated solu- tions of alkaline borates, produces on cooling a crystalline precipitate of B(HO)3. *4. If alcohol containing free boric acid be kindled, it burns with a green flame, best seen on stirring the mixture. Borates may be examined in this way by first adding strong H 2 SO 4 , to liberate tne B(HO) 3 . 5. If the solution of a borate be made distinctly acid with HC1, and turmeric paper dipped into it, the latter on gentle warming acquires a brown tint, which is turned blue by cailftic soda. 6. Moistened with H 2 SO 4 and heated in the lamp flame, a green coloration is produced. jr 100. OXALIC ACID. C 2 H 2 O 4 , c.w. 90. 1. BaCl 2 produces in neutral solutions a white precipi- tate of C 2 O 4 Ba, soluble in HNO 3 , in HCl,and in C 2 H 2 O 4 . 2. AgNO 3 produces a white precipitate of C 2 O 4 Ag 2 , soluble in HNO 3 , and in (NH 4 )HO. \ 3. CaCl 2 produces even in dilute solutions a white precipitate of C 2 O 4 Ca, soluble in HNO 3 , and in HC1, but nearly insoluble in (NH 4 )HO, and in acetic acid. On igniting C 2 O 4 Ca, a white residue of CaCO 3 is left, which effervesces on the addition of an acid. *4. Heated with strong H 2 SO 4 , effervescence takes TESTS FOR HYDROFLUORIC ACID. 103 place from the escape of a mixture ^of CO and CO 2 , and the former may be kindled at fh& ;mOitfh s>f the test tube, and will burn with a pa] blue flame. No blackening of the mixture occurs a 1 , in; the \case; of; organic acids, which yield CO on neating with H 2 3O 4 . 101. HYDROFLUORIC ACID. HF, c.w. 20. 1. BaCl 2 produces a white precipitate of BaF 2 , soluble in HC1, and sparingly in NH 4 C1. 2. CaCl 2 produces a gelatinous and almost trans- parent precipitate of CaF 2 , difficult to discern in the fluid, but made more apparent on addition of (NH 4 )HO. The precipitate is very difficultly soluble in HC1, even on boiling, and is nearly insoluble in acetic acid. *3. Heated with H 2 SO 4 , all fluorides are decomposed with evolution of HF, which is recognised by its power of etching glass. A very characteristic oily appear- ance is noticed whenever a fluoride is warmed with H 2 SO 4 in a test tube. (See -47.) The etching is best done by placing the fluoride in a platinum crucible, or small leaden cup, along with strong H 2 SO 4 , and covering the mouth with a waxed watch- glass, convex side downwards, on which a few scratches have been made with a needle. The concave side of the watch-glass is rilled with water to prevent the wax on the other side from melting, and the crucible or cup is then gently heated. On removing the glass and melting off the wax by gentle warming, the^ glass will be found etched at the unprotected parts.* * If the fluoride contain much silica, SiF^ is evolved instead of HF, and is detected by heating the substance with H 2 SO 4 in a test tube, and leading the evolved gas into water. Silica will separate out in flocculerit tufts, and H 3 SiF 6 will be found 'n solut-on. 104 PRACTICAL CHEMISTRY. 4. Heated with- a mixture of borax and HKSO 4 , on r ;'%a loog of pl^Jnym wire in the non-luminous gas flame, ' r BF 3 r is' produced', w.kich momentarily colours the flame ;; fceh/ 102. CARBONIC ACID. H 2 CO 3 or H 2 O -f CO 2 . \ I. BaCl 2 produces in neutral solutions a white preci- pitate qf BaCO 3 , soluble in acids with effervescence. * 2. Treated with dilute HC1, all carbonates at once evolve CO 2 with effervescence, and if this gas be con- ducted into lime-water it produces a turbidity from formation of CaCO^ (The experiment may be conve- niently performed by placing the carbonate and dilute acid in one test tube and the lime-water in another. As soon as the CO 2 has collected, it may be decanted into the lime-water tube care being taken to pre- vent any liquid from being decanted with it and on shaking the latter the lime-water will become turbid.) 103. SILICIC ACID. Si(HO) 4 , c.w. 96. 1. BaCl 2 produces a white precipitate of SiBa 2 O 4 , which is decomposed on addition of HC1, and Si(HO'4 separates out as a gelatinous precipitate. 2. HC1, added drop by drop to a strong solution of a silicate, produces a gelatinous precipitate of Si(HO) 4 , t>ut if added to a dilute solution or in large excess, no precipitate is obtained until the mixture has been eva- porated to dryness and ignited, when SiO 2 separates out, and this is not re-dissolved on addition of HC1. 3. Fused with NajCO, in a loop of platinum wire in TESTS FOR SILICIC ACID. 105 the non-luminous gas flame, effervescence occurs from the disengagement of CO 2 , and the bead is trans- parent on cooling, unless the Na 2 CO 3 be in excess. * 4. Fused with microcosmic salt on a loop of platinum wire in the non-luminous gas flame, solution does not take place, but the silica floats about on the bead undissolved. 1O4. The remaining six acids of this group are preci- pitated or decomposed by one or other of the group reagents for bases, and are therefore precipitated in the course of examination for bases, or expelled on the addition of HC1. The action of the base group- reagents is as follows : d / H 2 SO 3 decomposed by HC1, with evolution of 2 J S0 2 . ^ \ H 2 S 2 3 decomposed by HC1, with evolution of ^ \ SO 2 and separation of S. d / H 3 AsO 3 precipitated by H 2 S as As 2 S 3 (yellow). 5 } H 3 AsO 4 ^ \ HIO 3 decomposed by H 2 S, with formation of \ an iodide and separation of S. o. ( \ H,CrO 4 precipitated by (NH 4 ) 2 S as Cr 2 (HO) 6 . The following are additional tests for these acids : 105. SULPHUROUS ACID. H 2 SO 3 , c.w. 82. i. BaCl 2 produces a white precipitate of BaSO 3 , soluble in HC1. This solution, on addition of chlorine io6 PRACTICAL CHEMISTRY. water, yields a white precipitate of BaSO 4 , the sulphite being oxidized to sulphate. 2. AgNO 3 produces a white precipitate of AgSO 3 , which darkens on heating, from precipitation of Ag. * 3. Added to a mixture of zinc and HC1, H 2 S is pro- duced and recognized by its smell and blackening action on paper moistened with solution of a lead salt. 4. H 2 S decomposes free H 2 SO 3 with separation of sulphur. 106. THIOSULPHURIC (formerly called Hyposul- phurous) ACID. H 2 S 2 O 3 , c.w. 1 14. i. BaCl 2 produces a white precipitate of BaS 2 O 3 , soluble in HC1, with formation of sulphur as a yellow precipitate. *2. HC1, or H 2 SO 4 , produces no immediate precipitate, but on standing a short time sulphur is precipitated (yellow), and simultaneously SO 2 is evolved. 3. AgNO 3 produces a white precipitate of Ag 2 S 2 O 3 , which rapidly darkens in colour and becomes ulti- mately black from formation of Ag 2 S. These changes are hastened by heat. 4. Fe 2 Cl 6 produces a reddish coloration, but on heating it is decolorized, the ferric being reduced to ferrous chloride. 107. ARSENIOUS ACID. H 3 AsO 3 , c.w 126. *i. AgNO 3 produces in neutral solutions a yellow precipitate of Ag 3 AsO 3 , soluble in (NHJHO. If no precipitate appear at first owing to the solution not TESTS FOR IODIC ACID. 107 being neutral, add a few drops of a very dilute solu- tion of (NH 4 )HO until it appears. (See 75, 2.) 2. MgSO 4 + NH 4 C1 + (NH 4 )HO produce no pre- cipitate. (See also reactions for Arsenic, 75.) 108. ARSENIC ACID. H 3 AsO 4 , c.w. 142. * i. AgNO 3 produces in neutral solutions a light brown precipitate of Ag 3 AsO 4 . If necessary, add very dilute ammonia, as in the preceding case. 2. MgSO 4 + NH 4 C1 + (NH 4 )HO produce a white precipitate of MgNH 4 AsO 4 . (See also reactions for Arsenic, 75.) 109. IODIC ACID. HIO 3 , c.w. 176. 1. BaCl 2 produces a white precipitate of Ba(IO 3 ) 2 , soluble in HNO 3 . 2. AgNO 3 produces a white crystalline precipitate of AgIO 3 , readily soluble in (NH 4 )HO, but sparingly soluble in HNO 3 . 3. SO 2 produces at first a precipitate of I, which is converted into HI on addition of excess of the re-agent. * 4. On heating, iodates are decomposed, oxygen being evolved, and in some cases iodine is also given off in violet vapours. no. CHROMIC ACID. H 2 CrO 4 , c.w. 118-2. i. BaCl 2 produces a yellow precipitate of BaCrO 4 , soluble in HC1 and HNO 3 ,but insoluble in acetic acid. io8 PRACTICAL CHEMISTRY. 2. H 2 S in presence of HC1 reduces the solution to Cr 2 Cl 6 (green), with separation of S. In neutral solu- tions, Cr 2 (HQ) 6 is precipitated along with S. 3. SO 2 reduces solutions of chromates to the chromic salt, the colour of which is green. Chromates are likewise reduced by zinc and a dilute acid, by oxalic acid and dilute sulphuric acid, by strong H 2 SO 4 , by strong HC1, and by boiling the solution acidified with HC1 or H 2 SO 4 along with alcohol. 4. AgNO 3 produces a dark red precipitate of Ag 2 CrO 4 , soluble in HNO 3 and in (NHJHO. * 5. Lead acetate produces a bright yellow precipitate of PbCrO 4 , soluble in NaHO, but soluble with diffi- culty in dilute HNO 3 . (Sec also reactions for Chromium, 83.) Reactions of the Acids belonging to Group III. ill. Acids precipitated by AgNO 3 , and not by BaCl 2 . Hydrochloric, Hydrobromic,Hydriodic, Hydrocyanic, Hypochlorous, Nitrous, and Hydro sulphuric Acids. HYDROCHLORIC ACID. HC1, c.w. 36'5. 1 I. AgNO 3 produces a white curdy precipitate of AgCl, which becomes violet on exposure to light, The precipitate is insoluble in HNO 3 , but soluble in (NH 4 )HO, in KCN, in Na^O^ and also to some extent in NaCl. TESTS FOR HYDROBROMIC ACID. 109 * 2. Heated with H 2 SO 4 and MnO 2 , chlorides yield chlorine gas, recognized by its smell, bleaching action, and green colour. 3. Dry chlorides, when heated in a retort with H 2 SO 4 and K 2 Cr 2 O 7 , yield CrO 2 Cl 2 (chromium oxy- chloride), which distils over into the receiver as a dark red liquid, decomposed by addition of water or (NH^HO, yielding a yellow solution, which, on addi- tion of a lead salt, gives a yellow precipitate of PbCrO 4 . 112. HYDROBROMIC ACID. HBr, c.w. 81. i. AgNO 3 produces a pale yellow precipitate of AgBr, insoluble in dilute HNO 3 , soluble in strong (NH 4 )HO, and readily in KCN and Na 2 S 2 O 3 . * 2. Chlorine passed through a solution of a bromide decomposes it with liberation of Br, which dissolves in the liquid and colours it yellow. If this solution be shaken up with ether, the bromine is dissolved by it, and the yellow ethereal solution floats above the liquid which becomes colourless. If the ethereal solution be then separated from the liquid, and NaHO be added, the yellow colour disappears, and NaBr and NaBrO 3 are produced. On evaporation and ignition, oxygen is evolved and NaBr alone remains, which may be tested as in 3. 3. Heated with H 2 SO 4 and MnO 2 , bromides yield red vapours of Br, recognized by its powerful odour. 4. Heated in a retort with K 2 Cr 2 O T and H 2 SO 4 , dry bromides yield dark red vapours, which condense in the receiver to a liquid of the same colour, which no PRACTICAL CHEMISTRY. consists of pure bromine, and is decolorized on adding excess of (NH 4 )HO. (Compare Hydrochloric Acid test, ill, 3.) 113. HYDRIODIC ACID. HI, c.w. 128. 1. AgNO 3 produces a pale yellow precipitate of Agl, insoluble in dilute HNO 3 , and very difficultly soluble in (NH 4 )HO, but readily in KCN and Na 2 S 2 O 3 . 2. Cuprous sulphate * produces a dirty-white pre- cipitate of Cu 2 I 2 , which separates most completely if the solution be made slightly alkaline with Na 2 CO 3 . The reagent produces no precipitate in solutions of chlorides or bromides. 3. KNO 2 produces no reaction in solutions of iodides until a few drops of HC1 or H 2 SO 4 are added, when iodine is at once liberated and colours the solution yellow. If a little starch solution be now added, a deep blue coloration results from the formation of starch iodide. On warming the blue liquid the colour disappears, but reappears on cooling. The produc- tion of blue starch iodide is the most characteristic test for iodine. * 4. Chlorine water (or the gas) liberates iodine from iodides, but excess of Cl causes the formation of IC1 3 , which is colourless, and gives no blue coloration with starch solution. If therefore chlorine water be added drop by drop to a solution of an iodide mixed with starch solution, a blue coloration is produced, * Prepared by dissolving a mixture of two parts CuSO^ and five part; FeSC>4 in water, or by the action of SO a on CuSO4. TESTS FOR HYDROCYANIC A CID. i : i which disappears on further addition of the re- agent. 5. Free iodine (liberated by either of the above methods) is dissolved by CS 2 , forming a violet-coloured solution. If then, a solution of iodine be shaken up with CS 2 , the latter acquires a violet colour. Chloro- form may be substituted for CS 2 . 6. Heated with MnO 2 and dilute H 2 SO 4 , violet vapours of iodine are obtained, which colour paper moistened with starch, blue. 114. HYDROCYANIC ACID. HCN, c.w. 27. 1. AgNO 3 produces a white precipitate of AgCN, insoluble in HNO 3 , with difficulty in (NH 4 )HO, but readily in KCN and Na 2 S 2 O 3 . AgCN is decomposed on ignition, and metallic Ag remains ; this serves to distinguish it from AgCl, which is not decomposed on ignition. 2. If a solution of FeSO 4 , which has become oxi- dized by exposure to the air, be added to the solution of a cyanide made alkaline with NaHO, a bluish green precipitate is formed, which is a mixture of Prussian blue with the hydrated oxides of iron. On adding HC1, these last are dissolved, and the blue precipitate remains. . * 3. HC1 decomposes nearly all cyanides with evo- lution of HCN, recognized by its odour, resembling bitter almonds. If a cyanide be thus decomposed in a small porcelain basin, covered by a similar basin on which a drop of (NH 4 ) 2 S 2 (yellow) adheres, the .112 PRACTICAL CHEMISTRY. latter is converted into (NH 4 )CNS, which gives a blood-red coloration on addition of Fe 2 Cl 6 and HC1. NOTE. Hg(CN) 2 cannot be detected by the above methods. The dry substance is detected by igniting in a small tube, when cyanogen gas is evolved, or the solution is decomposed by H 2 S and filtered from the HgS ; the filtrate contains HCN. 115. HYPOCHLOROUS ACID. HC1O, c.w. 52-5. 1. AgNO 3 produces a white precipitate of AgCl. 2. Pb(NO 3 ) 2 produces a white precipitate, which changes in colour to red, and ultimately to brown froi?? formation of PbO 2 . 3. MnCl 2 produces a dark brown precipitate of MnO(HO) 2 . 4. Indigo and litmus solution are decolorized, espe- cially on addition of an acid. *5- Dilute acids decompose hypochlorites with evolu- tion of Cl. HNO 3 evolves HC1O from hypochlorites. 116. NITROUS ACID. HNO 2 , c.w. 47. 1. AgNO 3 produces a white precipitate of AgNO 2 , soluble in a large excess of water. 2. H 2 S, in presence of acid, produces a precipitate of S, and (NH 4 )NO 3 remains in solution. * 3. FeSO 4 , in presence of an acid, produces a black coloration from solution of NO in the FeSO 4 , (See also 113, test 3.) HYDROSULPHURIC A CID. 1 1 3 117. HYDROSULPHURIC ACID (Sulphuretted Hydro- gen). H 2 S, c.w. 34. J i. AgNO 3 produces a black precipitate of Ag 2 S, insoluble in dilute acids. 2. Lead acetate, even when highly dilute, produces a black precipitate of PbS. 3. Sodium nitro-prusside, in presence of NaHO, pro- duces a reddish violet coloration, even in very dilute solutions. The colour disappears in a short time. x *4. HC1 or H 2 SO 4 decomposes most sulphides with evolution of H 2 S, recognized by its disagreeable odour and by its blackening paper moistened with solution of lead. Reactions of the Acids belonging to Group IV. 118. Acids not precipitated by any reagent. Nitric, Chloric, and Perchloric Acids. NITRIC ACID. HNO 3 , c.w. 63. i. Nitrates when heated evolve oxygen, and in some cases nitrous vapours also. On fusing a nitrate and adding a fragment of charcoal, vivid deflagration occurs. (For the reactions of Nitric Acid, see 22.) A^ and large quantities of B(HO) 3 and HF. To precipitate, add water and then HC1 : if a precipitate remain, H 2 SO 4 was present. 2. To another portion of the neutralized solution add AgNO 3 : n precipitate indicates one or more of the following acids : (*) H Cl, H Br,- H I, H C N, H 4 Fe (C N) 6 , H 3 Fe (C N) 6> H 2 S. (6) H 3 PO 4 , H 3 As O 4 . H 3 A? O 3 , H 2 CrO 4 , Si(HO)4, B(HO) 3 , Ox,T, and CL To the precipitate add cold dilute HNO 3 . Acids under () are in- soluble, those under (b) soluble. ^* Ox, Ci, and T are contractions for oxalic, citric, and tarf aric acids (For the further separation of these organic acids, see Table K.) 1 2 n6 PRACTICAL CHEMISTRY. DETECTION OF ACIDS UNDER ((l). To a portion of the solution add ttarch paste and one drop of a solu- tion of N 2 O 3 in H 2 SO 4 . Blue coloration indicates HI. Add now chlorine water till the blue colour disappears, and shake with chloro- form. If this becomes reddish brown in colour, the presence of HBr is indicated. HC1 is de- tected in the presence of the others by boiling down the solution to dryness and distilling the resi- due with K 2 Cr 2 O 7 and H 2 SO 4 . See also note, page 117. DETECTION OF ACIDS UNDER (3). Test separately for each acid l>y the methods already given, SEPARATION OF H 3 AsO 3 , H 3 AsO 4 , AND H 3 PO 4 . Acidify solution with H Cl, add Na-jSOs, and heat until no smell of SO 2 is given off. Pass H 2 S through the hot solution, filter, and test for H 3 PO 4 with ammonium molyb- date : yellow precipitate indicates H 3 PO 4 . Precipitate another por- tion with magnesia mixture, and test both precipitate and nitrate for arsenic. TEST FOR THE REMAINING ACIDS BY THE FOLLOWING REACTIONS GIVEN UNDER EACH ACID. acid. Confirm Ox by test 4, 1OO, and HF by test 3, 101. For B ( HO) 3 by tests 4 and 5, 99. For Si(HO) 4 , by tests 2 and 4,1O3. ForH 2 CrO 4 , by tests4ands, 11O. For H 2 SO 3 by test 3, 1O5, and smell of SO 2 on adding HC1. For CO 2 by test 2, 1O2. For H 2 S 2 O 3 by tests 2 and 3, 1O6. For H C N by test For H 4 Fe(C N) 6 by tests 3 and 4, 126. For H 3 Fe(C N) 6 , by tests 2 and 3, 127. For H 2 S, by test 4, 117- *For H NOs, by tests b and c, 22. For HC1O3, by tests i and 2, 119. For Ox and HF, by CaCl 2 + acetic (/?) Acids in Insoluble Bodies. Some idea of the nature of the compound is generally obtained by the preliminary examination (Tables A and B). If not dissolved by the * If H I be present, it must first be removed by addition of Fe SO^ -r CuSO 4 . (See 113, 2.) ACIDS IN INSOLUBLE BODIES. 117 ordinary reagents, the substance must be fused with about four times its weight of a mixture of Na 2 CO3 and K 2 CO3. When cold, extract the fused mass with water, and filter if necessary. The filtrate contains the acid, and is neutralized with HC1 or HNOs, and examined by the methods given under (a). The sulphates of barium, strontium, and calcium are decomposed by boiling with a concentrated solution of Na 2 CO3. Filter, and examine the filtrate for the acid. Nitric acid and aqua regia oxidize sulphides to sulphates : hence the solution of a sulphide in these acids always contains H 2 SC>4. In such cases a separate portion of the substance must be examined for H 2 SC>4 by boiling with HC1, diluting with water, and then testing with BaGlg. NOTE. In mixtures of chlorides, bromides, and iodides, or any two of them, proceed as follows. Place a small quantity of the mixture in a test tube, add water and a few pieces of MnO 2 * (free from chlorides^ then one drop only of dilute sulphuric acid and boil: violet vapour indicates Iodides. Add another drop of the dilute acid and boil again. Proceed in this way till no more violet vapour is given off. Then ad J about 2 c.c. dilute sulphuric acid and boil : brown vapour indicates Bromides. Boil till all bromine is expelled, allow to cool com- pletely, add to the residue an equal bulk of strong sulphuric acid and warm : a green gas indicates Chlorides. Confirm by observing if a piece of' moistened red blotting-paper held in the mouth of the tube is bleached. * Powdered MnO 2 produces too much "bumping" to be used for this purpose. nS PRACTICAL CHEMISTRY. 121. GROUPING OF THE ORGANIC ACIDS. GROUP I. (TARTARIC ACID GROUP). Group reagent, CaCl 2 . Tar tar ic, Citric, and Malic Acids (Oxalic Acid, see 100). Acids which are precipitated by CaQ 2 in the cold or on boiling. GROUP II. (BENZOIC ACID GROUP). Group reagent, Benzoic and Succinic Acids. Acids which are not precipitated by CaCl 2 , but which give precipitates with Fe 2 Cl 6 in neutral solutions. GROUP III. Group reagent, AgNO 3 . Ferro -cyanic, Ferri-cyanic, Sulpho-cyanic, Acetic and Formic Acids. Acids precipitated by AgNO 3 in neutral solutions, and not by CaCl 2 , or Fe 2 Cl 6 . Acetates and Formates are only precipitated in concentrated solutions. TESTS FOR TARTAR1C ACID. 119 Reactions of the Organic Acids belonging to Group I. (Tartaric Acid Group.) Acids precipitated by CaCl 2 in the cold or on boiling. Tartaric and Citric Acids. 122. TARTARIC ACID. C 4 H 6 O 6 . i. CaCl 2 in neutral solutions produces a white pre- cipitate of C 4 H 4 CaO c , soluble in acids, and in am- moniacal salts. The precipitate is soluble in KHO, but is re-precipitated when the solution is boiled, and on cooling is re-dissolved. * 2. KC1 produces in solutions containing T in excess a white crystalline precipitate of C 4 H 6 KO 6 , soluble in mineral acids and alkalies, insoluble in acetic acid, The precipitation is promoted by stirring, or by addi- tion of alcohol. 3. Lime-water produces in neutral solutions a white precipitate C 4 H 4 CaO 6 (flocculent at first, afterwards crystalline), soluble in tartaric acid and NH 4 C1, but re-precipitated in crystals from these solutions after standing some time. 4. Add to .some calcium tartrate which has been washed two or three times by decantation (after pour- ing off the wash water as completely as possible), a drop or two of (NHJHO and a crystal of AgNO 3 , and heat the mixture in a test tube. A lustrous mirro* of silver will deposit on the tube. 5. Heated with strong H 2 SO 4 , the mixture darkens wo PRACTICAL CHEMISTRY. rapidly from separation of carbon, and SO 2 , CO, and CO 2 are evolved. 6. Heated to redness, the substance darkens in colour and gives off the characteristic odour of burnt sugar. 123. CITRIC ACID. C 6 H 8 O 7 . i. CaCl 2 produces no precipitate in neutral solutions in the cold, but on boiling, Ca 3 (C 6 H 6 O 7 ) 2 is precipitated, and is not soluble in KHO, but soluble in (NH 4 )HO. * 2. Lime-water produces no precipitate in cold neutral solutions, but on boiling, Ca 3 (C 6 H 5 O 7 ) 2 is precipitated. 3. AgNO 3 produces in neutral solutions a white floc- culent precipitate of C 6 H 5 Ag 3 O 7 , soluble in (NH 4 )HO : this solution does not blacken on boiling. 4. Heated with strong H 2 SO 4 , CO 2 and CO are evolved without any darkening in colour ; on continued heating, however, the mixture darkens, and SO 2 is evolved. 5. Heated to redness, irritating fumes are given off, readily distinguished from those given off by heating the preceding acid. i23a. MALIC ACID. C 4 H 6 O 5 . 1. CaCl 2 produces no precipitate in neutral solutions in the cold, but upon boiling, C 4 H 4 CaO 6 separates from strong solutions. The precipitate when heated with (NH 4 )HO and AgNO 3 causes no separation of silver. 2. Lime-water does not precipitate solutions of malic acid or of malatcs even on boiling. (Compare 123, 2.) TESTS FOR SUCCJNIC ACID. 121 * 3. AgNO 3 produces in neutral solutions a white granular precipitate of C 4 H 4 Ag 2 O 5 , which becomes grey on boiling. 4. Heated with strong H 2 SO 4 , CO 2 and CO are evolved, the fluid then darkens and SO 2 is evolved. Reactions of the Acids belonging to Group II. (Benzonic Acid Group.) 124. Acids precipitated by Fe 2 Cl c , and not by CaCl 2 . Benzole and Succinic Acid. BENZOIC ACID. C 7 H G O 2 . 1. Fe 2 Cl 6 produces, in neutral solutions, a buff- coloured precipitate of ferric benzoate, decomposed by (NH 4 )HO with formation of a more basic benzoate and ammonium benzoate. Ferric benzoate is soluble in HC1 with liberation of benzoic acid. 2. Heated with H 2 SO 4 benzoic acid does not blacken. *3. Heated in an open tube, the acid sublimes in needle-shaped crystals, and an irritating vapour is given off. When kindled, the crystals burn with a smoky flame. 125. SUCCINIC ACID. C 4 H 6 O 4 . i. Fe 2 Cl 6 produces, in neutral solutions, a reddish brown bulky precipitate of ferric succinate, soluble in mineral acids, and decomposed by (NH 4 )HO in a similar manner to ferric benzoate. 122 PRACTICAL CHEMISTRY. 2. Lead acetate produces a white precipitate of lead succinate, soluble in excess of the reagent and in HNO 3 . * 3. BaCl 2j in presence of (NH 4 )HO and alcohol, pro- duces a white precipitate of barium succinate ; this reaction serves to distinguish this acid from bcnzoic, which does not give a similar precipitate. 4. Heated in an open tube, the acid sublimes in fine needles ; and when kindled, the crystals burn with a bluish but not smoky flame. Reactions of the Acids belonging t Group III 126. Acids precipitated by AgNO 3 . in strong neu tral solutions. Ferro-cyanic, Ferri-cyanic, Sulpho-cyaaic, Acetic, and Formic Acids. FERRO-CYANIC ACID. H 4 Fe(CN) 6 . I. AgNO 3 produces a white precipitate of Ag 4 Fe(CN) , insoluble in dilute HNO 3 and in (NH 4 )HO, but soluble in KCN. *2. CuSO 4 produces a reddish brown precipitate of Cu 2 Fe(CN) 6 . 3. Fe 2 Cl 6 produces a deep blue precipitate of Prus- sian blue, insoluble in dilute mineral acids, soluble in oxalic acid, and decomposed by NaHO with separa- tion of Fe 2 (HO) 6 . 4. FeS0 4 gives a light blue precipitate, which rapidly darkens in colour by oxidation. TESTS FOR ACETIC ACID. 123 127. FERRI-CYANIC ACID. H 3 Fe(CN) 6 . i. AgNO 3 produces an orange-coloured precipitate of Ag 3 Fe(CN) 6 , insoluble in dilute HNO 3 , soluble in (NH 4 )HO and KCN. * 2. FeSO 4 produces a blue precipitate of Fe 3 Fe 2 CCN) 13 (Turnbull's blue), insoluble in acids, but decomposed by alkalies. 3. Fe 2 Cl 6 produces no precipitate, but alters the colour to a greenish brown. 127a. SULPHO-CYANIC ACID. H(CN)S. 1. AgNO 3 produces a white curdy precipitate of Ag(CN)S insoluble in dilute acids but soluble in (NH 4 )HO, and in solution of K(CN)S. From the solution in (NH^HO it crystalizes in shining scales. 2. CuSO 4 produces in strong solutions a black crystalline precipitate of Cu(CN) 2 S 2 which changes on standing to the cuprous salt Cu 2 (CN) 2 S 2 which is white. This change takes place at once by the addition of reducing agents (e.g. SO 2 or FeSO^) to the cupric salt. *3. Fe 2 Cl 6 produces a blood-red coloration from formation of a soluble sulpho-cyanate of iron. The colour is destroyed by addition of alkalies, by HgCl 2 and by many acids (e.g. HNO 3 , H 3 PO 4 , H 2 C 2 O 4 , H1O 3 ), but not by HC1 even when concentrated, 128. ACETIC ACID. C 2 H 4 O 2 . I. AgNO 3 produces, in neutral solutions, a white crystalline precipitate of C 2 H 3 AgO2, Soluble in (NH 4 )HO and in hot water. 124 PRACTICAL CHEMISTRY. 2. Fe 2 Clg, in neutral solutions, produces a deep red coloration. On boiling, all the iron separates as a light brown precipitate of basic acetate, and the fluid be- comes colourless. *3- A strong solution heated with H 2 So 4 and alcohol yields acetic ether, recognized by its characteristic and pleasant odour. 129. FORMIC ACID. CH 2 O 2 . 1. AgNO 3 produces, in neutral concentrated solu- tions, a white precipitate of CHAgO 2 , which rapidly darkens, especially on heating, owing to separation of Ag. 2. Fe 2 Cl 6 produces a similar reaction with formates as with acetates. *3. Hg(NO 3 ) 2 produces a white precipitate of the formate of mercury, which, however, soon becomes grey from separation of Hg. *4. Cold strong H 2 SO 4 decomposes formates with effervescence, CO being evolved. On heating, the gas comes off rapidly, and if lighted, burns with a blue flame. 130. Higher fatty .acid. STEARIC ACID. C ]8 H 36 O 2 . 1. Heated with alkalies, a soap is formed. If mutton fat (which contains the acid combined with glycerine) be heated with NaHO sodium stearate is formed, which dissolves in warm water and the glycerine separates. Soaps are decomposed by acids, which unite with the base, and the fat separates and may be dissolved in alcohol. From this solution the fat crystallizes in needles. i30a. CARBOLIC ACID OR PHENOL. C 6 H 5 (HO). i . Bromine water gives a yellowish white precipitate (Tribromo-phenol) even in dilute solutions. 2. Ferric chloride produces a violet coloration. 3. Heated with ammonia and a drop of sodium hypochlorita, a deep blue colour is produced, which becomes red on addition of aqjds. TESTS FOR MORPHINE. 125 i30b. SALICYLIC ACID. C 7 H 6 O 3 . 1. Heated with lime, phenol is evolved, easily recog- nized by its smell (like tar). 2. Ferric chloride produces even in very dilute solutions a deep violet colour, which disappears on addition of acids or alkalies. 3. Bromine water gives a white precipitate. (Com- pare 13Oa I.) TESTS FOR ORGANIC ALKALOIDS AND CERTAIN OTHER ORGANIC BODIES. 131. MORPHINE. C 17 H 19 NO 3 (Opium).* 1. KHO and (NH 4 )HO precipitate morphia from its salts, readily soluble in excess of KHO, less readily in excess of (NH 4 )HO. 2. Concentrated HNO 3 , added to powdered morphia or its salts, produces an orange red coloration, changing afterwards to yellow. The reaction is best seen on a porcelain crucible lid. Addition of SnCl 2 orof Na 2 S 2 O 3 destroys the colour. 3. A neutral solution of Fe 2 Cl 6 produces, when added to morphia (either in the dry state or in solution), a deep blue colour. 4. HIO 3 , when added either to free or combined morphia, is decomposed with liberation of I, which colours the liquid brown. Starched paper added to the solution becomes blue. Addition of (NHJHO deepens the brown colour. 5. H 2 SO 4 produces no coloration with morphia or its salts, but on adding a crystal of K 2 Cr 2 O 7 a bright green colour is produced. 1 132. MECONIC ACID, C 7 H 4 O 7 (Opium). i. Fe. 2 Cl 6 produces a deep red coloration, which does not alter on boiling (like ferric acetate), nor' on treat- ment with HgCl 2 (like iron sulphocyanide). * These names refer to the substances from which the alkaloids, &c. are derived. 126 PRACTICAL CHEMISTRY. 2. Lead acetate produces a white precipitate of lead meconate, insoluble in acetic acid. OPIUM SOLUTIONS : Opium contains both morphia and meconic acid ; hence, in examining solutions of opium, the reactions of both these substances are obtained. The reaction of Fe 2 Cl 6 with meconic acid is highly characteristic, and as meconic acid only occurs in opium, its detection in solutions of opium serves as an indirect indication of the presence of morphia. The two substances are separated by precipitating with lead acetate in pre- sence of acetic acid, and filtering from the lead meconate obtained. The filtrate contains acetate of morphia, If the lead meconate be washed thoroughly and then suspended in water and H 2 S passed through, PbS is precipitated, and meconic acid remains in solu- tion and may be obtained in crystals by evaporation. 133. STRYCHNINE. C 21 H 22 N 2 O 2 (Strychnos mix vomica and St. Ignatius' bean). i. Strong pure H 2 SO 4 dissolves strychnine without any discoloration of the solution, even when heated to 100 C. This solution produces certain characteristic colours, with the following reagents : (a) PbO 2 , blue coloration, changing to violet, then red, and finally yellow. (^) K 2 Cr 2 O 7 , blue coloration, changing soon to yel- lowish red. (<:) K 4 Fe(CN) 6 , violet coloration, changing less quickly. TESTS FOR BRUCINE. 127 \d] MnO ,, violet coloration changing to dark red. 2. Strong HNO 3 dissolves strychnine without dis- coloration unless heated, when the solution becomes yellow. 3. The taste of strychnine is intensely bitter, and is perceptible even in very dilute solutions. 4. Strychnine, when taken in poisonous doses, pro- duces tetanic convulsions. A most characteristic test, founded on this property, consists in injecting a trace of strychnine under the skin of a young frog, which is soon seized with tetanic convulsions, the body becomes rigid and arched, and the animal soon dies. 134. BRUCINE. C 23 H 2C N 2 O 4 4- 4 H 2 O (Strychnos mix vomica and St. Ignatius' bean). i. HNO 3 , when added to brucine, dissolves it, and produces an intense red coloration, which becomes yellow on heating. If SnCl 2 , Na 2 S 2 O 3 or (NH 4 ) 2 S be then added,-the colour changes to violet. (Compare corresponding reaction with Morphia, 131, 2.) 135. QUININE. C 20 H 24 N 2 O 2 (Cinchona Bark). 1. Chlorine water, when added to an acid solution of quinine or its salts, produces no coloration until (NH 4 )HO is added, when a bright green colour is pro- duced. If K 4 Fe(CN) 6 be added before the (NH 4 )HO, a red coloration is produced, soon changing to dirty brown. 2. KHO or (NH 4 )HO produces, in solutions of qui- nine salts, a white amorphous precipitate of quinine, 128 PRACTICAL CHEMISTRY. which, on standing, becomes crystalline. The precipU tate is scarcely soluble in KHO, but slightly soluble in (NH 4 )HO. If a mixture of a quinine solution and (NH 4 )HO be shaken up with ether, the precipitated quinine is dissolved. (Compare test 2 for Cinchonine, 136.) 3. Solutions of quinine or its salts are fluorescent in a high degree, and possess an extremely bitter taste. 136. CINCHONINE. C 20 H 24 N 2 O (Cinchona Bark). 1. Chlorine water produces no coloration, even on addition of (NH 4 )HO, which produces a yellowish precipitate. 2. KHO or (NH 4 )HO produces, in solutions of cin- chonine salts, a white amorphous precipitate of cin- chonine, insoluble in excess, and not dissolved when shaken up with ether. (Compare test 2 for Quinine, 135.) 3. Solutions of cinchonine have a bitter taste and an alkaline reaction. The salts are less soluble in water and alcohol than those of quinine. 137. URIC ACID. C 6 H 4 N 4 O 3 (Urine). 1. HNO 3 dissolves uric acid with effervescence, and the solution when evaporated to dryness and moistened with (NH 4 )HO assumes a bright purple colour (mur- exide). 2. NaaCOg solution dissolves uric acid. If this solu- tion be placed on paper moistened with AgNO s sola- TESTS FOR UREA. 129 tion, a brown stain is produced, owing to the reducing action of the uric acid. 138. UREA. CH 4 N 2 O (Urine). 1. HNO 3 , when added to solutions of urea, unites with it, forming the nitrate, which separates out in crystalline plates, which are tolerably soluble in hot water and alcohol, but nearly insoluble in HNO 3 . 2. KHO added to urea decomposes it on heating, and NH 3 is evolved. 3. Hg(NO3) 2 produces a white precipitate of vari- able composition. 4. HNO 2 decomposes urea at once into CO 2 , H 2 O, and N. CH 4 N 2 O + 2 HNO 2 = CO 2 + 4 N + 3 H 2 O. 5. When chlorine is passed into an aqueous solu- tion of urea the following reaction takes place CH 4 N 2 O + 6 Cl + H 2 = CO 2 -f 2 N -f 6 HC1. 139. CHOLESTERINE. C 26 H 44 O (Biliary Calculi). 1. KHO does not saponify cholestcrine, although in many of its properties it resembles oils and fats. 2. Boiling alcohol dissolves cholesterine, and on cooling it crystallizes out in colourless plates. 140. GRAPE SUGAR. CeH^Oe (Fruit, Honey, Dia- betic Urine, c.). 1. H 2 SO 4 forms with grape sugar a definite com- pound of a yellow colour. No charring occurs as in the case of cane sugar. 2. Solution of grape sugar, when mixed with a few K I 3 o PRACTICAL CHEMISTRY. drops of CuSO 4 and excess of KHO, produces on boiling a red precipitate (Cu 2 O), caused by the re- ducing action of the grape sugar. Solution of cane sugar does not give this reaction until boiled with a single drop of H 2 SO 4 , which converts it into grape sugar. 3. An aqueous solution of grape sugar when mixed with yeast and kept at a temperature of 27 undergoes fermentation, and CO 2 is evolved. The mixture is placed in a test tube provided with a conducting tube which dips into lime water ; over these a bell jar is placed (to prevent absorption of CO 2 from the air) and the apparatus kept for some hours at the necessary temperature. Turbidity in the lime water indicates the presence of grape sugar. (As CO 2 is sometimes given off by the yeast itself, it must be carefully tested before use in the manner described.) 141. ALBUMIN (White of Egg). 1. Boiling water causes albumin to coagulate and it becomes at the same time insoluble, so that if a solu- tion in cold water be heated, coagulation at once takes place. 2. HgCl 2 coagulates albumin, even when present in exceedingly small quantities. Coagulation also takes place with solutions of other salts, e.g. CuSO^ SnCl 2j AgN0 3 . 3. Most acids precipitate albumin from its solutions. HNO 3 acts most readily, and is therefore used to detect the presence of dissolved albumin. Acetic, TESTS FOR CASEIN 131 tartaric, and ortho phosphoric acids do not coagulate albumin, except in very concentrated solutions. CASEIN (Milk and Cheese;. 1. Casein is insoluble in water, but is kept in solu- tion in milk by a small quantity of free alkali. Such a solution does not coagulate by heat, but a film forms on its surface when heated. If this film be removed another forms in its place. 2. All acids (except carbonic) precipitate casein from its solutions ; the precipitate is soluble in excess of the reagent. 3. Infusion of rennet (the inner membrane of the stomach of a calf ) coagulates casein completely. 142. STARCH. C 6 H 10 O 6 (Potato, Wheat, &c.). See 48,4, 143. TANNIC ACID. C^H^O^. (Gall-nuts). 1. Solution of gelatin (isinglass) produces a yellowish flocculent precipitate. A piece of animal membrane gives the same precipitate. 2. Fe 2 Cl 6 produces a dark bluish black precipitate (ink). 144. GALLIC ACID. C 7 H 6 O 6 . (Gall-nuts). 1. Solution of gelatin produces no precipitate. 2. Fe 2 Cl 6 produces in neutral solutions a bluish black precipitate* 132 PRACTICAL CHEMISTRY. 145. Detection of CARBON, HYDROGEN, NITRO- GEN, CHLORINE, SULPHUR AND PHOSPHORUS, in ORGANIC BODIES : (a) CARBON. Mix the substance intimately with powdered CuO, place in a hard glass tube, fill up with small pieces of CuO, and close the tube with a cork, through which a conducting tube passes into a flask containing clear lime-water. On heating, the carbon is oxidized to CO 2 , which renders the lime-water turbid. Organic liquids are examined in a similar way : they are placed in a small bulb tube which is placed inside a hard glass tube and gently heated as soon as the CuO in front is red-hot. In this way the vapour passes over the CuO, and CO 2 is produced. (b} HYDROGEN. Proceed as in a (taking care to use perfectly dry CuO), but connect the conducting tube with a weighed CaCl 2 tube. Heat the hard glass tube as before, and then weigh the CaCl 2 tube. Gain in weight indicates that water has been formed by the oxidation of the hydrogen in the substance. When much water is present, it is visible in the bulb of the CaCLj tube. (c) NITROGEN. Mix the substance intimately with powdered soda-lime, and observe if NH 3 is given off on heating. Ascertain this by the smell and by the action on red litmus paper. (d] CHLORINE. Mix the substance intimately with quick-lime (free from chlorine), place the mixture in a test tube, and heat to redness. Extract the residue with water, filter, and add solution of AgNO 3 to the TESTS FOR SULPHUR. 133 filtrate. By this means organic chlorine compounds, which are not precipitated by solution of AgNO 3 , are decomposed, the chlorine uniting with the calcium to form CaCl 2 . (e) SULPHUR. Solid compounds may be tested by fusing them with a mixture of pure solid KHO and KNO 3 . The fused mass when cold is dissolved in water, acidified with HC1 and tested in the usual way with solution of BaCl 2 . Liquids are gently heated with strong HNO 3 . or with a mixture of HC1 and HC1O 3 , the solution is then diluted and tested with BaCl 2 . (/) PHOSPHORUS. Fuse the substance as in (e) with a mixture of KHO and KNO 3 , or heat it with strong HNO 3 . Add water and test as in 98, 2, 5, or 6. Phosphorus may also be detected by slightly charring the substance and proceeding as in 6O. 14-6. Detection of METALS in PRESENCE OF ORGANIC MATTER : The reactions given for the various metals are not reliable in presence of organic matter, and as metals such as arsenic, lead, mercury, &c., in cases of poison- ing have frequently to be detected when mixed with organic matter, it is necessary to state the methods which are used for this purpose. METHOD A. By means of dialysis. Add a little arsenious acid or some corrosive sub- limate to some minced meat (to represent the organic matter, such as stomach, liver, &c., of an actual poison- ing case), stir this in about 200 c.c. of distilled water and add a little pure hydrochloric acid. Place this 134 PRACTICAL CHEMISTRY. mixture on a dialyser and float it in a litre of water contained in a porcelain basin. After the lapse of twenty-four hours concentrate the contents of the basin on a water-bath, precipitate the arsenic or mercury by sulphuretted hydrogen, and confirm their presence by the tests already given. This method may be applied to the separation from organic matter of any crystalloid substance whether inorganic or organic. METHOD B. By destroying the organic matter. Place a similar mixture of meat with some metallic poison in a porcelain basin, add an equal bulk of strong pure hydrochloric acid, and warm on the water-bath, adding from time to time a few crystals of potassium chlorate and replacing the water removed by evapora- tion. Continue in this way until the organic matter has completely disappeared, heat to expel chlorine, filter, pass sulphuretted hydrogen through the warm solution, and test the precipitate by the reactions previously given. METHOD C Separation of arsenic from organic matter. Place a mixture of meat and arsenious acid in a glass retort, add some rock salt and pure sulphuric acid in quantity insufficient to decompose all the salt. Heat the retort and collect the distillate in a receiver kept very cold. The arsenic pressnt is converted by this treatment into trichloride, which being volatile passes into the receiver, and may be precipitated with sulphuretted hydrogen or recognised by other tests for arsenic. PART V. REACTIONS OF THE RARE METALS. 14-7. THE following rare metals are considered here by themselves, and their position in the various groups indicated. For their separation from each other and from the commonly occurring metals, a larger manual must be consulted. Tungsten, Thallium, Palladium, Rhodium, Os- mium, Ruthenium, Gold, Platinum, Iridium, Molyb- denum, Selenium,* Tellurium, Uranium, Indium, Beryllium, Zirconium, Cerium, Lanthanum, Didy- mium, Titanium, Tantallum, Vanadium, Lithium, Caesium, Rubidium. * Selenium, although a metalloid, is conveniently included here. GROUPING OF THE RARE METALS. 137 Grouping of the Rare Metals. 148. PRECIPITATED IN SILVER GROUP. Tungsten (as Tungstic Acid), Thallium (as Chloride). PRECIPITATED IN COPPER GROUP. Palladium, Rhodium, Osmium, Ruthenium. PRECIPITATED IN ARSENIC GROUP. Gold, Platinum, Iridium, Molybdenum, Tellurium, Selenium. PRECIPITATED IN IRON GROUP. Uranium, Indium (Thallium), as sulphides ; Beryl- lium, Zirconium, Cerium, Lanthanum, Didymium, Titanium, Tantalum, as hydrated oxides. PRECIPITATED AS SULPHIDES ON ADDING HC1 TO THE FILTRATE FROM THE IRON GROUP. Vanadium (Tungsten). FOUND IN POTASSIUM GROUP. Lithium, Caesium, Rubidium. i 3 8 PRACTICAL CHEMISTRY. Reactions of the Rare Metals of the Silver Group. 149. TUNGSTEN. W, c.w. 184. 1. HC1 produces a white precipitate of Tungstic acid H 2 WO 4 , insoluble in excess, but soluble in (NH^HO. H 2 SO 4 or HNO 3 produces the same pre- cipitate. 2. (NH 4 ) 2 S does not precipitate tungstates of the alkalies, until an acid be added, when tungsten tri- sulphide WS 3 is precipitated as a light brown powder. 3. K 4 Fe(CN) 6 in acid solutions yields a reddish brown coloration, and on standing, a precipitate of the same colour. 4. SnCl 2 produces a yellow precipitate. On adding HC1 and heating, the precipitate becomes blue. 5. Zn in presence of phosphoric acid produces a bright blue colour. 6. With microcosmic salt in the reducing flame a blue bead is obtained, which changes to red on heating with FeSO 4 . 150. THALLIUM. Tl, c.w. 203-5. 1. HC1 in strong solutions produces a white precipi- tate of T1C1, soluble in a large quantity of water. 2. (NH 4 ) 2 S produces a black precipitate of T1 2 S, insoluble in (NH 4 )HO, but soluble in HC1, H 2 SO 4J and HNO 3 . 3. KI precipitates Tl I (yellow) even in dilute solu- tions. 4. PtCl 4 precipitates a double chloride of thallium and potassium Tl 2 PtCl 6 of an orange colour. TSTS FOR PALLADIUM. 139 5 Zn added to thallium solutions precipitates the rnetal. 6. Thallium salts colour the lamp flame intensely green. Reactions of the Rare Metals of the Copper Group. 151. PALLADIUM. Pd, c.w. io6'6. 1. H 2 S precipitates PdS as a black powder, inso- luble in (NH 4 ) 2 S, and soluble in boiling HC1 and aqua regia. 2. Hg(CN) 2 precipitates palladium cyanide, Pd(CN)2, yellowish white and gelatinous, soluble in (NH 4 )HO and in HC1. 3. Kl produces a black precipitate of PdI 2 . 4. KC1 in strong solutions precipitates K 2 PdCl 6 in yellow needles, insoluble in absolute alcohol, but soluble in water, forming a dark-red fluid. 152. RHODIUM. Rh, c.w. 104-2. 1. H 2 S precipitates Rh 2 S 3 (brown), especially in warm solutions ; it is insoluble in (NH 4 ) 2 S, but soluble in hot HNO 3 . 2. KHO produces a yellowish precipitate of Rh(HO) 3 H 2 O, soluble in excess of the reagent ; on boiling the solution, brown Rh(HO) 3 is precipitated. 3. Zn produces a precipitate of metallic rhodium. 4. Dry compounds when ignited in a current of hy- drogen are reduced to metal, insoluble in aqua regi:i, but soluble on fusing with HKSO 4 . HO PRACTICAL CHEMISTRY. 153. OSMIUM Os, c.w. 199*1. 1. H 2 S in presence of acid precipitates OsS (brown- ish black), insoluble in (NH 4 ) 2 S. 2. OsO 4 (osmic acid) decolorizes indigo solution. 3. KI is decomposed with liberation of iodine. 4. Na 2 SO 3 yields a violet coloration, and after a time OsSO 3 (deep blue) separates out. 5. FeSO 4 precipitates black OsO 2 . 6. Zn in presence of acids precipitates the metal. 7. Dry compounds when ignited in a current of hy- drogen are reduced to metal. 8. The metal and the mono- and di-oxides, when heated in air, yield OsO 4 , which is recognized by its peculiar irritating smell (resembling chlorine). This is an exceedingly characteristic reaction. 154-. RUTHENIUM. Ru, c.w. 104-4. 1. H 2 S produces no immediate precipitate, but on standing, the solution turns blue, and brown Ru 2 S 3 is precipitated. 2. (NH 4 ) 2 S precipitates Ru 2 S 3 , difficultly soluble in excess. 3. KHO precipitates black sesquioxide Ru(HO) 3 , insoluble in excess, but soluble in acids. 4.. KCNS in pure solutions produces on standing, a red coloration, which becomes first purple, and then on heating violet. 5. Zn producej a blue coloration, and ultimately the metal separates out. TESTS FOR GOLD AND PLATINUM. 141 Reactions of the Rare Metals of the Arsenic Group 155. GOLD. Au, c.w. 197. 1. H 2 S produces in cold solutions black Au^S 3 , in hot solutions brown Au 2 S. Both precipitates are in- soluble in HNO 3 and in HC1, but dissolve in aqua regia. They dissolve in yellow, but not in colourless ammonium sulphide. 2. (NH 4 ) 2 S produces a black precipitate of Au 2 S 3 . 3. (NH^HO in strong solutions precipitates ammo- nium aurate (fulminating gold). 4. FeSO 4 produces a precipitate of the metal, as a brown powder, which, when rubbed, assumes a yellow colour and metallic lustre. 5. KNO 2 also produces a precipitate of the metal. 6. SnCl 2 + SnCl 4 produce a purplish precipitate (purple of Cassius) even in highly dilute solutions ; the precipitate is insoluble in HC1. 7. Heated on charcoal with Na 2 CO 3 before the blow- pipe, malleable yellow beads of the metal are ob- tained. (See also flame reaction, 59.) 156. PLATINUM. Pt, c.w. 197*5. 1. H 2 S produces in cold solutions on standing, brown PtS 2 ; on heating, the precipitate forms at once. It is soluble in a large excess of yellow ammonium sulphide ; it is insoluble in HC1 and in HNO 3 , but soluble in aqua regia. 2. (NH4) 2 S produces the same precipitate as H 2 S. 142 PRACTICAL CHEMISTRY. 3. NH 4 C1 produces a yellow crystalline precipitate of (NH 4 ) 2 PtCl 6 , more sparingly soluble in alcohol than in water. On heating the dried precipitate, metallic platinum is left in a finely divided state. 4. KC1 produces a yellow crystalline precipitate of K 2 PtCl 6 , also more insoluble in alcohol than in water. On heating the dried precipitate, a mixture of Pt and KC1 is obtained. 5. SnCl 2 produces no precipitate, but, in solutions containing HC1 in excess, yields a dark brown colora- tion due to the formation of dichloride. 6. FeSO 4 on long-continued boiling produces a pre- cipitate of the metal. 7. Fused on charcoal with Na 2 CO 3 before the blow-pipe flame, yields the metal as an infusible grey powder. (See also flame reaction, 59.) 157. IRIDIUM. Ir, c.w. 198. 1. H 2 S first removes the colour and sulphur sepa- rates out ; afterwards brown iridium sulphide is pre- cipitated. 2. (NH 4 ) 2 S produces the same precipitate, soluble in excess. 3. KC1 produces a dark brown precipitate of K 2 IrCl 6 , insoluble in a concentrated solution of the reagent. 4. NH 4 C1 produces a dark red precipitate of (NH 4 ) a IrCl 6 . 5. FeSO 4 decolorizes solutions of iridium, but pro- duces no precipitate. 6. Zn precipitates the metal as a black powder. TESTS FOR MOLYBDENUM. 143 158. MOLYBDENUM. Mo, c.w. 96. 1. H 2 S precipitates in warm solutions, after some time, brownish black MoS 3 , soluble in (NH 4 ) 2 S. 2. Zn, in acid solutions, produces a blue coloration, which changes to green, and lastly to black, when Mo 2 O 3 separates out. 3. KCNS, when added to an HC1 solution, produces on addition of Zn a crimson coloration. 4. Na 2 HPO 4 produces, in a nitric acid solution of ammonium molybdate, a yellow precipitate (in dilute solutions only after the lapse of some time). The pre- cipitation is aided \yy gentle heating. 159. TELLURIUM. Te, c.w. 128. 1. H 2 S produces, in presence of acid, a brown pre- cipitate of TeS 2 , easily soluble in (NH 4 ) 2 S. 2. H 2 O produces a white precipitate of H 2 TeO 3 . 3. KHO or Na 2 CO 3 , when added to an HCl solu- tion, produces a white precipitate of H 2 TeO 3 , soluble in excess of either reagent. 4. Na^SO 3 or SnCl 2 produces a black precipitate of the metal. 5. Zn precipitates the metal as a black powder. 6. Fused with Na 2 CO 3 , yields Na 2 Te, which, on treatment with HCl, yields H 2 Te, recognized by its very disagreeable odour. 160. SELENIUM. Se, c.w. 79-5. i. H 2 S produces, in presence of acid, a yellow pre- cipitate of doubtful composition, which darkens in colour on heating, and is soluble in (NH 4 ) 2 S. 144 PRACTICAL CHEMISTRY. 2. BaCl 2 produces a white precipitate of BaSe, soluble in HC1. 3. SO 2 and SnCl 2 produce in acid solutions a red precipitate of Se. 4. Heated in the reducing flame on charcoal, selenium compounds yield a highly characteristic smell, resem- bling that of horse-radish. Reactions of the Rare Metals of the Iron Group. 161. URANIUM. U, c.w. 120. 1. (NH 4 ) 2 S produces a brown precipitate of oxysul- phide, soluble in colourless ammonium sulphide to a black liquid. 2. Alkalies produce a yellow precipitate of hydrated oxide, insoluble in excess of the reagent. 3. (NH 4 ) 2 CO 3 produces a precipitate of the double carbonate of uranium and ammonium, soluble in excess of the reagent. 4. BaCO 3 produces a precipitate of hydrated oxide in cold solutions. 5. K 4 Fe(CN) 6 produces a reddish brown precipitate. 6. Fused with borax in the reducing flame, uranium compounds yield pale green beads. 162. INDIUM. In, c.w. 113-4. i. (NH 4 ) 2 S, in presence of (NH 4 )HO and tartaric acid, produces a white precipitate, which turns yellow on addition of acetic acid. TESTS FOR BERYLLIUM. 145 2. Alkalies produce a precipitate of the hydrate, in- soluble in excess of the reagent. 3. Alkaline carbonates precipitate the carbonate, soluble in excess of (NH 4 ) 2 CO 3? but not in excess of K 2 CO 3 or Na 2 CO 3 . 4. BaCO 3 precipitates solutions of indium com pletely. 5. K 4 Fe(CN) 6 produces a white precipitate. 6. Zn precipitates the metal in shining plates. 7. Heated with Na 2 CO 3 in the reducing blow-pipe flame, metallic beads are obtained. 8. Indium compounds impart a bluish violet tinge to the flame. 163. BERYLLIUM. Be, c.w. 9-3. 1. (NHj) 2 S precipitates the hydrate (flocculent) like alumina in appearance, but differs- from it in dis- solving on continued boiling with NH 4 Ci. Beryllium chloride is formed and ammonia is driven off. 2. KHO precipitates the hydrate, soluble in excess. If this solution be diluted and well boiled, the beryl Hum hydrate is re-precipitated. 3. Na 2 CO 3 precipitates the carbonate (white), soluble in a large excess. 4. (NH 4 ) 2 CO 3 precipitates the carbonate (white), easily soluble in excess of the reagent. 5. BaCO 3 in cold solutions precipitates the beryl- lium completely. 164. ZIRCONIUM. Zr, c.w. 89*6. i. (NH 4 ) 2 S produces a precipitate of hydrate, in- L 146 PRACTICAL CHEMISTRY. soluble in excess. KHO and (NH^HO produce the same precipitate, insoluble in excess of these reagents. 2. Alkaline carbonates precipitate the carbonate as a flocculent powder, soluble in a large excess of K 2 CO 3 and in a small excess of (NH 4 ) 2 CO 3 . From this latter solution the hydrate is re-precipitated on boiling. 3. BaCO 3 does not precipitate zirconium completely, even on boiling. 165. CERIUM. Ce, c.w. 92*2. 1. (NH 4 )HO produces a precipitate of a basic salt, insoluble in excess. 2. KHO produces a precipitate of the white hydrate, insoluble in excess of the reagent. The precipitate becomes yellow on exposure to the air. 3. (NH 4 ) 2 CO 3 produces a white precipitate, soluble in excess of the reagent. 4. C 2 H 2 O 4 precipitates cerous oxalate, insoluble in excess, but soluble in HC1. 5. BaCO 3 precipitates cerium completely after the lapse of some time. 6. Chlorine, passed through a solution mixed with sodium acetate, precipitates the peroxide (light yellow). 166. LANTHANUM. La, c.w. 92-9. 1. (NH 4 )HO produces a precipitate of a basic salt. 2. KHO precipitates the hydrate (white), which does not alter on exposure to air. 3. (NH 4 ) 2 CO 3 produces a precipitate, insoluble IP excess of the reagent. TESTS FOR DIDYMIUM. 147 4. C 2 H 2 O 4 precipitates the oxalate, insoluble in ex- cess, but soluble in HC1. 167. DIDYMIUM. D, c.w. 95. 1. (NH 4 )HO produces a precipitate of a basic salt, insoluble in excess of the reagent, but sparingly soluble in N.H 4 C1. 2. KHO precipitates the hydrate (white), which does not alter on exposure to air. 3. (NH 4 ) 2 CO 3 produces a white precipitate, insoluble in excess of the reagent, but soluble in NH 4 C1. 4. C 2 H 2 O 4 produces an almost complete precipitation of the oxalate, soluble in hot HC1. 5. Fused with microcosmic salt in the reducing flame, a reddish violet bead is obtained. 6. Fused with carbonate of soda in the oxidizing flame, greyish white beads are obtained. 168. TITANIUM. Ti, c.w. 50. 1. (NHJHO or (NH 4 ) 2 S precipitates H 2 TiO 3 (ti- tanic acid), insoluble in excess of either reagent. 2. KHO also precipitates H 2 TiO 3 , insoluble in excess. 3. BaCO 3 produces the same precipitate. 4. Na 2 S 2 O 3 when boiled with titanium solutions pre- cipitates them completely. 5. Zn produces, first a blue coloration, then a blue precipitate, which ultimately becomes white. 6. Fused with FeSO 4 and microcosmic salt in the reducing flame, a bright red bead is obtained. L 2 148 PRACTICAL CHEMISTRY. Reactions of the Rare Metals precipitated as Sulphides on adding HC1 to the nitrate from the Iron Group. 169. VANADIUM. V, c.w. 51 '3. 1. H 2 S produces no precipitate in acid solutions, but a blue coloration by reducing to a lower oxide. 2. (NH 4 ) 2 S produces in a solution containing H 2 SO 4 a brown precipitate of the sulphide, soluble in excess of the reagent. 3. NH 4 C1 in solutions of alkaline vanadates, pre- cipitates the metal completely as ammonium meta- vanadate (white). This precipitate loses ammonia on heating, and leaves a residue of V 2 O 6 . 4. K 4 Fe(CN) 6 produces, in acid solutions, a green precipitate, not dissolved by acids. 5. SO 2 , or oxalic acid, reduces acid solutions of V 2 O 6 to a lower oxide of vanadium, of a bright blue colour. 6. Zn, added to a solution of V 2 O 6 in H 2 SO 4 (diluted with H 2 O), produces a series of changes in colour, and ultimately a violet solution is obtained, which rapidly becomes brown in air by oxidation. Reactions of the Rare Metals found in Potassirm Group. 170. LITHIUM. Li, c.w. 7. 1. Pt.Cl 4 produces no precipitate. 2. Na 2 HPO 4 produces, on boiling the solution, a vhite precipitate of 2 Li 3 PO 4 + H 2 O. The precipitate TESTS FOR CAESIUM. 149 is soluble in HC1, and is not re-precipitated by (NH 4 )HO unless the solution be boiled. 3. Lithium compounds tinge the lamp flame bright crimson. 171. CESIUM. Cae, c.w. 133. 1. PtCl 4 produces a crystalline light yellow precipi- tate of 2(CaeCl) -+- PtCl 4 , insoluble in boiling water. (The corresponding potassium salt is dissolved by repeated treatment with boiling water.) 2. Tartaric acid produces a precipitate of the acid tartrate of caesium, more soluble in water than the cor- responding compound of rubidium. 3. Volatile caesium salts colour the flame violet. 172. RUBIDIUM. Rb, c.w. 85-4. 1. PtCl 4 produces a crystalline light yellow precipi- tate of 2 (RbCl) 4- PtClfc insoluble in boiling water. 2. Tartaric acid produces a precipitate of the acid tartrate of rubidium, which is much more insoluble in water than the corresponding caesium compound. 3. Volatile salts of rubidium colour the flame violet. PART VI. QUANTITATIVE ANALYSIS. EXAMPLE I. Determination of Filter Ash. WEIGH accurately a porcelain or platinum crucible which has been heated to redness and then allowed to become quite cold. Select six Swedish filter-papers of FIG. 20. uniform diameter, fold each as in the sketch, and then coil round it a thin platinum wire, leaving a loose end (A). Place the crucible on a glazed porcelain tile, QUANTITA TIVE ANAL YSIS. 1 5 1 kindle the point of the paper and hold it over the crucible until the paper is reduced to ashes. Detach the ashes carefully from the wire and allow them to fall into the crucible. Repeat this with the remaining five filter-papers, and if any ash has fallen on the porcelain tile sweep it with a feather into the crucible, which is then heated to redness until the contents no longer contain any black portions. Record the results thus : Crucible + ash from six filters (diam. = )...=- Crucible Ash from six filter-papers . . = grm. Weight of each filter ash (diam. = )...= grm. EXAMPLE II. Determination of Sulphuric Acid. Weigh out about 0*2 gramme of dry powdered potassium sulphate, or any other sulphate soluble in water. This is best done by placing the salt in a dry test-tube and weighing this with its contents, then carefully shake the required quantity into a dry beaker, and re-weigh the tube. The difference between the two weighings is the weight of salt taken. To the 152 PRACTICAL CHEMISTRY. contents of the beaker add distilled water and a few drops of pure hydrochloric acid, and heat over wire gauze with constant stirring till the liquid boils. Re- move the lamp and add drop by drop clear solution of barium chloride in very slight excess. Heat again to the boiling-point, stirring as before ; remove the lamp and allow the precipitate to subside. To the clear liquid add a drop of barium chloride solution to see if the precipitation is complete. If no further precipita- tion takes place, heat again to the boiling-point for some minutes, allow to settle and then pour off the liquid through a Swedish filter-paper, leaving the precipitate in the beaker. Now add hot distilled water, one or two drops of hydrochloric acid, and boil as before, allow to settle and again pour the clear portion through the filter.* Repeat this two or three times, and finally transfer with the aid of a wash bottle every trace of the precipitate to the filter, wash repeatedly with hot distilled water until a few drops of the filtrate give no turbidity with silver nitrate, showing that the excess of barium chloride and hydrochloric acid has been removed. When the washing is com- plete, get the precipitate as much as possible to the point of the filter, then dry it by placing the funnel with its contents in an air-bath. When dry, transfer the precipitate as completely as possible to a weighed porcelain or platinum crucible, burn the filter (with the same precautions as in the previous example), adding the ash to the contents of the crucible, and * This method is called "washing by decantation." QUANTITATIVE ANALYSIS. 153 heat to redness for some minutes, allow to become quite cold and weigh. Record the results thus : Tube + potassium sulphate . = Tube - potassium sulphate . = Potassium sulphate used . . = grm. Crucible -f barium sulphate -j- ash ..... Crucible . . . = Barium sulphate -{- ash . . = Filter ash ' Barium sulphate . Reaction : K 2 SO 4 + BaCl 2 = BaSO 4 + 2KC1 174-2 + 208 = 233 + H9' 2 Percentage composition=By calculation By analysis. Difference. K 2 44'39 S0 4 55'u 1 00 '00 154 PRACTICAL CHEMISTRY. EXAMPLE III. Determination of Hydrochloric Acid, Weigh from a tube into a beaker as before about 6*1 to 0*2 gramme pure common salt or other soluble chloride, dissolve in distilled water, add a few drops of pure nitric acid and warm the solution. To the warm solution add solution of silver nitrate drop by drop until the precipitation is complete. This is easily ascertained by allowing the precipitate to subside and then adding a drop of silver nitrate solution to the clear liquid. If no further precipitation take place, cover the beaker and keep it in the dark for some hours, when the supernatant liquid should be found quite clear. Now pour the clear liquid through a filter retaining the precipitate in the beaker. Add to it hot distilled water and a drop of nitric acid and heat (stirring constantly) till it boils. Allow to subside, then pour the clear liquid through the filter as before. Repeat this process twice, then transfer the silver chloride to the filter, wash with boiling water alone, until a few drops of the filtrate give no turbidity on addition of a drop of hydrochloric acid. Finally wash the precipitate into the point of the filter, cover the funnel and place it in an air-bath. When quite dry transfer the chloride of silver as completely as possible to a weighed porcelain crucible, keeping the filter-paper for separate treatment. Heat the crucible gently at first, and then more strongly till the chloride QUANTITATIVE ANALYSIS. 155 just begins to fuse, allow to become perfectly cold and weigh. Now fold the filter-paper as in the sketch, keeping the portion to which silver chloride adheres (A) in the FIG. 21. middle of the coil, which should be as tightly rolled up as possible. Wind round it a platinum wire (previously weighed on a watch glass), and light the filter-paper, holding it over the watch glass, and allow it to burn to ashes. When it ceases to glow apply a 156 PRACTICAL CHEMISTRY. lamp flame till it is completely burned, lay it with the platinum wire on the watch glass and weigh when cold. By this process the silver chloride is reduced to metallic silver, and is weighed as such, and the quantity of silver chloride to which it corresponds is calculated. (108 parts of silver correspond to 143*5 parts of silver chloride.) Record the results thus : Tube -\- common salt . . Tube - common salt = Common salt used ... . = grin. Crucible + silver chloride Crucible , Silver chloride . . \ ' . = grm. Watch glass + wire + ash + silver Watch glass -f- platinum wire . Ash + silver Ash , Silver . . Silver chloride from crucible Silver chloride from ash . Total silver chloride . , . = grm, QUANTITATIVE ANALYSIS 157 Reaction NaCl -f AgN0 3 = AgCI + NaNO 3 58-5 + 170 = 143-5 + 85 Percentage composition=By calculation. By analysis. Difference. Na 39-32 Ci 60-68 lOO'OO EXAMPLE IV. Determination of Phosphoric Acid. Weigh out into a beaker about 0*5 gramme of sodium phosphate, (Na 2 HPO 4 + I2H 2 O), dissolve it in water, and add magnesia mixture till the precipitation is complete (add ammonia if the liquid has not a strongly ammoniacal smell) and allow to stand for twenty-four hours. Filter and wash the precipitate with water containing one-fourth of its volume of ammonia solution, until a few drops of the filtrate when neu- tralized with nitric acid give no precipitate on addition of silver nitrate. Dry the precipitate in the air-bath, transfer to a weighed porcelain or platinum crucible adding the filter-ash to the contents of the crucible, heat gently at first, then strongly, and weigh when cold. Record the results as in Example II., substituting sodium phosphate for potassium sulphate, and mag- nesium pyrophosphate for barium sulphate. IS* PRACTICAL CHEMISTRY. Reaction : Na 3 HPO 4 + MgSO 4 + (NH 4 )HO = Mg(NH 4 )PO 4 + Na 2 SO 4 + H 2 O. 142 + 120 + 25 = I 37 + 142 + 18 On ignition the ammonium-magnesium phosphate is converted into magnesium pyrophosphate, Mg 2 P 2 O r . EXAMPLE V. Determination of Carbonic Anhydride. Weigh out from one to two grammes of Iceland spar (calcium carbonate), place it in the flask A (Fig. 22) along with a little water, then insert a small test-tube containing hydrochloric acid in such a way that it is supported by the side of the flask and without risk of the acid escaping till required to do so. Place strong sulphuric acid in the flask B till it is nearly half full, adjust the stoppers in each flask, and close the opening c with a piece of india-rubber tubing and a small piece of glass rod. Now weigh the double flasks accurately, and incline the apparatus so that a little of the hydrochloric acid in the test-tube mingles with the water and Iceland spar in A. Carbonic anhydride is at once evolved and is freed from moisture by passing through the sulphuric acid in B. When the evolution of gas slackens, allow more acid to escape from the tube, and continue this until the carbonate is completely dissolved. Then heat the flask A with a small flame so as to expel the carbonic QUANTITATIVE ANALYSIS. 159 anhydride dissolved by the liquid ; remove the stopper at C, attach an india-rubber tube to the exit tube of B and slowly draw air by suction through the apparatus until every trace of carbonic anhydride is removed. When the apparatus is quite cold replace the stopper at c and again weigh. The loss in weight represents the carbonic anhydride originally present in the spar. FIG. 22. A new and very light form of the apparatus is shown in Fig. 23. The weighed carbonate is placed in the tube A and hydrochloric acid in the bulb tube B. By closing the point of this tube with the finger while the acid is being added and then adjusting the hollow stopper, the acid is prevented from escaping. The drying tube C contains either strong sulphuric acid or calcium chloride, and its stopper is kept closed 160 PRACTICAL CHEMISTRY. until the apparatus is weighed. It is then opened and acid admitted to the tube by turning the stopper in B (so as to admit air) as often as may be required. Record the results thus : Tube + Iceland spar . ; . . . = Tube = Iceland spar used Carbonic anhydride apparatus . ,, after experiment Carbonic anhydride . . . . . = grm. Reaction : CaCO 3 + zHCl = CaCl 2 + CO 2 + H 2 O ioo + 73 = in + 44 + 18 Percentage composition = By calculation. By analysis. Difference. CaO 56 C0 2 44 EXAMPLE VI. Determination of Iron. Weigh from a tube into a beaker about o'2 gramme of ferrous ammonium sulphate * (FeS O 4 -f- (N H 4 ) 2 SO 4 -|- 6H 2 O), which has been previously powdered and dried by pressing between filter-paper. Dissolve the salt in water, add hydrochloric acid and a few drops of nitric acid (to convert the ferrous into ferric salt) and heat till it nearly boils. If sufficient nitric acid has been added the colour of the solution ought to change to brownish red. Allow to cool, add ammonia in slight excess, heat until the liquid nearly boils, stirring constantly and allow to settle. Pour off the clear hot * This salt contains i/yth of its weight of iron. QUANTITATIVE ANALYSIS. 161 jquid through a filter and repeatedly wash by decanta- tion, allowing the precipitate to subside each time. Then transfer the precipitate to the filter, wash thoroughly with boiling water and dry in an air-bath. Transfer the precipitate as completely as possible to a weighed crucible, add the filter ash to its contents, and heat to redness for some time. Allow to cool and weigh the iron sesquioxide (Fe 2 O 3 ). [The student will be able to judge from the previous examples how to record the results.] EXAMPLE VII. Determination of Calcium. Weigh out about 0*2 gramme of Iceland spar and dissolve.it in dilute hydrochloric acid in a beaker covered all the time with a watch glass, to prevent loss by spirting. When completely dissolved, wash any liquid adhering to the watch glass into the beaker, add ammonia in slight excess, warm the solution, and to the hot liquid add solution of ammonium oxalate until complete precipitation is effected. Now add ammonia until the fluid smells of it, cover the beaker and leave it in a warm place for at least twelve hours. Pour the clear liquid through a filter, leaving the precipitate in the beaker, wash it several times with boiling water and finally transfer it to the filter. Any precipitate adhering to the beaker is removed by rubbing with a glass rod tipped with india-rubber. Dry the precipitate in the air-bath, transfer to a M 162 PRACTICAL CHEMISTRY platinum crucible (to which the filter ash is added) and heat at first gently over the Bunsen lamp, finally to a red heat. Now heat the crucible over a Bunsen blowpipe flame for a considerable time, and repeat this until the crucible ceases to lose weight. This serves to convert the calcium oxalate into oxide (CaO), from the amount of which the percentage of calcium in the spar is calculated. EXAMPLE VIII. Determination of Copper. Weigh out about D'2 gramme of crystallized copper sulphate (CuSO 4 -f- 5H 2 O), into a porcelain basin, dissolve in water, and heat till nearly boiling. Add now pure solution of potash drop by drop until no further precipitation occurs, and continue the heating for some little time, but without allowing the liquid to boil. When the precipitate has subsided, pour the clear colourless liquid through a filter, wash three times by decantation using boiling water each time, then collect on the filter, wash thoroughly with hot water and dry in the air-bath. Detach the dried precipitate as completely as possible from the paper, place it in a weighed porcelain crucible, and burn the filter with the same precautions as with silver chloride (page 155), add the ash to the contents of the crucible and moisten it with one or two drops of nitric acid. Carefully expel the excess of acid by gentle heating, and ignite until the copper nitrate formed is converted QUANTITA TIVE ANAL YSIS. 163 into oxide. When quite cold, weigh, and calculate the amount of copper in the salt from the weight of copper oxide found. EXAMPLE IX. Preparation of Potassium Permanganate Solution for Determination of Iron, &c. Place between 5 and 6 grammes of powdered potassium permanganate in a glass-stoppered bottle capable of holding rather more than a litre. Add now a litre of water, and shake the bottle until the salt is completely dissolved. When this is the case add the solution to a glass-stoppered burette, filling up to the top and then running out the solution till the liquid stands exactly at the uppermost line marked o. Weigh accurately about 0*5 gramme of dry powdered ferrous ammonium sulphate (FeS O 4 (N H 4 ) 2 S O 4 + 6H 2 O) ; place it in a beaker of about half a litre capacity, and dissolve (without the aid of heat) in about 200 c.c. of water. When dissolved add dilute sulphuric acid, place the beaker on a white porcelain tile and run in the permanganate solution from the burette, stirring constantly until the solution is distinctly pink coloured. When the reaction is seen to be nearly over, the per- manganate must be added drop by drop. Now read off the point at which the permanganate stands in the burette and record the result. Weigh out a second portion of the double sulphate and repeat the process until you obtain three or four closely agreeing experiments. M 2 164 PRACTICAL CHEMISTRY. Record the results thus : Exp. I. gramme, double salt required c.c. permanganate. > II. ,, c.c. III. c.c. IV. c.c. Therefore looc.c. permanganate = grm. double salt. [As the salt contains one-seventh of its weight of iron, the results are easily calculated for metallic iron.] Reaction : 2 KMn0 4 + ioFeSO 4 + 8H 2 SO 4 = 2MnSO 4 +5Fe 2 (SO 4 ) 3 + K 2 SO 4 + 8H 2 O EXAMPLE X. Determination of Iron ( Volumetric). Weigh out about 0*1 gramme of clean iron wire and dissolve it with the aid of heat in dilute sulphuric acid in a flask through which a current of carbonic anhydride is continually passed. When complete solution has been effected, allow to cool in the current of carbonic anhydride, dilute with water, and triturate with permanganate solution as in the previous ex- periment. In exact experiments a correction must be made for the impurities in the wire, which usually amount to 0*4 per cent., in other words every deci- gramme of iron wire corresponds to only 0-0996 gramme pure iron. Repeat this experiment two or three times and compare the results with those obtained in Example IX. QUANTITATIVE ANALYSIS. 165 EXAMPLE XI. Determination of Oxalic Acid. Weigh out between I and 2 decigrammes of pure crystallized oxalic acid (C 2 H 2 O 4 + 2H 2 O), place in a beaker with about 200 c.c. of water, add dilute sulphuric acid and warm gently. Add permanganate from the burette, stirring constantly, until the pink colour is permanent even on warming the solution. Repeat the experiment two or three times and compare the results with the preceding ones. Every 112 parts of iron as ferrous salt correspond to 126 parts of crystallized oxalic acid. 2 KMnO 4 + 5C 2 H 2 4 + 3H 2 SO 4 = K 2 SO 4 + 2MnSO 4 -f ioCO 2 + 8H 2 O. EXAMPLE XII. Preparation of Standard Oxalic Acid. Weigh out exactly 63 grammes of pure crystallized oxalic acid. This is best done by accurately weighing a large watch glass, adding 63 grammes to the weights, and then oxalic acid to the watch glass until the balance is again in equilibrium. With the aid of a dry funnel transfer the whole of the crystals to a litre flask, wash the watch glass and funnel, and nearly fill with distilled water. Dissolve by shaking the flask, and when quite dissolved fill up exactly to the mark on the neck, and shake again thoroughly. Test the strength by means of the standard permanganate 166 PRACTICAL CHEMISTRY. solution previously made, using for this purpose about 10 or 20 c.c. of the oxalic acid solution. Now weigh out in the manner already described, 21*3714 grammes of pure recently ignited sodium carbonate, transfer it to a quarter-litre flask, dissolve in water, fill up to the mark and shake thoroughly. Pour this solution into a burette and allow 20 c.c. ( = I gramme Na 2 O), to flow out into a porcelain basin containing about looc.c. water. Add a few drops of neutral litmus solution, and then the standard solution of oxalic acid until the blue colour disappears. Heat the solution to the boiling-point, again add standard acid till the colour changes from purple to bright red, and repeat this until the latter colour is constant. Repeat the experiment with another 20 c.c. of sodium carbonate solution, and if the two experiments agree, compare them with the previous ones in which the strength was estimated by means of permanganate. Use the standard acid for determining the amount of alkali in accurately weighed quantities of caustic soda, caustic potash, soda ash, &c., &c. EXAMPLE XIII. Prepa lation of Standard Siilphuric Atict. Measure out a litre of water and place it in a stoppered bottle, then run into it from a burette 34 c.c. of strong pure sulphuric acid, add 20 c.c. of water, shake thoroughly, and allow to become quite cold. Fill a burette with the diluted acid and ascertain, just QUANTITA TIVE ANAL YSIS. 167 as in the previous example, how much of it is required to saturate 20 c.c. of the sodium carbonate solution prepared in XII. From this calculate the weight of sulphuric anhydride in the litre. Take now 20 c.c. of the acid and determine the amount of sulphuric acid by precipitation with barium chloride as in II. Com- pare the two results, and if they agree well, take the mean as the quantity of sulphuric acid present in the solution z which must now be diluted so as to contain exactly 40 grammes SO 3 in the litre. The amount of water to be added is easily calculated. Thus if a litre is found to contain 41*6 grammes SO 3 , instead of 40, then by the proportion 40 : 41*6 : : 1000 = 1040, it is necessary to add 40 c.c. water to each litre of the acid. Use the standard acid so prepared for determining the amount of alkali in soda ash, &c. EXAMPLE XIV. .Preparation of Standard Soda. Dilute a solution of pure soda until the specific gravity as indicated by the hydrometer is about I "05, and place it in a burette. Now add 50 c.c. of the standard acid (XIII.) to about 100 c.c. of water contained in a basin, add a few drops of litmus solution, and then run in from the burette caustic soda solution until the colour just changes to blue. The solution is required of such a strength that i c.c. acid is exactly neutralised, by i c.c. soda. By proceeding as above the soda will 1 63 PRACTICAL CHEMISTRY. be stronger than the acid, and must be diluted with the proper quantity of water. If, for example, 46 c.c, of soda are equal to 50 c.c. of acid, then 4 c.c. of water must be added to every 46 c.c. of the solution. By means of the standard soda ascertain the amount of acid in accurately weighed quantities of acetic acid, sulphuric acid, potassium bisulphate, &c. EXAMPLE XV. Determination of Chlorine ( Volumetric). Weigh out accurately 2 '3944 grammes of pure silver nitrate and dissolve in one litre of distilled water. (This is a convenient strength for the determination of the amount of chlorides in water, for other pur- poses a stronger solution is employed.) Place the silver solution in a burette and allow it to run drop by drop into 50 c.c. of ordinary water contained in a porcelain basin, and to which one or two drops of a solution of pure potassium chromate have been added. When the liquid becomes faintly red in colour read off the amount of silver nitrate used ; the number of cubic centimetres used indicates the amount of chlorine in 100,000 parts of the water examined. The potassium chromate serves to indicate when the precipitation of silver chloride is complete, since none of the red silver chromate is formed until all of the highly insoluble silver chloride is precipitated. QUANTITA TIVE ANAL YSIS. 169 EXAMPLE XVI. Estimation of Chlorine in Bleaching Powder. Weigh out exactly 4*95 grammes of pure dry arsenic trioxide, place it in a flask with some water, heat gently, and add pure crystallized sodium carbonate from time to time until complete solution has been effected. Allow to cool, transfer carefully to a litre flask, and fill up to the mark on the neck. This forms a deci-normal solution of arsenious acid, and may be used for the determination of iodine, chlorine, bleaching powder, &c. Reactions : As 2 3 -f 4l + 2H 2 = 4HI 4-As 2 6 . As 2 3 + Ca(C10) 2 = CaCl 2 + As 2 O 5 . Weigh out exactly 3*55 grammes of bleaching powder into a mortar, add water, and thoroughly mix with the aid of the pestle. Decant off the turbid fluid into a litre flask, leaving the sediment behind. Add more water to this, triturate again, transfer as before to the litre flask, and repeat this until all the bleaching powder has been transferred to the litre flask, then fill up to the mark. Shake thoroughly, and before allowing it to settle withdraw looc.c. by means of a pipette, and place it in a beaker. Add now from a burette the standard arsenic solution, stirring constantly, and testing from time to time (with the aid of a glass rod) the action of the fluid on strips of paper moistened i;o PRACTICAL CHEMISTRY. with a mixture of potassium iodide and starch solu- tion. At first the prepared paper is deeply marked blue, but as the arsenious acid is added the marks become fainter and fainter, and it is easy to determine the exact point when the paper ceases to be tinged. Read off from the burette the number of cubic centi- metres of arsenic solution used ; this number indicates at once the percentage of chlorine in the sample of bleaching powder. EXAMPLE XVII. Determination of Manganese Dioxide. Weigh accurately from 3 to 4 grammes of dry, finely-powdered pyrolusite, place it in the flask A, Fig. 22 (see page 159), add a little water, and about 6 grammes neutral sodium oxalate. Put about 50 c.c. strong sulphuric acid in the flask B, adjust the stoppers, and weigh the entire apparatus on an accurate balance. Now attach an india-rubber tube to the exit tube of B and gently suck air from A, so as to cause sulphuric acid to pass over from B into A. As soon as this is done an evolution of car- bonic anhydride begins ; the gas is dried by passing through the acid in B, and escapes by the exit tube. When the evolution of gas slackens repeat the suction, and continue in this way until every trace of pyrolusite has been dissolved. When this is the case draw more acid over into A, so as to heat the liquid and expel any dissolved carbonic anhydride, then QUANTITATIVE ANALYSIS. 171 remove the stopper at c and draw air through the apparatus to replace the carbonic anhydride. When perfectly cold weigh the apparatus again ; the loss in weight indicates the carbonic anhydride expelled, from the amount of which the quantity of manganese dioxide in the pyrolusite may be calculated. Reaction : MnOo + 2 H 2 S0 4 + Na 2 C 2 4 = MnSO 4 + Na 2 S0 4 + 2CO 2 + 2 H 2 O. 87 + 196 + 134 = 151 + 142 + 88 + 36. EXAMPLE XVIII. Determination of Lead. Weigh out about 0*5 gramme of re-crystallized lead nitrate, dissolve in water, add ammonium carbonate in slight excess and a few drops of ammonia, heat gently, and after some time, filter. Wash the precipi- tate .with distilled water 'and dry in the water-bath. Transfer the precipitate as completely as possible to a weighed porcelain crucible, ignite gently at first and then strongly, allow to cool and weigh the lead oxide (PbO) remaining. The filter must be burned separately as in the case of silver, and when thoroughly ignited may be added to the contents of the crucible. . 172 PRACTICAL CHEMISTRY. EXAMPLE XIX. Determination of Zinc. Weigh out about I gramme of zinc sulphate crystals (ZnSO 4 + 7 H 2 O), dissolve in water and heat nearly to the boiling-point, add by degrees solution of sodium carbonate till there is a slight excess, and then heat to the boiling-point. Allow to settle, decant the clear liquid, boil the residue with more water two or three times, then transfer to the filter, wash thoroughly,, and dry the precipitate. Transfer the dry precipitate as completely as possible to a crucible, add to it the ashes of the filter-paper, heat to redness, and when cold weigh the residue of zinc oxide (ZnO.) When these examples have been completed, the student may practise other determinations, using the same reactions. Thus, by weighing out quantities of barium chloride, silver nitrate, and magnesium sul- phate, and precipitating their solutions with sulphuric acid, hydrochloric acid and sodium phosphate re- spectively, the methods described in II., III., and IV. may be used for the determination of barium, silver, and magnesium. APPENDIX A. TABLE OF THE ELEMENTS, WITH THEIR SYMBOLS AND COMBINING WEIGHTS. Element. Symbol. Aluminium ..... Al ...... 27*3 Antimony ...... Sb ..... 120 Arsenic ....... As ...... 75*2 Barium ...... Ba ...... 137 Beryllium ...... Be ...... 9'3 Bismuth ...... Bi ...... 210 Boron ...... , B ...... n Bromine ...... Br ...... 80 Cadmium . ..... Cd ...... 112 Caesium ...... Cs ...... 133 Calcium ...... Ca ...... 40 Carbon ....... C ...... 12 Cerium ....... Ce ...... 92*2 Chlorine ...... Cl ...... 35-5 Chromium ..... Cr ...... 52 Cobalt ....... Co ...... 587 Copper ...... Cu . . . . . 63'$ Didymium . . . . D ...... 95 Erbium E 1116 174 PRACTICAL CHEMISTRY. Element. Symbol. Combining Fluorine F 19 Gold ..*.... Au ....... 197 Hydrogen H ...... i Indium In H3'4 Iodine I 127 Indium Ir 198 Iron Fe 56 Lanthanum La 92*9 Lead ....... Pb 207 Lithium Li 7 Magnesium Mg 24 Manganese Mn 55 Mercury Hg 200 Molybdenum Mo .... .96 Nickel Ni ...... 587 Niobium Nb 94 Nitrogen N 14 Osmium Os 199-1 Oxygen O 16 Palladium Pd io6'6 Phosphorus P 31 Platinum Pt 197-5 Potassium K 39*1 Rhodium Rh 104*2 Rubidium Rb 85-4 Ruthenium Ru ...... 104-4 Selenium . . ; . ; . . Se 79-5 Silver Ag . . . . , 108 TABLE OF THE ELEMENTS. 175 Element Symbol. Silicon ....... Si ...... 28 Sodium ....... Na ...... 23 Strontium ...... Sr ...... 87-5 Sulphur ...... S ...... 32 Tantalum ...... Ta ...... 182-3 Tellurium ...... Te ...... 128 Thalli im ...... Tl ...... 203-5 Thorium ...... Th ...... 1157 Tin ....... . Sn ...... 118 Titanium ...... Ti ...... 50 Tungsten ...... W ...... 184 Uranium ...... U ...... 120 Vanadium : .... V ...... 51-3 Yttrium ...... Y ...... 6r6 Zinc ........ Zn ...... 65*2 Zirconium . ..... Zr ..... 89-6 APPENDIX. APPENDIX B. WEIGHTS AND MEASURES. Measures of length. Measures of iveig'it. I metre = 10 decimetres (dcm.). i gram = 10 decigrams. i metre 100 centimetres (cm.). i gram = 100 centigrams, i metre = 1000 millimetres (mm,)- i gram = 1000 milligrams. 1000 metres = i kilometre. 1000 grams = i kilogram. 100 metres = i hectometre. 100 grams i hectogram. 10 metres = i decametre. . 10 grams = i decagram. i inch = 2 "539954 centimetres, i grain = 0*06479895 gram i foot = 3-0479449 decimetres. i ounce) i yard = 0-91438348 metre. (Troy.)/~ 3 1 i mile = t '6093149 kilometre. i pound\ (Avd.)j='453S9265kilog. i cwt. = 50 "80237689 kilogs Measures of capacity. i litre = i cubic decimetre. i cub. in.= 16-3861759 cubic a 'itre = 1000 cubic centi- centimetres. metres. i cub. ft. = 28-3153119 cubic i x dtrs = i kilolitre or stere. decimetres. i gallon ] =4 - 54345796Qlitr es (70,000 grains.)' APPENDIX. 177 APPENDIX C TREATMENT OF SILVER RESIDUES. SILVER residues generally consist chiefly of silvei chloride, and may be treated as follows. Add hydro- chloric acid in slight excess to precipitate any solu- ble silver salt that may be present, and allow the precipitate to subside. Pour off the supernatant liquid, and wash the precipitate several times (by de- cantation) with water. Place the precipitate in an evaporating basin, add solution of caustic soda and one or two pieces of sugar, and boil the mixture, stirring constantly. As soon as the precipitate settles rapidly to the bottom, leaving the liquid clear, the heating is stopped, and the precipitate washed with common water (by decantation) till no longer alkaline, and finally once or twice with distilled water. The moist precipitate is then dissolved in pure nitric acid, and evaporated to obtain crystals of silver nitrate. If any choloride has escaped reduction, the solution in nitric acid will .not be clear, in this case it must be diluted and filtered. When the quantity of silver residues is small, it is better to dissolve the finely divided silver in the smallest quantity of nitric acid possible, dilute with distilled water, and filter the solution. N TABLE IRON GROUP IN PRESENCE OF PHOSPHORIC Manganese, Zinc, Chromium, Barium, Strontium, and Mag- To the filtrate from the sulphides of the Cu and As groups add (N H 4 ) H O (till some time and filter. Wash well with water containing (N H 4 ), S, and finally RESIDUE. NiS,CoS,Si,0 2 , (Ca F 2 ). Wash and test a portion with micro- cosmic salt. A white powder remaining undissolved indi- cates Silicic Acid j the bead becoming blue in- dicates Cobalt. Examine the re- mainder of the pre- cipitate for Ni, as in Table F. * As Ca F 2 is gnaringly soluble in HC1, it may be found in this residue. Its presence is de- tected by inciner- ating a portion of the residue, and heating it 'with strong H a SO 4 . Evolution of HF indicates Calcium Flu- oride. 1 METHOD I. Cr is ABSENT. Boil down with a few drops of H NO; (yellow coloration indicates the presence of Iron). Add dilute H a S O 4 , boii again and filter. RESIDUE. FILTRATE. ( Ba S O 4> Sr S O,,f Ca S O .). Proceed as in Table G. III. RESIDUE. 1 Add equal bulk of Alcoho' and filter. FILTRATE. 1 Ca S O 4 Wash, dissolve in water, add (N H 4 ) H and (NH 4 ) 2 C 2 4 . White pre- cipitate indicates Calcium. Confirm by flame reaction. RESIDUE. Evaporate to a small bulk, dilute with water, cautiously add Fe 2 C1 6 (until a drop of the solution gives a buff coloured precipitate with (N H 4 ) H 0, neutralize with Na 2 CO 3 , add a few crystals of sodium .icetate and a few drops of acetic acid, boil for 15 minutes, and filter the hot liquid I FILTRATE. ! Wash with boiling water, boil precipitate with Na 2 C O 3 , filter, acidify the filtrate v/ith H Cl, add (N H 4 ) H O till alkaline, white precipitate in- dicates Aluminium. RESIDUE. 1 Mn, Zn, Mg (traces of Co and Ni). Add (NH 4 )HO, NH 4 C1, and (NH 4 ) 2 S. filter. FILTRATE. Zn, Mn (traces of Co and Ni). Wash, dissolve in H Cl, and proceed as in Table F. Method I. 1 Boil down to a small bulk, and add Na 2 H P O 4 , white precipitate indicates Mag* nesium. ACID, &c. Separation of Iron, Nickel, Cobalt, Aluminium, nesium, in presence of Phosphoric, Boric, and Oxalic Acids. alkaline), N H 4 Cl and (N H 4 ) 2 S. Warm the mixture gently in a flask, shake for once with water alone. Treat the precipitate with cold dilute H Cl, and filter. FILTRATE. Cr,Al,Fe,Zn,Mn,Ca,Ba,Sr,Mg,H 3 Po 4 ,H 3 BO 3 and H 2 C,,O 4 . Test small portions for Cr, Fe, H 3 P O 4 , H 3 B O 3 and H a C 3 O 4 . Adopt Method I. if Cr be absent ; Method II. if present. METHOD II. Cr is PRESENT. Bv,il down with a little K Cl O 3 till it smells of chlorine. Test for Ba, Sr, and Ca ! in a portion of the solution, as in 1. To the remainder add Fe a Clf,, as in I ; : evaporate down nearly to dryness, dilute with water, add Na 2 CO 3 , or K H O, till j just neutral or slightly acid, when perfectly cold, add Ba C O 3 in slight excess, ! place in a small flask, close with a cork, shake well, and allow to settle. Filter. RESIDUE. r FILTRATE. Fe 2 (H0) 6 , Cr 2 (HO) 6 , A1 2 (HO) 6 . Proceed as in Table F., Method II. RESIDUE. Mn, Zn (traces of Co and Ni). Proceed as in Table F., Method II. Mn, Zn, Mg (traces of Co and Ni). Add a few drops of H Cl, boil to expel CO 2 , add (NH 4 )HO (till alkaline) N H 4 Cl, and (N H 4 ) 2 S. Filter. |_ FILTRATE Boil with H 2 SO 4 to remove Ba and Sr. Filter, precipitate the Ca in the fiitiate by adding (NH 4 )HO and (NH 4 ) a C 2 O 4 . Filter. Concentrate the filtrate, add Na 7 IIPO 4 , white precipitate indicates Magnesium, N 2 i8o APPENDIX. APPENDIX D. TREATMENT OF PLATINUM RESIDUES. THESE residues generally contain the platinum as double chloride of potassium or ammonium. Any soluble platinum chloride is precipitated by addition of ammonium chloride, and the filtered precipitate is dried and ignited. The residue is then thoroughly washed with hot water, and dissolved by boiling with a mixture of three parts of hydrochloric to one of nitric acid. It is best to pour off the solution from time to time and add fresh acids. When all the platinum is dissolved the solution is evaporated nearly to dryness ; hydrochloric acid is added (to expel the nitric acid), and the liquid is evaporated to complete dryness on the water bath. The residue is dissolved in water or preserved dry in bottles. QUESTIONS AND EXERCISES. 3. (a) How many grams of oxygen can be obtained by heating 1 ,080 grams of mercuric oxide ? (b] Twenty-four grams of oxygen are required ; how much mercuric oxide must be heated ? 4-s (a) What is left when potassium chlorate is heated ? (ft) How many grams of potassium chlorate must be heated in order to yield 500 grams oxygen ? (c) How much potassium chloride will remain 'after heating 247 grams potassium chlorate ? 5. How is potassium chlorate distinguished from potassium chloride ? 6. (a) Describe the best means of preparing oxygen ? (b} The residue in a flask used to prepare oxygen, consisted of 13 grams of potassium chloride ; how much chlorate had been heated, and how much oxygen evolved ? (c} Why is manganese dioxide mixed with potassium chlorate in preparing oxygen ? 182 PRACTICAL CHEMISTRY. 8. (a) What happens when a taper is burned in oxygen ? () How is the presence of carbon dioxide generally ascertained ? 9. (a) 1*4 gram of charcoal containing 96 per cent. of carbon is burnt in oxygen ; how much carbon dioxide is produced ? (b] Contrast the burning of a candle with the breath- ing of an animal in air. 10. How much oxygen is required to burn 42 grams of sulphur to sulphur dioxide ? 11. How much phosphoric anhydride can 'be ob- tained by burning 2*34 grams phosphorus in oxygen, and how much phosphoric acid would be obtained by dissolving the anhydride in water ? 13. (a) How much hydrogen can be got by dis- solving 15 ounces of zinc in sulphuric acid ? () Which will yield the greater quantity of hydro- gen when dissolved in acids-- 162 grams zinc, or 140 grams iron ? Why ? 14. How much water will be formed by burning all the hydrogen evolved by dissolving 3*42 grams zinc in sulphuric acid ? 15. How much zinc sulphate should be got by dis- solving i Ib. zinc in sulphuric acid ? 19. How is nitrogen prepared ? 21. One hundred grams of potassium nitrate are heated with sulphuric acid ; how much nitric acid and hydrogen potassium sulphate will be obtained ? 22. What are the best tests for nitric acid ? QUESTIONS AND EXERCISES. 183 24. (a) How is ammonia gas prepared, and what is the residue from its preparation ? (b) How much ammonium chloride must be used to obtain 291 grams ammonia gas ? 25. How is ammonia tested for ? 27. (a) How is ammonium nitrate prepared ? (ff) How much ammonium nitrate can be got from 327 grams of nitric acid ? 28. How much nitrous oxide can be obtained from 437 grams of ammonium nitrate ? 29. How can nitrous oxide be distinguished from oxygen ? 30. If 7 grams of copper be acted on with nitric acid, how much nitric oxide will be obtained, and how much copper nitrate will remain ? 31. Mention some of the properties of nitric oxide. 32. (a} How is carbon dioxide prepared from marble ? (b) How much hydrochloric acid will be needed exactly to decompose a kilogram of marble ? 33. What reaction takes place when carbon dioxide and lime-water are brought together ? Give the equa- tion. 34. (a) How is carbon monoxide prepared from formic acid, and from oxalic acid ? (ft) How much carbon monoxide should be obtained from 12 grams oxalic acid? 35. What happens when a mixture of carbon monoxide and dioxide is shaken up with caustic soda ? 1 84 PRACTICAL CHEMISTRY. 36. (a) Give the equation for the preparation of chlorine from common salt, manganese dioxide, and sulphuric acid ? (b] What is the action of hydrochloric acid on manganese dioxide ? (c] How much chlorine can be got by heating 560 grams manganese dioxide with salt and sulphuric acid ? (d} How much can be got by heating the same quantity with hydrochloric acid ? 37. (a) Mention the chief properties of chlorine. (B) How does chlorine act as a bleaching agent ? 38. How much common salt is needed to prepare 109 kilograms of hydrochloric acid ? 39. (a) What is meant by " neutralizing " an acid solution ? (b] What is meant by the terms " acids," " bases," and " salts ? 40. (a] Give the equation for the preparation of bleaching powder from lime and chlorine. (ft) How much chlorine is needed to convert 5 tons of slaked lime into bleaching powder ? 41. (a) What happens when each of the following acids is brought in contact with calcium hypochlorite, sulphuric, nitric and hydrochloric ? () How can a solution of hypochlorous acid be distinguished from a solution of chlorine ? 4-2. (a) Give the equation for the preparation of iodine from potassium iodide. (b} Thirty-seven grams of iodine were obtained ; how much potassium iodide was used ? QUESTIONS AND EXERCISES. 185 43. What is the best test for free iodine ? 4-4.. (a) How is sodium iodide prepared ? (b] How much sodium iodide can be got from 20 grams of iodine ? 45. How much bromine can be got from 85 grams potassium bromide ? 46. How much sodium, bromate and bromide, should be obtained by dissolving 97 grams bromine in caustic soda ? 47. (a) How is hydrofluoric acid prepared ? (b) What is its action upon glass, and how may it be used to etch glass ? 48. (a) What is the action of strong sulphuric acid upon copper ? (b) How much sulphuric acid and copper must be used to obtain a kilogram of sulphur dioxide ? (c} What is the action of sulphur dioxide upon nitric acid, and upon potassium chromate solution ? 50. (a) Describe in detail how sulphuretted hydrogeii gas is prepared. (b) How much ferrous sulphide must be used to obtain 42 grams of the gas ? 51. (a) How does sulphuretted hydrogen serve to divide the metals into groups ? (b) How could you separate by means of sulphur- etted hydrogen a solution containing copper, iron, and sodium ? 52. How is sodium hydrate prepared ? 54. (a) What is meant by the terms " oxidizing " and " reducing " flames ? iS6 PRACTICAL CHEMISTRY. (b) How is each obtained, and for what is each used? 55. (a) How would you distinguish, by means of the blow-pipe, salts of the following metals copper, cobalt, lead, zinc, antimony, aluminium, chromium, and strontium ? (6) "What happens when mercuric chloride is heated on charcoal along with sodium bi-carbonate in the blow-pipe flame ? 56. Name the substance which produces each oi the following reactions : (a} Bead brittle, soluble in nitric acid, sulphide black. (b) No metallic bead, but incrustation of oxide on charcoal (brown). (c) Bead malleable, soluble in nitric acid ; solution gives with sulphuric acid a white precipitate. (d) No metallic bead ; after heating, moistening with cobalt chloride and re-heating gives a pink residue. (e) Red bead, soluble in nitric acid ; solution becomes bright blue on adding ammonia : borax bead blue. (/) No metallic bead: borax bead amethyst coloured in reducing flame. 57. (a) How may a Bunsen lamp flame be substituted for the flames obtained by the mouth blow-pipe ? (&) Which is the hottest part of the Bunsen lamp flame, and which part has most reducing power ? (c) How are metallic films and metallic beads obtained by Bunsen's flame reactions ? 58. How can arsenic, mercury, thallium, cadmium, QUESTIONS AND EXERCISES. 187 gold and platinum be distinguished by Bunsen's flame reactions ? 59. Give Bunsen's reactions for zinc, mercury, bis- muth, lead, copper, and iron. 60. Give Bunsen's tests for phosphorus and sulphur, and state the reactions involved in each. 61a. Six substances gave the following reactions when examined by means of Table A. Name each substance. (a) Water evolved on heating (reaction acid). Sub- stance dissolved in water, and on adding HC1 a gas was evolved smelling of burning sulphur. None of the group reagents gave a precipitate. Flame coloration, yellow. (b} Nitrogen tetroxide evolved on heating. Sub- stance dissolved in water, and gave a white precipitate with hydrochloric acid, soluble in hot water. (c} Oxygen evolved on heating. Substance insolu- ble in water, but dissolved by heating with hydro- chloric acid with evolution of chlorine. No precipi- tate with HC1, or with HC1 -f H 2 S, but on adding (NH 4 )HO + NH 4 C1 + (NH4) 2 S a flesh-coloured pre- cipitate was obtained. (d] Carbon dioxide evolved on heating strongly. Substance insoluble in water, but on adding HC1 a gas was evolved with effervescence, which rendered lime- water turbid. No precipitate with the first three group reagents, but a white one with (NH 4 ) HO + NH 4 Cl + (N H 4 ) 2 C O 3 . Flame coloration dull red. 1 88 PRACTICAL CHEMISTRY. (e) White sublimate on heating. Substance dis- solved in water and gave no precipitate with BaCl 2 , but a white one with AgNO 3 , insoluble in NHO 3 . No precipitate with any of the group reagents, but gave a smell of ammonia on heating with NaHO. (/) Oxygen evolved on heating. Substance soluble in water, gave no precipitate with the first three group reagents, but a white one with (NH 4 ) HO + NH 4 C1 + (NH^COg. Flame coloration green. 61/3. Six substances gave the following reaction when examined by means of Table B. Name each sub- stance. (a) White and unfused on heating. Yields when heated on charcoal with sodium carbonate, a mal- leable bead which leaves a black mark on paper. () Infusible and disappears on heating. Defla- grates when heated with potassium nitrate. (c) Infusible when heated. Microcosmic salt bead colourless. Blue mass after heating on charcoal, moistening with CoCl 2 and heating again. (d) White and infusible. When heated with H 2 SO 4 , evolves gas which etches glass. (e) Infusible, but darkens in colour on heating. Fused with NaHCO 3 , yields a brittle bead. (/) Fused but not volatilized on heating. Yields malleable metallic bead when heated on charcoal with NaHCO 3 . 62. (a) Give the four group reagents, and state the metals precipitated by each. (b) What are the metals of Group V. ? QUESTIONS AND EXERCISES. 189 63. Write out the tests for silver. 64. Write out the tests for mercurous salts. 65. 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