IRLF 
 
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 SB 35 M33 
 
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REESE LIBRARY 
 
 ,j\ n_~ A Ji * i 
 
 UNIVERSITY OF CALIFORNIA. 
 ?iveJ iHrvrs \fT 
 
EXPERIMENTS 
 
 ARRANGED FOR 
 
 STUDENTS IN GENERAL CHEMISTRY 
 
 BY 
 
 EDGAR F. SMITH AND HARRY F. KELLER, 
 it 
 
 PROFESSOR OF CHEMISTRY, UNIVERSITY OF PROFESSOR OF CHEMISTRY, MICHIGAN MINING 
 
 PENNSYLVANIA, PHILADELPHIA. SCHOOL, HOUGHTON. 
 
 SECOND EDITION, ENLARGED, WITH 37 ILLUSTRATIONS. 
 
 PHILADELPHIA: 
 
 P. BLAKISTON, SON & CO 
 
 1012 WALNUT STREET. 
 
|)Q$ 
 
 COPYRIGHT, 1891, BY P. BLAKISTON, SON & Co. 
 
 PRESS or WM F. FELL & Co., 
 I22O-24 SANSOM ST., 
 
 PHILADELPHIA. 
 
PREFACE. 
 
 This little work is designed as a guide for beginners in chemistry. 
 The arrangement of the course is such as the authors have used with 
 success in the instruction of their classes ; its object is not to dispense 
 with the supervision of an instructor, but rather to assist him. Although 
 reference is made to Richter's "Inorganic Chemistry," any other text- 
 book on the subject can be employed in its stead. The experiments 
 have been collected from various sources, and no claim is made for 
 originality. 
 
CONTENTS. 
 
 PART I. 
 
 CHAPTER PAGE 
 
 I. APPARATUS, MANIPULATIONS AND OPERATIONS, 9-10 
 
 II. GENERAL PRINCIPLES 10-11 
 
 III. HYDROGEN, 12-14 
 
 IV. CHLORINE, BROMINE, IODINE, FLUORINE, 14-19 
 
 V. OXYGEN, SULPHUR, . . 19-25 
 
 VI. NITROGEN, PHOSPHORUS, ARSENIC, ANTIMONY, ." 25-32 
 
 VII. CARBON AND SILICON, BORON, 32-34 
 
 PART II. 
 
 VIII. POTASSIUM, SODIUM [AMMONIUM], 35-38 
 
 IX. CALCIUM, STRONTIUM, BARIUM, 39-41 
 
 X. MAGNESIUM, ZINC, 41-42 
 
 XL MERCURY, COPPER, SILVER, GOLD, 43-46 
 
 XII. ALUMINIUM, TIN, LEAD, BISMUTH, 46-50 
 
 XIII. CHROMIUM, MANGANESE, IRON, NICKEL, COBALT, 50-56 
 
NON-METALS. 
 
 FIG. i. 
 
 CHAPTER I. 
 
 APPARATUS, MANIPULATIONS AND OPERATIONS. 
 
 1 i ) The Bunsen burner and the blowpipe. 
 
 i. Make a borax bead. 2. Dissolve a very minute quantity of man- 
 ganese dioxide in it. 3. Heat in the oxidizing flame (?). 4. In the 
 reducing flame (?). 5. Heat oxide of lead on charcoal in the reducing 
 flame. 6. In the oxidizing flame. 
 
 (2) Working with glass tubing and rods. 
 
 i. Cut various lengths of rods and tubing. 2. Round the sharp edges 
 by softening and turning the ends in the lamp. 
 
 (3) Construct a wash-bottle (Fig. i). 
 
 i. Soften a sound cork by rolling it under your foot on 
 a clean floor. 2. Bore two parallel holes through it by 
 means of a cork-borer. These perforations should be cyl- 
 indrical and of less diameter than the glass tubes they are 
 to receive. Use a rat-tail file in enlarging them. 3. Cut 
 suitable lengths of glass tubing. 4. Draw the longer one 
 to a fine point after softening in the flame. 5. Bend the 
 tubes in an ordinary fish-tail burner, and round the sharp 
 edges. 6. Fit the different pieces together. 
 
 (4) Arrange some other form of apparatus for practice. 
 
 (5) The balance. 
 
 i. Weigh an object by placing it on the left-hand pan of the balance, 
 and a weight judged about equal on the right-hand pan. Should the 
 latter be found too heavy, replace it by the next smaller one \ if too light, 
 by the next heavier one. Then add systematically the smaller weights, 
 until the needle points to the middle of the scale. The final adjustment 
 is made with the rider. In adding or removing weights, the supports 
 must always be raised. 
 
 (6) Measuring vessels. 
 
 i. Measure off 10 cc. of water (a] in a cylinder, () in a burette, (c) 
 
 9 
 
10 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 in a pipette. Always read the lower meniscus. 2. Measure off similarly 
 50, 100 and 200 cc. of water, and determine their weight. 3. Measure 
 the volume of 50 grms. of oil of vitriol, and of 65 grms. of muriatic acid. 
 What are the specific gravities of these substances ? Note the relation 
 between weight and volume in the metric system. 
 
 (7) Chemical operations : Solution, evaporation, crystallization, precip- 
 itation, filtration, washing and drying. 
 
 i. Place into a test-tube pure sodium carbonate, into another cobalt 
 chloride, and add distilled water to each. Stir. What occurs ? 2. To 
 calcium carbonate, add water. Is there any change ? Now add a little 
 hydrochloric acid. What action has it ? 3. Pour 5 cc. of strong hydro- 
 chloric acid upon powdered manganese dioxide ; observe appearance and 
 odor. Note, too, in each case, whether heat has any effect. Distinguish 
 between chemical and mechanical solution. 4. Heat the cobalt chloride 
 and the calcium carbonate solutions, each in a separate dish, on an iron 
 plate, until the liquids are completely driven off (?). 5. Dissolve potas- 
 sium chlorate in hot water, and allow to stand and cool (?). 6. To a 
 portion of the cobalt chloride solution, add a solution of soda; boil. 7. 
 Allow to settle and filter. 8. Wash the precipitate "until pure water runs 
 through the filter (?). 9. Heat the filter until perfectly dry. 
 
 CHAPTER II. 
 
 GENERAL PRINCIPLES. 
 
 (i) Changes in matter. 
 
 i. Rub a glass rod with a piece of cloth, then touch particles of paper 
 with it (?). 2. Through an insulated spiral of stout copper wire pass a 
 current from two Bunsen cells. Place a piece of wrought-iron a nail 
 will answer inside the spiral, and bring iron filings in contact with it. 
 What happens? Interrupt the current and note the result ; repeat. 3. 
 Heat a platinum wire in the non-luminous flame ; is there any change ? 
 What is the effect of removing it ? 
 
 Are the original properties of the substances in the above experiments 
 altered, after the action of the forces of electricity, magnetism and heat 
 has been stopped ? 
 
 4. Mix intimately four parts, by weight, of finely powdered sulphur 
 with seven parts of very finely divided iron (filings). Pass a magnet over 
 a portion of the mixture. Another portion treat with carbon disulphide in 
 
GENERAL PRINCIPLES. II 
 
 a test-tube. Then heat the remaining portion in a tube over a gas flame. 
 Note carefully what occurs in each case. Powder the mass FIG a 
 resulting from the last operation in a dry mortar. Can you 
 extract from it any iron with a magnet, or dissolve out any 
 sulphur with carbon disulphide? What inference do you draw 
 from the facts observed ? 5. Decompose water in Hofmann's 
 apparatus by an electric current. The water should be acid- 
 ulated with sulphuric acid to make it a conductor of elec- 
 tricity. A current from four to six Bunsen cells is required. 
 To the gas, of which a larger volume has collected, apply a flame, and to 
 the other a glowing spark (?) 6. Heat oxide of mercury in a tube of 
 hard glass (Fig. 2). Apply the spark test (?). 7. Rub some sulphur 
 and mercury together in a mortar (?) 8. Heat sugar in a dry test-tube, 
 at first gently, and then more strongly. Note color and odor. 9. Mix 
 dry soda and tartaric acid in a mortar. Is there any action ? What 
 occurs when you add water ? Point out in what respect the changes 
 involved in experiments 1-3 differ essentially from those in 4-9. By 
 what general names can you distinguish the two different kinds? With 
 which does chemisty concern itself? Define chemistry. 
 
 Through what agencies have the results been obtained in experiments 
 4-9 ? Has any gain or loss of matter occurred in any of them ? 
 
 (2) The products resulting from 5 and 6 cannot be further simplified, 
 *. e., decomposed into dissimilar substances. They are elements* What 
 are water and red oxide of mercury ? 
 
 i. Dissolve in a little nitric acid, the black powder obtained by heating 
 an intimate mixture of powdered sulphur and copper filings. f 
 
 Evaporate the solution nearly to dryness, take up in water FIG. 3 . 
 and filter. What remains on the filter ? Place the filtrate in 
 a beaker, dip the platinum electrodes of a battery (one or two ~*" P 
 Bunsen cells) into it (Fig. 3), and allow the current to act for 
 ten minutes. What do you observe upon the platinum foil, 
 forming the negative pole ? What changes have the copper 
 and the sulphur undergone in this experiment ? 
 
 (Study pp. 18-27, i n Richter's Chemistry.) 
 
 (3) Metals and non-metals. (See Richter, p. 20.) 
 
 * The instructor should here develop the idea of element more fully. 
 
 f A better substitute would be finely divided copper ; such as may be obtained by the 
 reduction of black cupric oxide in a current of hydrogen gas (see page 14). 
 
12 
 
 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 CHAPTER III. 
 
 HYDROGEN. H. 
 
 (i) Put several pieces of granulated zinc into a test-tube and pour 
 dilute sulphuric acid upon them. What occurs? 
 
 FlG . 4 . (2) Arrange the apparatus shown in Fig. 4. 
 
 The flask should contain about 15 grms. of Zn, 
 and dilute H 2 SO 4 is poured through the funnel 
 tube. When all the air in the apparatus has been 
 displaced (ask for precautions /) collect six test- 
 tubes full of the gas over water. 
 
 (3) What are its properties ? Will it burn ? 
 Support combustion? Is it lighter than air? 
 (4) i. To learn what becomes of hydrogen when it burns in air, 
 arrange apparatus as in Fig. 5. The gas is led from the evolution flask 
 A, into a bottle containing concentrated H 2 SO 4 , and then passes through 
 a tube filled with pieces of CaCl 2 . The gas which escapes is free from 
 moisture. Burn it under a cold glass jar. What do you obtain ? 2. Fill a 
 small flask with a mixture of one vol. of H and five vols. of air ; cork ; 
 invert the flask several times to mix the gases; wrap a towel around it and 
 bring its mouth to a flame. Result ? 
 
 FIG. 5 . (5) Hydrogen is not the only product of 
 
 the action of Jf 2 SO 4 itpo?i Zn. 
 
 Pour some of the liquid remaining in 
 the flask, in which H was generated, into 
 a porcelain dish. Evaporate to about 
 one-third of the original bulk ; allow to 
 stand several hours. You will now dis- 
 cover that the solution is full of colorless 
 crystals. These are zinc sulphate or white vitriol a salt, ZnSO 4 -j- 7H 2 O. 
 Write the equation of the reaction. 
 
 (6) Determine the weight of H generated by a given weight of Zn. 
 
 A piece of Zn (not more than .02 grm.) is accurately 
 weighed, and placed under a funnel in a beaker (Fig. 6). 
 The latter is then nearly filled with water, so that the en- 
 tire funnel is under the surface. A test-tube containing dilute 
 H,SO 4 is lowered over the stem of the funnel. Hydrogen 
 appears and collects in the tube. When all the Zn has dis- 
 appeared,* transfer the tube containing the H to a larger vessel, 
 holding water. Measure the volume of the gas by marking 
 
 FIG. 6. 
 
 This may be hastened by bringing a spiral of platinum wire in contact with the Zn. 
 
HYDROGEN. 1 3 
 
 the tube where the inner and outer levels of water are even, and then 
 weighing or measuring the quantity of water that it will hold to that 
 mark. Note the temperature of the water, and the height of the baro- 
 meter. 
 
 The weight of the H is found by multiplying the vol. by the wt. of 
 i cc., /.<?., .0000896 gr. Before this can be done, however, it is necessary 
 to reduce the volume of the gas to o C. and 760 mm., as the above value 
 has been determined under these conditions. If v =r volume observed, 
 t temperature, and p = pressure, then 
 
 v X p 
 
 (i + at) X 
 
 and W 
 
 X .000x5896 . 
 
 a = .003665.* 
 
 To calculate the quantity of Zn necessary to generate a 
 unit of H, we say 
 
 FIG. 7. 
 
 Wt of H : Wt of Zn : : I : x. 
 
 x here stands for the equivalent weight of Zn. 
 
 The equivalent weights of some other metals, such as 
 Fe, Cd, Mg, can be determined in the same manner. Mag- 
 nesium gives the most satisfactory results. 
 
 (7) Decompose water by electrolysis and test the products. 
 
 (8) Take a small piece of sodium, wrap it in paper and place it, with 
 forceps, under the mouth of a test-tube 
 
 filled with water, and inverted in water 
 (Fig. 7) contained in a dish. Repeat this 
 until the test-tube is filled with the gas. 
 Test it for H. 
 
 What becomes of the metal ? Write n 
 the reaction. 
 
 (9) Construct the apparatus shown in 
 Fig. 8. 
 
 Water is heated to boiling in the flask 
 
 A, and the steam led over iron filings or wire, heated to redness in a tube 
 of hard glass. Care must be taken to prevent steam from condensing in 
 any part of the tube. Collect the escaping gas over water. Test it 
 for H. 
 
 * Tension of aqueous vapor is here neglected. 
 
EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 Is the iron changed ? Equation ? 
 
 FIG (10) Into a weighed tube 
 
 of hard glass (6-8 inches in 
 length) place a weighed quan- 
 tity (1-2 grams) of cupric 
 oxide ; connect the tube with 
 a CaCl 2 -tube of known weight 
 (Fig. 9). Pass a current of dry 
 H over the CuO and heat. 
 After the change is complete, cool and determine the loss in weight of 
 tube -)- CuO, and the gain in the CaCL 2 tube. Explain the reaction. 
 
 Problems. i. How much H can be obtained from Zn and 299 grms. 
 of H 2 SO 4 ? 2. How much Zn and H-jSO^ are necessary to furnish 100 
 grms. of H ? 3. Suppose you have found that .015 grm. of Mg yield 
 15.2 cc. of H at 20 C. and 750 mm., what is the equivalent weight of 
 that metal ? 4. How many cc. of H can be obtained from 2 grms. of 
 Na and water? 5. How many grms. of H 2 O can be decomposed by 
 5 grms. of Fe; by how much is the weight of the latter increased ? 6. 
 10 grms. of CuO will yield how much Cu upon heating in H? 
 Give a brief summary of what you have learned about H. 
 (Study Richter, pp. 39-47-) 
 
 FIG. 
 
 CHAPTER IV. 
 
 FIRST NATURAL GROUP OF ELEMENTS CHLORINE, BROMINE, IODINE, 
 
 FLUORINE. 
 
 CHLORINE. Cl. 
 
 (i) Into a test-tube put MnO 2 and concentrated HC1. Note what 
 happens both before and after heating. 
 
 (2) Use apparatus (shown in Fig. 10) for preparing 
 larger quantities of Cl. The MnO 2 should be in the 
 form of small lumps (not powder.) Heat the mix- 
 ture gently, pass the Cl through a small quantity of 
 water and collect it either by downward displace- 
 ment or over warm water. 
 
 Write the reaction. How many atoms of Cl are 
 liberated ? How many molecules ? 
 
 (3) i. What is the normal condition of this ele- 
 ment? 2. Is it lighter than air? 3. Is it inflamma- 
 
 4. Does it support combustion? 
 
 ble? 
 
FIRST NATURAL GROUP OF ELEMENTS CHLORINE. 15 
 
 To obtain answers to these questions, fill 5 test-tubes with dry Cl, and 
 proceed as under H. 
 
 (4) Again fill 5 bottles with the dry gas. Cover them with glass plates. 
 Into i throw a little pulverized antimony. 
 
 Into 2 carefully introduce a piece of phosphorus. 
 
 Into 3 insert tissue paper moistened with oil of turpentine. * 
 
 Into 4 introduce colored flowers. 
 
 Into 5 pour a little litmus solution. 
 
 What are the results ? 
 
 (5) Fill a small-sized flask one-half with chlorine, the other half with 
 hydrogen. Wrap a towel about the flask and apply a flame to its open 
 mouth (?). Care. 
 
 (6) Invert a bottle filled with Cl over water saturated with the same 
 gas. What follows in the course of a few hours' exposure to sunlight ? 
 Can you account for results in experiments (4), (5) and (6) ? Why should 
 the gas be collected over warm water ? 
 
 (7) Determine the weight of a litre of chlorine. Arrange apparatus 
 as shown in Fig. n. 
 
 In the evolution-flask place a mixture of equal 
 weights of salt and manganese dioxide. Add sul- 
 phuric acid, previously diluted with its own volume 
 of water (pour the acid into the water ! ). Heat 
 gently. Chlorine is evolved, and dried by passing 
 it through concentrated sulphuric acid, after which it 
 is led into the perfectly dry flask c. When this is 
 filled, which you ascertain by the color of the gas in 
 the neck, slowly withdraw the tube and cork the 
 flask at once. Weigh the flask. Read the barometer and thermo- 
 meter. Determine, also, the weights of the flask filled with air and with 
 water. 
 
 Calculation : 
 
 Capacity of flask, a 
 
 Temperature, t 
 
 Pressure, p 
 
 Flask filled with air, . w 
 
 Cl, \ . w ' 
 
 Wt. of a litre of air, *- 2 93 grm. 
 
 " Cl, ....'.'..... x 
 
 * It is well, when the turpentine is old, to gently warm it, and then saturate the tissue 
 paper. 
 
1 6 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 The weight of the air filling the flask is a xP x - OOI2 93 ^ The 
 
 (i + .00367 t) 760 
 
 ence between this and w is the weight of the vacuous flask. Subtract 
 this from w'. The remainder is the weight of the Cl, (W). Reduce the vol. 
 
 of the Cl to o C. and 760 mm. (see under H) ; it is v = ax p 
 
 ( i + .003671)760 
 
 11 i ^ r T^ W X IOOO 
 
 and the weight of i litre, x = - 
 
 How much heavier is one litre of Cl than an equal vol. of H ? 
 
 Write the reaction involved in the above method for preparing Cl. 
 
 Problems. i. How many litres of Cl can be obtained from i kilo of 
 MnO 2 and HC1? 2. What weight of salt is required to prepare 100 
 litres of Cl ? 3. How many pounds of sodium sulphate and manganese 
 sulphate will be formed in the preparation of 100 litres of chlorine gas? 
 4. Calculate the number of grams of Cl that 2 litres of water will ab- 
 sorb, provided the latter takes up twice its volume of Cl ? Write out 
 your deductions from the above experiments on Cl. 
 
 (Read Richter, pp. 49-52.) 
 
 HYDROGEN CHLORIDE. HC1. 
 
 (1) Repeat the explosion of equal vols. of Cl and H. Quickly cover 
 the mouth of the flask, and immerse it under water. Does the latter rise ? 
 Put a drop of the liquid on the tongue and note the taste. Add some 
 blue litmus solution. Is there any change? 
 
 (2) The product of the union of H and Cl is a colorless gas. It is 
 called hydrogen chloride. It is usually prepared by the action of sulphuric 
 acid upon salt, thus: 
 
 2NaCl + H 2 S0 4 = 2HC1 + Na 2 SO 4 . 
 or, better, NaCl + H 2 SO 4 = HC1 -f NaHSO 4 . 
 
 The apparatus employed here is the same as that used for making Cl 
 (Fig. 10). 
 
 (3) Determine the properties of HCl as under H and CL 
 
 What new property appears here ? Fill a long dry glass tube with the 
 gas, and quickly bring it into a basin containing water colored blue with 
 litmus. What happens? What does HCl gas yield on dissolving in 
 water ? 
 
 (4) In the preparation of H by the action of Na upon water, it was 
 observed that the liquid became soapy to the touch, and acquired the 
 property of turning red litmus blue. Prepare such a solution. To it add 
 a few drops of litmus, and then an HCl solution (gradually) from a 
 
FIG. 12. 
 
 FIG. 13. 
 
 FIRST NATURAL GROUP OF ELEMENTS HYDROGEN CHLORIDE. 17 
 
 burette, until the blue color just begins to turn. Evaporate the resulting 
 
 liquid to crystallization. Dissolve and recrystallize the product. It 
 
 appears in cubes, and has the taste of common salt. It 
 
 does not affect either red or blue litmus. We say it is 
 
 neutral in reaction. The substance is chloride of sodium 
 
 or common salt. What is a salt? An acid? A base ? 
 
 How can you obtain HC1 and Cl from NaCl ? 
 
 (5) Burn Hin an atmosphere of Cl, and Clin hydrogen. 
 Generate chlorine as already described (p. 14) and 
 
 collect it in a large cylinder. Into this introduce a burn- 
 ing jet of hydrogen (Fig. 12). Does it continue burn- 
 ing ? What is the appearance of the flame ? To show the combustion of 
 Cl in H arrange apparatus as in Fig. 13. 
 
 (6) To determine the weight of a litre of HCl, proceed 
 exactly as under chlorine. 
 
 (7) Determine the composition of HCl by volume. 
 
 i. Fill a perfectly dry and graduated tube with HCL Close 
 the open end with the thumb, and opening the tube for a 
 moment, quickly pour in about 10 cc. of sodium amalgam 
 (see sodium, p. 36). Close the tube at once with the thumb, 
 slightly moist, and shake well. Invert the tube in a large 
 beaker of water, and remove the thumb. The amalgam will 
 drop into the water, and the latter will rush up into the tube, 
 filling it nearly half full. Immerse the tube so that the water 
 in it and that in the beaker are on the same level. This is 
 done to measure the hydrogen under atmospheric pressure. 
 Read the residual volume of the gas and measure the volume 
 of the mercury. 
 
 Calculation : 
 
 Capacity of tube, a 
 
 Vol. of mercury, b 
 
 Vol. of H c 
 
 a b 
 
 (8) Add a solution of HCl to solutions of silver nitrate ; of mercurous 
 nitrate ; and of lead acetate. What do you observe in each case? Boil 
 the precipitate formed in the lead solution with water. Cool, and note 
 result. 
 
 3 
 
i8 
 
 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 FIG. 14. 
 
 BROMINE. Br. 
 
 (1) Allow a drop of bromine to fall upon a heated watch glass ; cover 
 it quickly with a beaker. What is the color of the vapor ? Dissolve one 
 drop of bromine in each of the following solvents contained in test-tubes : 
 water, alcohol, ether, carbon disulphide, and chloroform. Note the 
 relative solubilities, and the color of each solution. 
 
 (2) i. Pass Cl through an aqueous solution of potassium bromide. 
 What happens? 2. To one portion of the product add a few drops of 
 
 CS 2 , and agitate the mixture ; what is the result ? 
 3. To another portion of the solution, containing 
 free Br, add a few drops of a starch solution.* 
 Result ? 
 
 (4) Devise a method for preparing bromine 
 from KBr. 
 
 (5) Prepare hydrobromic acid. In a small flask 
 cover 2 grams of amorphous phosphorus with 4 
 grams of H 2 O, and from a funnel, provided with 
 
 a stop-cock, gradually allow 20 grams of Br to run- in.f 
 
 The gas is purified by conducting it through a U-tube, containing moist- 
 ened pieces of phosphorus and glass (Fig. 14), and led into water to 
 obtain the aqueous solution. 
 
 (6) Add aqueous HBr to solutions of AgNO 3 , HgNO 3 and Pb(NO 3 ) 2 
 do the resulting bromide precipitates differ much from the corresponding 
 chlorides ? 
 
 IODINE. I. 
 
 (1) i. Place an iodine crystal upon a warm plate, and note color of 
 vapor. 2. Test the solubility of iodine in the same solvents as were used 
 with bromine ; what are the colors of the resulting solutions? 
 
 (2) i. Pass Cl through a solution of KI. Test the resulting liquid 
 with ether, carbon disulphide and starch solution (as with Br). 2. Repeat 
 this experiment, substituting Br- water for the Cl. Avoid excess of Cl as 
 well as Br (?). 
 
 What conclusion do you draw from these experiments relative to the 
 affinity of the halogens for potassium ? 
 
 * The starch solution necessary for this purpose can be prepared as follows : One gram 
 of starch is well ground in a mortar, with very little water, to creamy consistence. It is 
 then poured into 200 cc. of boiling water. Allow to subside, decant the clear superna- 
 tant liquid and use it for the test. 
 
 f As it is rather difficult to weigh bromine upon a balance, calculate the volume corres- 
 ponding to the weight given and measure out the same in a cylinder. 
 
SECOND NATURAL GROUP OF ELEMENTS OXYGEN. 19 
 
 (3) Pass hydrogen sulphide gas (H 2 S) into 50 cc. of water, and add 
 powdered iodine till the brown color no longer disappears. Warm, filter 
 (?) and distil the filtrate. The product is what ? 
 
 How is gaseous HI prepared ? 
 
 (4) Precipitate solutions of silver nitrate (AgNO 3 ), mercurous nitrate 
 (HgNO 3 ), lead nitrate (Pb(NO 3 ) 2 ), and mercuric chloride (HgCl 2 ) with 
 KI. Note result in each case. 
 
 FLUORINE. Fl. 
 
 (1) In a lead dish (or platinum crucible) place i gram of pulverized 
 fluor spar (CaFl 2 ). Add cone. H 2 SO 4 ; cover the dish or crucible with a 
 watch-glass coated with paraffin, through which some characters have been 
 drawn with a fine point. Heat gently for a few minutes. 
 
 What do you observe on removing the paraffin ? 
 
 (2) Can you liberate Fl from a fluoride? 
 
 Problems. i. How much NaBr, H 2 SO 4 and MnO 2 are necessary to pro- 
 duce i cu. metre of Br vapor at 20 C and 745 mm. ? 2. What per cent, of 
 HI does a liquid contain, which represents a solution of 50 litres of the 
 gas in i litre of H 2 O? 3. 10 grms. of CaFl 2 will give what weight of 
 HF1 ? 4. How much salt and sulphuric acid will be required to prepare 
 6 litres of muriatic acid of sp. gr. 1.17? What volume would the HC1 
 in these six litres occupy at 735 mm. pressure and 22C? 5. What is 
 the percentage of hydrochloric acid in a solution of which 17 cc. dis- 
 solve exactly 2 grams of metallic magnesium ? What is the volume of 
 hydrogen liberated at 760 mm. and o ? 
 
 CHAPTER V. 
 
 SECOND NATURAL GROUP OF ELEMENTS OXYGEN, SULPHUR, 
 SELENIUM, TELLURIUM. 
 
 OXYGEN. O. 
 
 (1) Preparation. i. Weigh the hard glass tube a (Fig. 15), and intro- 
 duce a weighed quantity (about .5 grm.) of red oxide of mercury. Ignite 
 strongly; collect the liberated gas, and measure it. Weigh the tube 
 with the residue. What are the products of the ignition ? 
 
 (2) Prepare more of the gas, as follows : Mix equal parts of KC1O 3 
 and pulverized MnO 2 ; heat in a tube of hard glass or small retort. 
 Collect the gas in bottles over water (Fig. 15). 
 
20 
 
 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 FIG. 15. 
 
 Into No. i lower a piece of ignited sulphur on an iron spoon. Note 
 result. Add water after the combustion (?). 
 
 Into No. 2 introduce a small piece of burn- 
 ing phosphorus (care !). Proceed as in No. i. 
 
 Into No. 3 introduce ignited charcoal. Treat 
 as before. Add now a few drops of blue litmus 
 to the contents of each bottle. Any change? 
 
 Into bottle No. 4 introduce a fine watch 
 spring, previously heated at one end and 
 dipped into powdered S. Result ? 
 
 Is oxygen heavier or lighter than air? 
 it color, taste, or odor ? 
 it support combustion ? 
 
 Has 
 
 Will it burn ? Does 
 
 What other methods can be used for preparing O ? 
 (3) Determine the weight of a litre of O. 
 
 FIG. 
 
 Arrange apparatus shown in Fig. 16; a is a 
 tube of hard glass, whose weight is known ; it 
 contains a weighed amount of KC1O 3 (about 0.3 
 grm.). The bottle A is filled with water, b is a 
 clip and d a beaker. The exit tube should con- 
 tain water as far as clip b at the beginning of the 
 experiment. Open the clip, heat a to bright 
 redness, and receive the water displaced by the 
 O in d. When no more gas is evolved, cool ; allowing the rubber tube 
 to dip under the water of the jar. Some of the water will be drawn back 
 into the bottle (?). Measure the volume of the water in d. Note the 
 temperature of the air, and the height of the barometer. Weigh a, con- 
 taining residue (KC1). 
 
 Calculation : 
 
 Weight of the tube, a 
 
 Weight of KC1O 3 and tube, b 
 
 Weight of KC1O 3 , b a 
 
 Volume of H 2 O collected, v 
 
 Barometric pressure, p 
 
 Temperature, t 
 
 Aqueous tension at t, ... p 7 
 
 Weight of KC1 and tube c 
 
 V X(P P') (b c) x iooo 
 
 Vo_. 
 
 "(I + .003670x760 
 
 and 
 
 .Vo 
 
 Dissolve the residual KC1 in water, and to its solution add nitrate of 
 silver (?). How does KC1O 3 behave under like conditions? 
 
SECOND NATURAL GROUP OF ELEMENTS WATER. 21 
 
 (4) Give a summary of" your work upon O. 
 
 Problems. i. How much O, by wt. and vol., can be obtained from 
 54 grms. of HgO? 2. Heat will expel what vol. of O from 2.45 grms. 
 of KC1O 3 ? 3. How much HgO is necessary to yield i cu. d. m. of O ? 
 4. How many times is O heavier than H ? 
 
 OZONE. O 3 . 
 
 (i) Pour water on clean pieces of phosphorus to half cover them ; 
 invert a large, clean jar over this and allow to stand for several hours'. 
 Test the air under the jar for ozone. For this purpose use paper impreg- 
 nated with a mixture of starch paste and potassium iodide. What occurs ? 
 
 (Read Richter, p. 85-89:) 
 
 COMPOUNDS OF OXYGEN AND HYDROGEN. 
 
 WATER. H 2 O. 
 
 (1) Arrange the distillation apparatus (Fig. 17) and prepare about 100 
 cc. of distilled water. Note its taste and odor. Test it for chlorides 
 with AgNO 3 . Does it leave a residue FlG . I7 . 
 
 upon evaporation ? What action has 
 it on litmus? 
 
 Apply all these tests to a natural 
 water (except rain). 
 
 (2) i. Heat a little vegetable mat- 
 ter in a dry test-tube. 2. Heat fresh 
 meat in the same manner. 3. Care- 
 fully heat crystals of zinc or copper 
 
 sulphate in a test-tube. What happens in these experiments? 4. Expose 
 clear crystals of sodium phosphate, on a watch-crystal to the air. 5. Do 
 the same with pieces of calcium chloride. Results ? 
 
 (3) Determine the quantitative composition of water. 
 
 1. The composition of water by weight follows from the experiment of 
 reducing oxide of copper described under H. 
 
 2. The relative volumes with which O and H unite to form water, are 
 determined either by analysis or synthesis. The former has been per- 
 formed in electrolyzing water. 
 
 3. Fill a eudiometer (Fig. 18) with water. Through a rubber tube 
 admit about 50 cc. of O and then a like volume of H. (If the eudiome- 
 ter is not graduated, mark these with rubber bands.) Close the open 
 end with your thumb, leaving some air to serve as a cushion beneath it, 
 and pass the spark. Remove the thumb, and pour in enough water to 
 
22 
 
 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 FIG. 18. 
 
 make the levels equal in both limbs. What is the amount of the contrac- 
 tion ? What is the residual gas ? Test it. 
 
 (4) Determine the weight of a litre of steam, Con- 
 struct apparatus shown in Fig. 19. The flask a is closed 
 with a cork. C is a vessel containing melted paraffin. 
 A small glass tube is weighed and filled with water (not 
 more than .02 grm.). Heat the vessel Cwith a Bunsen 
 burner until the temperature of a is constant (?). Now 
 drop the tube containing the water through the mouth 
 of the flask (the bottom of which should be protected 
 with a layer of asbestos) and quickly re-cork. When 
 the fall of water in the graduated tube ceases, read the volume of gas, 
 and note the temperature and pressure of the air. 
 The calculation is analogous to that used under O. 
 (5) Perform experiment 2, p. 100 in Richter. 
 How many volumes of steam result from the com- 
 bination of 2 vols. of H and i vol. of O? 
 
 How would you deduce the molecular formula of 
 water from the preceding experiments ? 
 
 FIG. 19. 
 
 (1) Add moist hydrated barium peroxide to cold dilute H 2 SO 4 . Filter. 
 What does the filtrate contain ? 
 
 (2) i. Add a solution of KI, containing starch, to a portion of this 
 liquid (?). Ferrous sulphate hastens the reaction. 2. Cautiously add a 
 dilute solution of potassium permanganate to another portion (?). 
 
 COMPOUNDS OF OXYGEN AND CHLORINE. 
 
 (1) Make a dilute solution of caustic potash, and conduct chlorine into 
 it until the latter is no longer absorbed. Treat one portion of the product 
 with HC1, and another with H 2 SO 4 . What results ? 
 
 (2) Mix 10 grms. of quicklime with 25 cc of H 2 O. After the slaking 
 FIG. 20. is finished, conduct Cl into the mixture until it is no 
 
 longer absorbed. 
 
 Add HC1 to one portion and H-jSO* to a second 
 portion. 
 
 What is set free ? Does it bleach ? 
 
 (3) Pass Cl into a hot concentrated solution of KOH 
 till it ceases to be absorbed (Fig. 20). What separates 
 upon cooling? Recrystallize the product from water. Will it give off O 
 
SECOND NATURAL GROUP OF ELEMENTS SULPHUR. 23 
 
 upon heating? Try the action of HC1 upon a crystal. Allow a drop of 
 cone. H-jSOi to fall upon a small crystal and warm gently (?). Care ! 
 Observe carefully the behavior of KC1O 3 upon heating (?). 
 
 SULPHUR. S. 
 
 (i) Place a few grams of powdered S in a dry test-tube, and heat 
 gradually. Observe and describe the changes which occur. 
 
 (3) Dissolve a little S in CS 2 and allow to stand till the liquid has 
 evaporated. What remains? 
 
 (4) Determine the sp. gr. of S (Fig. 21). Water, previously boiled is 
 introduced into a flask. It is essential that the neck of the flask should 
 be narrow. Weigh the flask, then place an additional 10 grm. 
 weight upon the right-hand pan of the balance and small pieces FIG. 21. 
 of S upon the left-hand pan, until the pointer is again in the 
 middle. Now introduce the S into the flask. Carefully remove 
 
 water above the mark and re-weigh the flask with its contents. 
 The loss in weight will represent the weight of a volume of water 
 equal to that of 10 grms of S. The latter divided by the former 
 is the specific gravity of the sulphur. 
 
 (5) Prepare the monoclinic modification of S by melting about 10 grms. 
 of the ordinary variety in a covered Hessian crucible. Cool ; and as 
 soon as a solid crust has formed upon the surface, pierce it and allow the 
 still liquid portion of the contents to run out. Note the shape of the 
 crystals upon the sides of the crucible. 
 
 (6) To obtain the plastic variety, heat 10 grams of S in a test-tube 
 above 230 C., and pour the mass into cold water. 
 
 Test the solubility of the product in CS 2 . Preserve a portion of it for 
 several days. Does it change ? 
 
 (7) To a strong solution of yellow potassium sulphide, add HC1. 
 What are the properties of the separated sulphur ? 
 
 Give a brief outline of the element sulphur ; compare it with the pre- 
 viously studied elements. 
 
 SULPHUR AND HYDROGEN. 
 
 (8) Hydrogen sulphide is formed with difficulty from its elements, but 
 is readily obtained by the action of acids upon sulphides, thus : 
 
 FeS + H 2 SO 4 = FeSO 4 + H 2 S or Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. 
 
 The apparatus to be used is the same as that employed in preparing 
 hydrogen. 
 
 (9) What are the properties of H 2 S ? Is it soluble in water ? Does it 
 
24 
 
 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 burn ? What are its products of combustion ? Hold a porcelain plate in 
 the flame; what results ? Pass the gas into solutions of chromic acid, per- 
 manganic acid, and ferric chloride. What changes are observed ? How 
 do these last-named reactions show its reducing power? What happens. 
 to the aqueous solution of the gas when exposed to the air ? What action 
 has the aqueous solution upon litmus? To what class of compounds does 
 it, therefore, belong? 
 
 (10) Pass H 2 S through solutions of the following salts, viz. : CuSO 4 , 
 SbCl 3 , Pb (NO 3 ) 2 , AsCl 3 , and Zn (C 2 H 3 O 2 ) 2 . Note results carefully. 
 Can sulphides be prepared in another manner? (See Chap. II, i.) 
 (n) Determine the composition of hydrogen sulphide. 
 
 Into a bent tube of hard glass, filled with mercury, 
 (Fig. 22), introduce dry hydrogen sulphide.* Place a 
 piece of tin in the bent portion, and heat it. Is the 
 volume of the gas changed after the experiment, and 
 what becomes of the piece of tin ? Test the gas re- 
 maining in the tube. Do your results enable you to 
 deduce the molecular formula of H 2 S ? (See Richter, 
 3d ed., p. in.) Trace the similarity between H 2 S and H 2 O. Write a 
 summary of your experiments on H 2 S. 
 
 
 (12) 
 
 FIG. 23. 
 
 SULPHUR AND CHLORINE. 
 
 Sulphur Mono chloride. Prepare this compound by conducting 
 dry chlorine over molten sulphur. The product which distils 
 over is collected in a dry test-tube, kept cold by immersion 
 in ice water. 2. Redistil the product. Determine its boiling 
 point in an apparatus similar to that pictured in Fig. 23. 
 Note the color and odor of the product. Expose some of it 
 to the air on a watch-glass. Add water to another portion 
 contained in a test-tube. Note carefully what happens. 
 Write the reaction, and examine for all the products. 
 
 SULPHUR AND OXYGEN. 
 
 (13) Burn sulphur in the air. Result? Burn FeS 2 in the air. What 
 are the properties of the resulting compound ? It is sulphur dioxide SO 2 . 
 
 (14) Fit a small flask, as indicated in Fig. 24. Place copper turnings 
 in it, then add H 2 SO 4 (strong) through the funnel tube. Warm. Is the 
 product the same as that obtained in 13? Is it soluble in water? Has 
 the aqueous solution the same properties as the gas? 2. Pass some of 
 
 The instructor should perform this experiment. 
 
NITROGEN GROUP NITROGEN. 25 
 
 the gas into solutions of potassium dichroraate and potassium permanga- 
 nate acidulated with H 2 SO 4 . Repeat these experiments Fir ,. 24 
 with the aqueous solution instead of the gas. What 
 happens in each case? 3. Test the aqueous solution of 
 SO 2 with litmus. What is this solution commonly called ? 
 4. Fill a dry jar with SO 2 gas ; introduce colored flowers. 
 Note the result. 
 
 (15) What is the formula of sulphurous acid? How 
 many series of salts can it form ? How would you desig- 
 nate the different sodium salts ? Add HC1 to a solution 
 of Na 2 SO 3 . What follows? Evaporate the solution to 
 dryness and examine the residue. What is it? Write the reaction. 
 
 SULPHUR TRIOXIDE SO 3 and SULPHURIC ACID H 2 SO 4 . (Read Rich- 
 ter, p. 189). 
 
 (16) i. Prepare sulphuric acid as described in Richter, p. 191. Study 
 the product carefully. 2. Dilute a portion of it with water ; what happens ? 
 2. Test a portion of this diluted solution with litmus (?). 4. Another 
 portion neutralize with NaOH and evaporate. What is the residue ? 
 Does it contain any S? Prove this. 5. Add BaCl 2 to a third portion of 
 the solution. What is the precipitate ? Is it soluble in water or in hydro- 
 chloric acid ? 6. What is the action of strong H 2 SO 4 upon wood or 
 paper? Explain the cause of this action. 
 
 (17) How many series of salts can sulphuric acid form. Prepare 
 (NH 4 ) 2 SO 4 , Na 2 SO 4 , NaHSO 4 and CuSO 4 . (Read Richter, pp. 189-200). 
 
 CHAPTER VI. 
 
 NITROGEN GROUP NITROGEN, PHOSPHORUS, ARSENIC, ANTIMONY 
 AND BISMUTH. 
 
 NITROGEN. N. 
 
 (i) Preparation. i. In a dish swimming on water place a piece of 
 phosphorus and ignite it ; invert a beaker glass over it (Fig. 25). What 
 FIG 25. FIG 26 becomes of the P ? When the latter has ceased 
 burning, restore the level of the water, and note 
 the decrease in the volume of the air. Test the 
 residual gas with a burn ing taper. 2. Heat gently 
 in a small flask or retort a mixture of i part KNO 2 , 
 i pt. NH 4 C1, i pt. K 2 Cr 2 O 7 , and 3 pts. of H 2 O; 
 collect the gas over water. Fill five bottles with 
 this gas. 
 
 (2) Has it color, taste, odor? Does it burn or 
 4 
 
26 
 
 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 FIG. 
 
 FIG. 28. 
 
 support combustion ? Is the gas heavier than air ? Does it unite readily 
 with other elements? 
 
 (3) Determine the weight of a litre of nitrogen. A round-bottomed 
 flask is fitted, as shown in Fig. 26. Pour about 30 cc. of water into it, 
 and insert the rubber cork to the mark. Boil the 
 water, while the clip is open, until all the air has been 
 expelled from the flask. Steam should be allowed to 
 escape for about five minutes. Now close the tube 
 with the clip, and remove the flame. Cool and 
 weigh the flask. Read the temperature and baro- 
 metric pressure in the balance-room. 
 
 Connect the flask with the tube, b, of the aspirator, 
 containing N, and arranged as in Fig. 27. The 
 rubber tube, a, is made to dip under water, and the 
 clip is gradually opened, allowing N to enter the flask. Now raise the 
 vessel containing the water into which the rubber tube 
 dips, so that the water in it is at a higher level than that 
 in the aspirator. Close the clip. Disconnect the flask 
 and open the clip for a moment, to establish atmospheric 
 pressure in the flask. Weigh. The calculation is identical 
 with that given for O. 
 
 Wnat is the ratio between the weights of equal volumes 
 of N and H? 
 
 (4) Is air a chemical compound ? 
 How would you determine the weight of a litre of air ? 
 (5) i. Determination of the Oxygen in Air by the Pyrogallate Method, 
 At the atmospheric temperature and pressure 
 measure off 50 cc. of air in the Hempel burette, 
 shown in Fig. 28. Connect this at c with the 
 capillary of a Hempel's compound pipette (Fig. 
 29) containing an alkaline solution of pyrogallate 
 of potash. Open the clips and transfer the air to 
 the pipette by raising the tube, a. When this is 
 accomplished, and the capillary of the pipette is filled with water from b, 
 close the clips again. Disconnect the apparatus. Shake 
 the pipette for several minutes so as to bring gas and 
 absorbent in intimate contact. Reconnect pipette and 
 burette, and force the residual gas into the latter. 
 Restore atmospheric pressure and read the volume. 
 What does the loss represent ? 
 
 2. Explosion Method. To 40 cc. of air contained in 
 
 FIG. 29. 
 
NITROGEN GROUP AMMONIA. 27 
 
 the burette add 40 cc. of pure H. Pass this mixture into the Hempel 
 explosion pipette shown in Fig. 30. Close the stop-cock, d, and the clip, 
 c, then connect the platinum electrodes with an inductor and pass a spark 
 What takes place ? Measure the volume of the gas remaining. How 
 much of the contraction was due to O ? What is the composition of the 
 gas after the explosion ? 
 
 (Study Richter, pp. 116-125.) 
 
 NITROGEN AND HYDROGEN. 
 AMMONIA. 
 
 (6) Preparation. Heat an intimate mixture of finely powdered ammo- 
 nium chloride and caustic lime in a flask ; conduct the evolved gas 
 through a tube filled with small pieces of lime, and collect it in jars or 
 test-tubes over mercury. 
 
 What is the object of the lime in the tube? Why can you not dry 
 the gas by passing it through H 2 SO 4 or CaCl 2 ? Why should it be col- 
 lected over mercury ? 
 
 (7) Is ammonia gas combustible? Does it FIG. 3 i. FIG. 32. 
 support combustion ? i. Conduct NH 3 through 
 
 a glass tube, and insert this into a wider tube 
 filled with oxygen (Fig. 31). Apply a flame. 
 The ammonia gas will ignite and continue to 
 burn. 2. Heat concentrated ammonia water in 
 a beaker until there is an abundant disengage- 
 ment of gas, then conduct a rapid current of 
 oxygen through the liquid, and lower a glowing 
 spiral of platinum into the beaker (as in Fig. 32). 
 What happens ? 
 
 Note the odor of NH 3 (caution ?). Is it lighter than air? Soluble in 
 water ? 
 
 (8) Prepare an aqueous solution of ammonia. 
 What are its properties ? 
 
 Add red litmus to some of the solution (?), and then neutralize care- 
 fully with dilute HC1. Evaporate to dryness. Compare the product 
 with the ordinary NH 4 C1. Test it for Cl (?). Heat a little of it with 
 sodium hydroxide (?). Heat another portion on a platinum foil (?). 
 
 (9) Determine the weight of a litre of NH^ 
 
 Fill a dry flask with the gas by upward displacement, and proceed 
 exactly as under chlorine. What is the density of NH S ? 
 
 To determine the quantitative composition of NH 3 , perform experi- 
 ments T and 2 on pp. 130 and 131, in Richter. 
 
 Write out summary. (Read Richter, pp. 125-131.) 
 
28 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 NITROGEN AND THE HALOGENS. 
 
 (n) Pour a saturated alcoholic solution of iodine into strong ammonia 
 Weuer. Collect the precipitate on a filter and wash it with water. Open 
 the moist filter; tear it into small pieces and spread these on a board. 
 After they have become dry, touch them with the end of a rod (?). Ask 
 for instructions / (Read Richter, pp. 132-133.) 
 
 NITROGEN AND OXYGEN. 
 
 (12) Hyponitrous oxide N 2 O. i. Place about 5 grms. of ammonium 
 nitrate in a small retort ; add a little water, and apply heat. Collect the 
 product over warm water. 2. Test it with a glimmering chip ; 3. with 
 burning phosphorus ; 4. with burning sulphur. 5. Mix equal volumes of 
 this gas and of H, and apply a flame. What other gas does it resemble 
 in its properties? (Read Richter, pp. 212-213.) 
 
 (13) Nitric oxide NO. i. Pour dilute HN0 3 (sp. gr. 1.2) upon 
 copper turnings contained in an evolution flask. Cool, and allow the 
 red fumes, which form at first, to escape ; then collect the colorless pro- 
 duct over water. 2. What occurs when this gas comes in contact with 
 the air? Is it the O or the N of the air that acts upon the gas? 3. 
 Apply the tests given under (12) to this gas (?). How can NO be dis- 
 tinguished from oxygen ? 4. Fill a cylinder with NO, and add a few 
 drops of CS 2 , shake well and bring a flame to the mouth of the vessel (?). 
 5. Pass a current of NO into a strong solution of ferrous sulphate. What 
 occurs? After the solution has become saturated with the gas heat it to 
 boiling (?). 6. Pass the gas into a solution of potassium permanganate (?). 
 
 (14) Nitrogen trioxide N 2 O 8 . (Read Richter, pp. 205-206.) 
 Nitrous Add HNO 2 . (Richter, p. 206.) 
 
 (15) Nitrogen tetroxide, N 2 O 4 , and dioxide, NO 2 . i. Heat 10 grams 
 of dry lead nitrate in a test tube ; condense the escaping vapors in a 
 well-cooled receiver. What are the vapors, and what is the condensed 
 liquid ? Note the color. 2. What is the action of cold water, and of 
 aqueous solutions of the alkalies upon N 2 O 4 ? What do these reactions 
 indicate in respect to the composition of this compound ? (Richter, pp. 
 207-208.) 3. What is its action upon potassium iodide? 
 
 (16) Nitrogen pentoxide, N 2 O 5 . (Richter, p. 205.) 
 
 NITRIC ACID. HNO 3 . 
 
 i. Preparation. In a retort heat a mixture of sodium nitrate and sul- 
 phuric acid in proportions corresponding to the equation (?) : 
 NaNO, + H 2 SO 4 = N*aHSO 4 -f HNO 3 . 
 
NITROGEN GROUP PHOSPHORUS AND HYDROGEN. 29 
 
 Collect the product in a cold receiver. 
 
 2. What are the physical properties of HNO 3 ? Color ? Odor ? 
 Action on litmus (dilute with H 2 O) ? 3. What action has it on indigo ? 
 Upon the skin ? 4. Notice the effect of the acid upon the following 
 metals: Cu, Fe, Pb, Zn, Sn. Write the reaction for each one. 5. 
 Cover powdered sulphur with the acid, and warm (?). Dilute with water, 
 filter, and test the liquid with BaCl 2 (?). 6. Add a few drops of HNO 3 to 
 a solution of ferrous sulphate (?) ; warm the solution (?). 
 
 Problems. i. Required i cu. m. of N. How much air is to be de- 
 prived of O ; and how much P must be burned, if 62 pts. of P unite 
 with 80 pts. of O ? 
 
 2. How much HNO 3 , containing 46 per cent, of water, may be ob- 
 tained from 1,700 grms. of NaNO 3 , and how much water must be taken ? 
 
 3. How many grams of NH 3 will be absorbed by 5 litres of H 2 O, if 
 the latter absorbs 500 times its volume of the gas ? 4. Ten litres of 
 water having absorbed 700 times their volume of ammonia, what are the 
 least amounts of NH 4 C1 and CaO necessary for producing this solution ? 
 
 * PHOSPHORUS. P. 
 
 (1) i. Determine the physical properties of the active and the red 
 varieties. 2. Allow a small piece of the active variety to ignite in the air. 
 Will the red variety do this? 3. Throw a small piece of the yellow 
 variety into a jar of dry Cl (?). Repeat with the red variety (?). 
 4. Bring a small dry piece of active P in contact with iodine (?). 5. 
 Heat a flask containing a small piece of P and water until the former is 
 melted, then pass a current of oxygen through a delivery tube into the 
 melted phosphorus (?). Care ! (Study Richter, pp. 133-136.) 
 
 PHOSPHORUS AND HYDROGEN. 
 
 (2) Phosphine PH 3 . i. Fill a flask almost full with a moderately 
 concentrated NaOH solution. Add a few pieces of P, and heat carefully. 
 When the air in the neck of the flask has been expelled by the escaping 
 gas, insert a cork with a delivery tube the other end of which dips under 
 warm water. What becomes of the gas as it escapes into the air ? Write 
 the reaction involved. 
 
 (Richter, pp. 136-139.) 
 
 Is there any similarity between PH 3 and NH 3 ? 
 
30 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 PHOSPHORUS AND THE HALOGENS. 
 
 (3) i. Pass a current of dry CO 2 gas into a retort, the bottom of which 
 is covered with dry sand. When all the air has been expelled, introduce 
 some well-dried pieces of P, and replace the CO 2 by a stream of dry Cl. 
 Connect the neck of the retort with a Liebig's condenser, and collect 
 the product in a receiver. It is phosphorus trichloride. What are its 
 properties? Pour some of it into water (?). 
 
 2. Place a little PC1 3 in a dry test-tube, and pass a stream of dry Cl 
 upon its surface. What is the result ? 
 
 PHOSPHORUS AND OXYGEN. 
 
 (Richter, pp. 214-219.) 
 
 (4) i. Prepare phosphorus pentoxifte, P 2 O 5 , by burning a carefully dried 
 piece of P under a dry bell -jar. 2. Drop a portion of the product into 
 water (?). 
 
 (5) Orthophosphoric acid, H 3 PO 4 ; metaphosphoric acid, HPO 3 ; and 
 pyrophosphoric acid, H 4 P 2 O 7 . How are these acids obtained ? How many 
 series of salts are derived from them ? By what names would y<Ju distin- 
 guish the different salts ? 
 
 i. Dissolve some Na 2 HPO 4 in water and test che solution with AgNO 3 , 
 and FeCl 3 . What do you observe in each case ? 2. Dissolve fused 
 Na 2 HPO 4 in water, and perform the same tests with its solution. 3. Heat 
 salt of phosphorus (NaNH 4 HPO 4 ) until it no longer effervesces ; cool, 
 crush the residue in a mortar, and dissolve it in water. How does this 
 solution behave upon treating with the above reagents? 4. Acidify a 
 portion of the last-named solution with acetic acid, and add a solution of 
 albumen to it. Result? 
 
 (6) Phosphorus trioxide P 2 O 3 , and phosphorous acid H 3 PO 3 . 
 
 Pour PC1 3 into water. Evaporate the solution to syrupy consistency (?). 
 (Study Richter, p. 216.) 
 
 (7) Hypophosphorous acid H 3 PO 2 . 
 
 Heat pieces of phosphorus in a porcelain dish with a moderately 
 strong baryta solution (see p. 29). When- no more PH 3 is formed, cool, 
 filter, and pass CO 2 into the solution until it shows a neutral reaction to 
 litmus. Toward the end, the solution should be warmed. Filter and 
 evaporate to suitable concentration. Hypophosphite of barium will 
 crystallize. 
 
 How may the free acid be obtained from this salt ? 
 
NITROGEN GROUP ANTIMONY. 31 
 
 ARSENIC. As. 
 
 (1) Study the physical and chemical properties of this element. 
 (Richter, pp. 142 and 143.) Are they analogous to those of phosphorus ? 
 
 i. In a tube of hard glass heat a small piece of As to redness. Result ? 
 2. Heat As with the oxidizing flame upon charcoal (?). 3. Dissolve 
 powdered As in strong HNO 3 (?). 
 
 ARSENIC AND HYDROGEN. * 
 
 (2) Perform Marsh' s test for As* 
 Arrange the apparatus shown in Fig. ^ 
 
 33. To the mixture of Zn and dilute 
 H 2 SO 4 contained in a, add a small por- 
 tion of the solution to be tested for As. 
 The liberated gas contains H and AsH 3 
 (arsine). It is passed through c, filled 
 with CaCl 2 (?), and then through d, a 
 tube of hard glass, contracted at several places. After all the air has 
 been expelled from the apparatus, ignite the hydrogen. If As is present 
 it will burn with a bluish white flame, and white vapors will be given off. 
 Hold a cold porcelain plate in the flame (?). Heat the tube d, as shown 
 in the figure (?). 
 
 Great care must be exercised in performing this test, as the arsine gas is 
 extremely poisonous ! 
 
 ANTIMONY. Sb. 
 
 (i) Study this element in the same manner as As. Distinguish between 
 SbH 3 and AsH 3 . 
 
 i. Treat the metallic mirrors obtained in Marsh's apparatus, with a 
 freshly prepared solution of hypochlorite of sodium : As dissolves readily, 
 while Sb is scarcely acted upon. 2. Heat a piece of the tube in which a 
 mirror has formed, in the flame of the Bunsen burner. Dissolve the pro- 
 duct in dilute, warm HCI, and add H 2 S water (?). 3. Treat the spot 
 formed upon a cold porcelain plate with yellow ammonium sulphide, and 
 evaporate the solution at a gentle heat (?). 
 
 Problems. (i) How much P can be obtained from 250 grms. of 
 bones? (See Richter, p. 134.) (2) 10 grms. of P give what vol. phos- 
 phine ? (3) What is the weight of the product remaining, after evapor- 
 ating a solution of 10 grms. of As in HNO 3 ? 
 
 * Ask for instructions. 
 
32 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 CHAPTER VI. 
 
 CARBON GROUP-CARBON AND SILICON. 
 CARBON. C. 
 
 (i) How many allotropic modifications of this element are known? 
 What are their principal properties? i. Boil a dilute litmus solution 
 with powdered animal charcoal; filter. Result? 2. Substitute indigo 
 for the litmus in the preceding experiment (?). 
 3. Determination of the composition of coal. 
 
 i. Volatile matter and coke. Weigh out 2 grms. of powdered coal in 
 a platinum crucible provided with a well-fitting 
 cover. Heat with a large flame, until the escaping 
 gases cease to burn between the lid and the 
 crucible. A blast lamp flame is applied for a 
 minute longer. Cool and weigh. Loss in weight 
 represents the volatile matter. The residue is 
 called coke. , 
 
 2. Ash. A second portion of coat (i grm.) is 
 gently heated over the Bunsen flame, until the 
 volatile constituents are expelled. The heat is 
 then raised and the lid of the crucible placed in the position indicated in 
 Fig. 34. The residue is the ash. 
 (Read Richter, pp. 150-152.) 
 
 CARBON AND HYDROGEN. 
 
 (2) Methane (Marsh gas) CH 4 . 
 
 i. Preparation. Heat a dried mixture of sodium acetate and sodium 
 hydroxide in an iron tube.* Collect the gas over water. Note its color, 
 odor and taste. Does it burn? 2. Mix i vol. of it with 7 to 8 
 times its vol. of air and explode by applying a flame. (Ask for in- 
 structions !) 
 
 How would you determine the molecular weight of this compound ? 
 
 (3) Make a eudiometric combustion of i vol. of CH 4 with 2 vols. of 
 O as described in Richter, p. 121. 
 
 (4) Ethane C 2 H 6 . (Richter, p. 153.) 
 
 (5) Acetylene C 2 H 2 . Light a Bunsen burner at the base and turn it 
 
 *A hard glass tube will answer. 
 
CARBON GROUP SILICON. 33 
 
 down, so that the flame is small. Acetylene can be recognized, among 
 the products of combustion, by its characteristic odor. 
 
 (6) CARBON AND THE HALOGENS. (Richter, p. 160.) 
 
 CARBON AND OXYGEN. 
 
 ( 7 ) Carbon , dioxide CQ 2 . 
 
 i. Preparation. Upon pieces of marble, contained in an evolution 
 flask, pour dilute HC1 (i HC1 : 1-2 H 2 O). Conduct the resulting gas 
 through water and through cone. H.jSO 4 . It may be collected either by 
 downward displacement of the air, or over mercury. 2. Note color, 
 taste and odor of this gas. Is it soluble in water ? How does its weight 
 compare with that of air ? Does it burn or support combustion ? 
 3. Conduct a current of CO 2 into a solution of NaOH, evaporate the 
 liquid, and test the residue for Na 2 CO 3 (?). 4. To different portions ,of 
 Na 2 CO 3 solution, add solutions of MgSO 4i BaCl 2> Pb (NO 3 ) 2 , ZnSO 4 (?). 
 
 (Study Richter, pp. 228-232.) 
 
 (8) Carbon monoxide CO. 
 
 Preparation. i. In a tube of hard glass heat zinc dust to faint redness, 
 while conducting a slow current of CO 2 over it. In what respect does 
 the product differ from CO 2 . 2. Heat crystals of oxalic acid with 
 cone. H 2 SO 4 in a flask, and wash the product with a NaOH solution. 
 Write the reaction. Study the properties of this gas. (Richter, p. 233.) 
 
 (9) Carbon disulphide CS 2 . 
 
 Perform some of the experiments indicated in Richter, p. 234. 
 
 (10) CARBON AND NITROGEN. 
 
 i. In a dry test-tube heat a nitrogenous carbon compound with a small 
 piece of K. Cool and add water. KCN is formed and can be tested 
 with AgNO 3 . 2. Convert a portion of the KCN into KCNS by evapo- 
 rating with (NH 4 ) 2 S. Test with FeCl 3 . 3. To a solution of FeSO 4 add 
 potassium ferrocyanide. What results ? 4. What is the action of the 
 ferrocyanide upon solutions of ferric salts ? 
 
 (n) Study the nature of flame. Make the experiments described'in 
 Richter, pp. 155-160. 
 
 SILICON. Si. 
 
 (i) Preparation. Make an intimate mixture 1 of i grm. magnesium 
 powder and 4 grms. of finely powdered quartz-sand. Heat this to 
 bright redness in a wide tube of hard glass. It is best to use the blast 
 lamp for this purpose. The part of the tube containing the mixture 
 should be rotated in the flame. The residue, after a few minutes' heating, 
 is allowed to cool, and treated with water containing HC1. The product 
 5 
 
34 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 consists of amorphous silicon and undecomposed quartz. 2. Test the 
 action of the following reagents upon Si : sulphuric, nitric and hydro- 
 fluoric acids, potash solution and chlorine. (Read Richter, p. 161.) 
 
 SILICON AND OXYGEN. 
 
 (2) Silicon dioxide (Silica, Quartz) SiO 2 . 
 
 i. Test its solubility in the various acids and alkalies. 2. Fuse a mix- 
 ture of i grm. of finely powdered quartz with 4 grms. of Na. 2 CO 3 , in a 
 platinum crucible. Dissolve the product in water. 3. To a portion of 
 this solution add HC1, and evaporate to complete dryness. Take up the 
 residue with water and filter off the insoluble portion. 4. To another 
 portion of the aqueous solution of the fusion add NH 4 C1. (?). Make a 
 bead of salt of phosphorus ; bring a fragment of a silicate or of quartz into 
 it, and heat in the blow-pipe flame for a few minutes (?). 
 
 BORON. B. 
 
 (1) Preparation similar to that of Si. What are its properties? Does 
 it unite directly with other elements ? Is it known in several allotropic 
 modifications ? What is the valency of this element ? 
 
 (Read Richter, pp. 240 and 241.) 
 
 BORON AND OXYGEN. 
 
 (2) Boric Acid -BO 3 . 
 
 i. Dissolve borax in 5 parts of boiling water, add HC1 to acid reaction, 
 and allow to cool. What crystallizes out of the solution ? Dry some of 
 the product by pressing it between filter paper. Test its solubility in 
 water and in alcohol. What do you observe on igniting the alcoholic 
 solution ? Moisten a piece of turmeric paper with an aqueous solution of 
 boric acid, and dry at a gentle heat. What happens? 
 
 Problems. (i) How much CO 2 results from the combustion of 12 
 grms. of carbon ? (2) How much CO 2 will an indefinite quantity of 
 CaCO 3 give, when acted upon by 4.666 grms. of muriatic acid, contain- 
 ing 30 per cent, of pure HC1 ? (3) How many cubic decimeters of CO 
 can be obtained from 90 grms. of oxalic acid ? (4) What amount of SiO. 2 
 can be obtained from 2 grms. of Wollastonite (CaSiO 3 ) ? (5) What is 
 the theoretical quantity of boric acid obtainable from 15 grms. of borax 
 (Na 2 B 4 O 7 -f ioH 2 O)? 
 
METALS OF THE ALKALIES POTASSIUM AND OXYGEN. 35 
 
 METALS. 
 
 CHAPTER VII. 
 
 METALS OF THE ALKALIES POTASSIUM, SODIUM, [AMMONIUM]. 
 
 POTASSIUM. K. 
 
 (1) Preparation. Arrange apparatus as shown in Fig. 35. Into a tube 
 of hard glass, c, introduce a porcelain boat containing about i grm. of a 
 mixture of 138 pts. (i FJG 35 
 
 mol.) of dry (?)K 2 CO 3 
 and 72 pts. (3 at.) of 
 Mg powder. Pass a cur- 
 rent of dry H over it, 
 and after all the air has 
 been displaced in the 
 apparatus (?), light the 
 escaping gas ; heat the 
 part of the tube surround- 
 ing the boat to incipient redness. Observe the brilliant metallic mirror 
 which is formed, and drive it away from the boat by increasing the 
 temperature : it is potassium. Note also the green color of the vapor 
 and the violet coloration it imparts to the burning hydrogen. What is the 
 residue left in the boat ? Test its reaction with litmus (?). 
 
 Formulate the reaction involved in this method of preparation. 
 
 (2) i. Cut a piece of K with a knife, and observe the color and lustre of 
 the fresh surface. Care ! 2. To ascertain whether the metal is fusible, heat 
 a small piece of it in a stream of H. 3. Is it heavier or lighter than water ? 
 
 (3) i. Expose a thin slice of K to the air. What takes place? 2. 
 Throw a small piece of it upon H 2 O (?). In this experiment it is ad- 
 visable to use a tall beaker and to cover the same with a glass plate. 3. 
 What is the action of the halogens upon K? Ask for instructions. 
 
 POTASSIUM AND OXYGEN. 
 
 (4) Preparation of Potassium Hydroxide. In an iron vessel dissolve 
 50 grms. of crystallized Ba(OH) 2 in 160 cc. of water. Cautiously add 
 a hot concentrated solution of 20 grms. of K 2 SO 4 until a sample of the 
 supernatant liquid is no longer precipitated by either K 2 SO 4 or Ba(OH) 2 . 
 Filter rapidly through a plaited filter, and evaporate the solution in an 
 iron or silver dish over a large flame. Continue heating the residue till it 
 appears in a state of quiet fusion. During this operation protect the eyes 
 with a glass plate. Now pour the product upon a clean iron surface, and 
 
36 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 while still warm put it into a bottle provided with a well-fitting stopper. 
 Examine its fracture and color. Try its solubility in water and in alcohol. 
 What is the reaction of the aqueous solution with litmus ? What is an alkali ? 
 
 Salts. 
 
 (5) Potassium Chlorate. KC1O 3 . (See p. 22). 
 
 (6) Potassium Nitrate. KNO 3 . To a hot concentrated solution of 20 
 grms. of NaNO 3 add a solution of 18 grms. of KC1. Boil. What sepa- 
 rates from the warm mixture ? What crystallizes from the mother liquor 
 on cooling ? Recrystallize the latter product. Examine its crystalline form. 
 Is it more soluble in hot than in cold water ? Explain the method of 
 preparation. 
 
 (7) Into a red-hot platinum crucible throw small portions of an intimate 
 mixture of 10 grms. of KNO 3 and i^ grms. of charcoal powder. What 
 takes place? Write the reaction. What is gunpowder ? 
 
 Reactions. 
 
 (8) Use KNO 3 for the following tests. 
 
 i. Place a little of the salt upon the end of a clean platinum wire and 
 introduce it into a non-luminous flame. What color do you observe? 
 View the flame through a cobalt glass (?). 2. To the aqueous solution 
 of the potassium salt add HCland boil. Concentrate by evaporation and 
 add PtCl 4 . What is the composition of the resulting precipitate ? Try 
 its solubility in hot and in cold water, also in alcohol. 3. To the con- 
 centrated solution of the salt add a saturated solution of tartaric acid ; 
 either at once, or on shaking, a white crystalline precipitate appears (?). 
 
 SODIUM. Na. 
 
 (1) How is this metal usually prepared ? 
 
 (2) Study its physical and chemical properties (Richter, p. 285). 
 Wherein does it differ from K? 
 
 (3) Prepare Sodium Amalgam. 
 
 To 500 grms. of dry mercury, contained in a Wedgewood mortar add 
 gradually 5-10 grms. of Na in thin slices. Perform this operation in a 
 good draught chamber, as the union of the two metals is attended with 
 the evolution of light and heat, and poisonous vapors are given off. Stir 
 well with the pestle, allow to cool, and transfer the product to a well- 
 stoppered bottle. What is its action on H 2 O or dilute H 2 SO 4 ? 
 
 SODIUM AND OXYGEN. 
 
 (4) Preparation of Sodium Hydroxide solution. 
 
 Add a little water to 10 grms of fresh quicklime contained in an iron 
 (or porcelain) vessel. Cover the latter, and in a second iron pot dis- 
 solve 25 grms. of soda ash (Na 2 CO 3 ), using about 100 cc. of water. 
 
METALS OF THE ALKALIES SODIUM AND OXYGEN. 37 
 
 Heat the solution to boiling ; stir the quicklime which should have 
 broken up to a white powder with enough water to form a thin paste 
 (milk of lime), and add this gradually to the boiling liquid. Stir well 
 with an iron wire ; transfer the mixture to a bottle ; cork, and allow it to 
 stand. After the supernatant liquid has become perfectly clear, decant it 
 by means of a glass siphon filled with water. It should be preserved in 
 a tightly corked bottle (?). Test a few drops of the solution with BaCl 2 (?). 
 What should the solution contain, and of what does the precipitate, from 
 which it was separated, consist? Write the equation representing the re- 
 action. 
 
 (5) Determine the amount of NaOH contained in the solution. 
 Measure off accurately 20 cc. into a porcelain dish ; add a drop or two 
 
 of phenolphthalein solution, and dilute with water. From a burette care- 
 fully add dilute hydrochloric acid until the red color has just disappeared. 
 Read off the volume of the acid used ; it is the exact quantity needed to 
 neutralize the alkali : 
 
 NaOH -f HC1 = NaCl + H 2 O; 
 
 that is, 40 pts. (i mol.) of NaOH require 36.5 pts. (i mol.) of HC1, 
 and if we know the weight of the HC1 contained in the volume of the 
 dilute acid consumed, a simple proportion will give the weight of the 
 alkali in 20 cc. of the solution. The strength of the acid is determined 
 as follows: In a porcelain dish, dissolve 1.06 grms. of pure Na 2 CO 3 , 
 previously ignited and accurately weighed; add a little phenolphthalein, 
 heat to boiling and introduce acid from the burette until the liquid 
 remains colorless after continued boiling. The carbonate is then exactly 
 neutralized : 
 
 Na 2 CO 3 + 2HC1 = 2NaCl + CO 2 + H 2 O. 
 
 It takes, therefore, 73 pts. of HC1 for 106 pts. of Na 2 CO 3 . Suppose, now, 
 20 cc. of the acid had been used to decolorize the indicator, then i cc. 
 would equal iffi = .53 grms. of Na 2 CO 3 , or .365 grms. of HC1. The 
 latter number is the standard or strength of the dilute acid. 
 
 The phenolphthalein takes no part in these reactions ; it merely indi- 
 cates by its change of color the complete neutralization of the alkali. 
 Why is it necessary to boil the solution when the acid is standardized 
 with a carbonate ? 
 
 Salts. 
 
 (6) Sodium chloride. NaCl. 
 
 Purify common salt. Grind 50 grms. of salt in a mortar with 150 cc. 
 of water. Filter into a beaker, and conduct HC1 gas into the solution, 
 
38 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 as shown in Fig. 36. Pure NaCl separates out. Collect it on a platinum 
 no. 3 6. cone, remove the liquid with the aid of a filter 
 
 pump, and dry the salt by warming it in a porce- 
 lain dish, while stirring it with a glass rod. 
 (7) Sodium carbonate. Na 2 CO 3 . 
 Recrystallize some of the commercial carbo- 
 nate. Heat a portion of the product in a porce- 
 lain dish. What do you observe ? 
 
 Reactions. 
 
 (8) Use the purified chloride for the tests, i. What color do sodium 
 salts give to the flame? 2. Mix a drop of the aqueous solution with 10 
 drops of a PtCl 4 solution on a watch-glass. Evaporate very carefully to 
 a small volume. On cooling, a red colored salt crystallizes out in long 
 monoclinic needles (?). Is it soluble in water? in alcohol? 3. Are 
 there any salts of sodium which are not soluble in water ? Can com- 
 pounds of sodium be precipitated by any reagent ? 
 
 AMMONIUM. 
 
 (1) What is the composition of ammonium? Can it be obtained in a 
 free state? (See Richter, p. 295.) 
 
 (2) Dissolve commercial sal ammoniac in a little water, add ammonia 
 in slight excess, warm, filter if a precipitate is formed, and evaporate to 
 crystallization; stir constantly. Ammonium chloride is thus obtained 
 in the form of a fine powder. 
 
 Reactions. 
 
 (3) i. On a piece of platinum foil heat successively small portions of 
 the chloride, the sulphate, and the nitrate. What occurs in each case ? 
 2. Mix a little NH 4 C1 with burnt lime in a small mortar. Note the odor 
 of the escaping gas and its reaction with litmus. 3. Heat a small por- 
 tion of NH 4 C1 with a caustic soda solution. What is given off? Explain 
 the action of strong bases upon ammonium salts. 4. Add PtCl 4 to a 
 solution of NH 4 C1. Result? 5. To a concentrated solution of the 
 ammonium salt add tartaric acid and shake the mixture (?). 6. Do com- 
 pounds of ammonium impart a color to the flame ? 
 
 Compare the metals of the alkalies with each other. How can the com- 
 pounds of potassium, sodium, and ammonium be distinguished ? 
 
 Problems. i. How much KNO 3 is theoretically obtainable from 2 kilos 
 
METALS OF THE ALKALINE EARTHS CALCIUM. 39 
 
 of Chili saltpetre of 97%, and what amount of Sylvite containing 98% of 
 KC1 is required? 2. Suppose that 75 cc. of dilute HNO 3 were required 
 to saturate 50 cc. of a potash lye; further, that 10 cc. of the acid neu- 
 tralized i. 06 grms. of Na 2 CO 3 , what amount of KOH would the lye con- 
 tain? 3. In the valuation of a pearl ash (impure K 2 CO 3 ), 29.1 cc. of a 
 sulphuric acid were used to neutralize 5 grms. of the sample ; the acid 
 contained 98 grms. of H 2 SO 4 per litre ; calculate the percentage of im- 
 purities in the product. 4. Required the minimum amount of marble that 
 should be burnt to liberate the NH 3 from 50 grms. of NH 4 NO 3 . 
 
 CHAPTER VIII. 
 
 METALS OF THE ALKALINE EARTHS CALCIUM, STRONTIUM, BARIUM. 
 
 CALCIUM. Ca. 
 CALCIUM AND OXYGEN. 
 
 (1) i. Ignite 2 grms. of powdered marble in a platinum crucible to 
 the highest temperature obtainable with the aid of the blast lamp. Con- 
 tinue this for 15 minutes, occasionally stirring the mass with a platinum 
 wire ; what is the residue ? Explain the reaction. 2. Add about 5 cc. 
 of water to the product. What do you observe? Test the reaction of 
 the product with litmus paper. 
 
 (2) i. Prepare lime water. To the slaked lime obtained from 20 
 grms. of quicklime (see p. 37) add i litre of water ; transfer the mixture 
 to a bottle. Cork tightly, shake and allow to stand. When the solution 
 has become clear, draw it off by means of a siphon ? What does it con- 
 tain ? Of what does the undissolved portion consist ? 2. Place a por- 
 tion of the lime water on a watch glass and expose to the air(?). 3. 
 Through a second portion blow air from your lungs (?). 4. Conduct a 
 stream of CO 2 through a third portion and observe carefully the successive 
 changes. Explain them. 5. What takes place upon boiling the clear solu- 
 tion which is obtained as the final product in the preceding experiment ? 
 
 Salts. 
 
 (3) Calcium Chloride. CaCl 2 . 
 
 i. Evaporate some of the spent acid of a CO 2 generator to dryness. 
 What is the residue? 2. Expose a little of the salt to the air (?). 3. 
 What use have you made of CaCl 2 previously? 4. Prepare porous CaC/^ 
 (CaCl 2 -f- 2H 2 O). Dissolve the residue obtained in i in lime water, 
 filter, and neutralize exactly with HC1. Evaporate the filtrate to dryness 
 
40 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 in a porcelain dish, and heat the residue for some time on the sand-bath. 
 The solution of the product must show a neutral reaction. 
 
 (4) Calcium Hypochlorite.C*(C\Q\. (See p. 22.) 
 
 (5) Calcium Sulphate. CaSO 4 . 
 
 i. Carefully heat a few grms. of gypsum in a porcelain dish until the 
 water of crystallization is completely expelled. Pulverize the residue. 
 What happens when it is made into a paste with water and allowed to 
 stand ? 
 
 Reactions. 
 
 Use the pure CaCl 2 for the following tests : 
 
 i. Introduce a small portion of the salt into the Bunsen flame by 
 means of a platinum wire (?). 2. To the aqueous solution add (NH 4 ) 2 
 CO 3 . Result ? 3. To another portion add dilute H 2 SO 4 . What is the 
 composition of the precipitate? Why does it not form in very dilute 
 solutions? 3. Add (NH 4 ) 2 C 2 O 4 to the filtrate from the CaSO 4 (?). 
 * t 
 
 STRONTIUM. Sr. 
 
 Reactions. 
 
 i. What color is imparted to the Bunsen flame by compounds of this 
 element ? 2. Add a CaSO 4 solution to the solution of a strontium salt (?). 
 
 BARIUM. Ba. 
 
 Reactions. 
 
 i. Observe what color Ba compounds give to the flame. Moisten the 
 sample with HC1 before heating it (?). 2. To a portion of the aqueous 
 solution of the chloride add (NH 4 ) 2 CO 3 . What results? 2. Add dilute 
 H 2 SO 4 to a second portion (?). 
 
 Point out how the elements of this group may be distinguished (#) 
 from those of the preceding group ; (b} from each other. 
 
 Problems. i. How much nitric acid of 20 per cent, will effect the so- 
 lution of i grm. of Iceland spar (CaCO 3 ) ? How much CO 2 is given off, 
 and what volume would it occupy at 20 C. under a pressure of 750 mm. ? 
 2. Suppose .5 grm. of sulphur were dissolved in HNO 3 , what quantity 
 of BaCl 2 must be added until it ceases to produce a precipitate ? 3. One 
 grm. of a mineral consisting of the carbonates of Ca, Sr, and Ba, in the 
 proportion of their molecular weights, will leave what weight of the mixed 
 sulphates on treating and evaporating with an excess of H 2 SO 4 ? 
 
MAGNESIUM GROUP MAGNESIUM. 41 
 
 CHAPTER IX. 
 
 MAGNESIUM GROUP MAGNESIUM, ZINC, CADMIUM. 
 MAGNESIUM. Mg. 
 
 (1) Examine the metal in the forms of ingot, ribbon and powder* 
 Note its color, lustre and specific gravity. 2. Introduce a piece of the rib- 
 bon into the flame with the forceps (?). What is the product ? 3. Treat 
 a piece of the ribbon with dilute H 2 SO 4 . Reaction ? 
 
 Salts. 
 
 (2) Magnesium Chloride. MgCl 2 . 
 
 Prepare the ANHYDROUS salt. Dissolve about 50 grms. of the crystal- 
 lized (?) chloride and 50 grms. of NH 4 C1 in as little water as possible. 
 Evaporate to dryness in a porcelain dish. Reduce the mass while hot to 
 small pieces in a mortar, dry it carefully, so as to remove every trace of 
 moisture. It is best to do this by heating small portions of the material 
 in a porcelain crucible until it no longer sinters. A small sample should 
 not give off moisture when heated in a dry test-tube. Be careful also to 
 prevent re-absorption of moisture. Quickly transfer the warm powder to 
 a platinum crucible provided with a well-fitting cover. Heat, at first 
 gently, to expel the NH 4 Cl, then increase the temperature until the mass 
 is in a state of quiet fusion. It is the anhydrous salt which, being 
 extremely hygroscopic, should be preserved in a tightly stoppered bottle. 
 It should dissolve in water to a clear liquid. 
 
 Why cannot the anhydrous chloride be obtained by evaporation of the 
 aqueous solution ? 
 
 (3) Magnesium Sulphate. Mg SO 4 -f yH 2 O. 
 
 Recrystallize some of the commercial salt. What is the form of the 
 crystals ? Taste ? 
 
 (4) Reactions. 
 
 i. Heat a portion of the sulphate or chloride on a platinum wire in 
 the Bunsen flame ; moisten with Co(NO 3 ) 2 solution and heat again. A 
 pink-colored mass results. 2. Add some caustic soda to a little of the 
 solution of the chloride (?). The resulting precipitate dissolves on ad- 
 dition of an ammonium salt (?) 3. Mix a second portion of the chloride 
 solution with NH 3 and NH 4 C1, add Na 2 HPO 4 and agitate the liquid. What 
 is the composition of the precipitate ? Examine it with the aid of a lens. 
 6 
 
42 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 ZINC. Zn. 
 
 (1) How is this metal obtained from its ores? 
 
 (2) Study the physical and chemical properties of Zn (see Richter, p. 
 316). i. Treat a small piece of pure Zn with dilute H 2 SO 4 (?). 2. Re- 
 peat this experiment, substituting the impure commercial metal. What 
 difference do you observe ? What causes it ? 
 
 (3) Granulate commercial zinc. Melt 100 grms. of the metal in a well- 
 covered Hessian crucible. The blast lamp maybe used for this purpose, 
 but it is better to perform the operation in a wind furnace. The crucible 
 is then removed from the source of heat, and allowed to cool until the 
 melted metal no longer takes fire when the cover is lifted. Pour the 
 metal, in a thin stream, into a pail filled with cold water. Drain the 
 product and dry at a moderate heat. 
 
 Salts. 
 
 (4) Zinc sulphate. Zn SO 4 -f 7H 2 O. (See p. 12). i. Prepare some 
 of this salt and recrystallize it carefully from water. 2. Examine the 
 crystals. What other salt have you prepared that exhibits similar forms ? 
 Is there any analogy in the composition of the two salts ? 
 
 Reactions. 
 
 (5) i. Heat a small piece of Zn on charcoal in the oxidizing flame. (?) 
 2. Moisten the incrustation obtained with a drop of Co(NO 3 ) 2 , and heat 
 again. Result ? 3. To a solution of ZnSO 4 add (NH 4 ) 2 S. What is the 
 color of the precipitate ? Try its solubility in dilute HC1 and in HC 2 H 3 O, 
 (acetic acid). 4. Study the action of caustic alkalies, e. g., NaOH upon 
 the Zn solution. 
 
 How could you distinguish between Zn and Mg ? What differences 
 are there between this and the preceding groups ? 
 
 Problems. i. What is the strength of a sulphuric acid of which 20 cc. 
 dissolve exactly .048 grm. of Mg? 2. Suppose it was found that i grm. 
 of Zn gave with H 2 SO 4 , 325 cc. of H at 16 C. and 755 mm., and, 
 further, that .369 grm. of Mg produced the same amount of the gas. 
 Knowing the atomic weight of Mg to be 24, and remembering that the 
 two sulphates are isomorphous, how is it possible to deduce the at. wt. of 
 Zn from the data given ? 
 
MERCURY AND OXYGEN. 43 
 
 CHAPTER X. 
 
 MERCURY, COPPER, SILVER, GOLD. 
 MERCURY. Hg. 
 
 (i) Study the physical and chemical properties of the metal. Wherein 
 does it differ from the other metals? 
 
 MERCURY AND OXYGEN. 
 
 ( 2 ) Mercuric oxide. Hg O . 
 
 How is this substance prepared? What is its behavior on heating? 
 
 Mix a little powdered S with dry Na. 2 CO 3 and HgO. Ignite the mix- 
 ture in a dry test-tube. Extract the residue with water, filter, acidify 
 with HC1 and add BaCl 2 . What has become of the oxide of mercury in 
 this experiment? 
 
 Salts. 
 
 (3) Mercurous Nitrate. HgNO 3 . 
 
 An excess of metallic mercury (use 10-15 g rms O i g treated in the cold 
 with moderately strong HNO 3 until the formation of crystals is no longer 
 noticeable. Redissolve the crystals by warming, filter, and allow to 
 crystallize. 
 
 To prepare a solution of the salt take it up with water acidulated with 
 HN0 3 (?). 
 
 (4) Mercuric chloride. HgCl 2 . 
 
 Dissolve about 5 grms. of Hg in aqua regia. Evaporate to dryness on 
 a water bath. Place the residue into a small dry flask, cover the latter 
 with a watch-glass, and heat cautiously on a sand-bath. What is the 
 sublimate formed in the upper part of the flask ? Dissolve it in four parts 
 of boiling water and allow to crystallize. 
 
 Reactions. 
 
 (5) Mercurous compounds. Use the solution of the nitrate, i. Add 
 a few drops of HC1 to 2 or 3 cc. of the solution. What takes place ? 
 Filter, and add NH 3 to the precipitate (?). 2. Add stannous chloride 
 to another portion of the nitrate solution (?). 3. In a third portion 
 immerse a slip of Cu foil. Examine the stain on the metal ; is it changed 
 when you hold it in the flame? 
 
 (6) Mercuric compounds. The chloride will answer for the tests. 
 
 i. Pass H 2 S through a dilute solution and observe the gradation of 
 colors through which the precipitate passes. What is the final product ? 
 
44 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 2. Add SnCl 2 , drop by drop, to the mercury solution. Explain the 
 changes which occur. 
 
 COPPER. Cu. 
 
 i. Preparation. Ignite the pure oxide in a current of dry H (see p. 
 14). Examine the color and the lustre of the product \ test its solu- 
 bility in HC1, H 2 SO 4 (both strong and dilute), and HNO 3 . Write equa- 
 tions representing the reactions. 
 
 Salts. 
 
 (2) Copper Sulphate. CuSO, -f sH 2 O. 
 
 To 10 grms. of Cu in a flask add 45 grms. of cone. H 2 SO 4 , and heat. 
 When the metal has completely disappeared and the gas (?) ceases to be 
 given off, allow to cool, place the white crystalline residue (?) into a 
 porcelain dish, rinse the flask with hot water. Now add a few drops of 
 HNO 3 to the hot water solution, and filter. From the filtrate the sulphate 
 crystallizes on standing. Recrystallize the product. 
 
 Does this salt suffer decomposition on exposure to the atmosphere? 
 Heat a small quantity in a porcelain crucible, first moderately, then more 
 strongly (?). 
 
 (3) Sulphate of Copper and Potassium. CuK 2 (SOt) 2 -f- 6 H 2 O. 
 Prepare solutions of 10 grms. of blue vitriol and 7 grms. of K 2 SO 4 , 
 
 both saturated at 70. The latter should also contain a few drops of 
 H 2 SO 4 . Mix the solutions ; on cooling the double salt separates in 
 whitish-blue crystals. Examine their form. 
 
 Reactions. 
 
 (4) Use either of the salts you have prepared. 
 
 i. Mix a little of the salt with Na 2 CO 3 , and heat on charcoal in the 
 reducing flame (?). 2. Make a borax bead and dissolve a minute quan- 
 tity of a Cu-compound in it. What color does it give (a) in the oxidiz- 
 ing flame? (<) in the reducing flame? (r) when the bead is reduced with 
 a small piece of tin ? 3. Through a dilute Cu-solution pass H 2 S. Is the 
 resulting precipitate soluble in HC1 or in HNO 3 ? 4. Add ammonia, 
 drop by drop, to the solution. What changes do you observe? 5. To 
 a portion of the very dilute solution add potassium ferrocyanide (?). 
 
 (5) To a solution of copper sulphate in a porcelain dish add a small 
 piece of Zn. Allow to stand over night. Note the result. Has the Zn 
 disappeared ? Does the solution contain any of this metal ? In what 
 form ? Where is the Cu ? 
 
 (6) Repeat the experiment, weighing the copper sulphate (.5 grm.) 
 and the Zn (.2 grm.). Add.HCl in quantity sufficient to insure the 
 
SILVER. 45 
 
 entire solution of the Zn, collect the Cu on a filter, wash with alcohol, 
 dry, heat gently and weigh it in a porcelain crucible. The filtrate should 
 be colorless. 
 
 Compare the weight of the metallic Cu obtained with that of the Zn 
 employed (?). How does the found Cu accord with the calculated amount 
 of that metal in .5 grm. CuSO 4 .5H 2 O? 
 
 Repeat the experiment using Cd in place of Zn. Compare the weights 
 of the metals as before. What deduction can you make ? 
 
 SILVER. Ag. 
 
 (1) Prepare pure Silver from a coin. 
 
 Dissolve a 25-cent piece in nitric acid of sp. gr. 1.2, filter (?), and 
 evaporate the blue (?) solution to dryness. Fuse the residue till it 
 blackens, extract with 250 cc. of water ; filter. Now add ammonia in 
 large excess, and then, cautiously, a sodium bisulphite solution (of about 
 40 %) until on boiling a small portion of the liquid, it is completely 
 decolorized. 
 
 The greater part of the Ag separates from the solution on standing in 
 the cold ; it is well crystallized. The remainder may be precipitated by 
 warming to 70. Digest the product with strong ammonia (?), wash, dry 
 and ignite it. 
 
 Examine the metal carefully. What are its physical and chemical 
 characteristics? 
 
 SILVER AND SULPHUR. 
 
 (2) Silver Sulphide. Ag. 2 S. 
 
 Into a dilute solution of AgNO 3 (see next experiment), containing 
 about 2 grms. of the metal, pass H. 2 S. When the liquid smells of the gas, 
 filter off the black precipitate, wash it with water and dry at 100. 
 
 Salts. 
 
 (3) Silver Nitrate. AgNO 3 . 
 
 Dissolve the Ag obtained in (i) in dilute HNO S and evaporate to dry- 
 ness on the water bath. Dissolve the residue in 80 cc. of distilled water, 
 and preserve the solution in a dark bottle (?). What is its reaction with 
 litmus ? 
 
 Reactions. 
 
 (4) i. Compounds of Ag on charcoal before the blow-pipe give a 
 white metallic globule (?). 2. To a silver solution use the nitrate add 
 HC1. Collect the precipitate on a small filter, wash, dissolve it in am- 
 monia, and add an excess of HNO 3 to the solution (?). Explain these 
 
46 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 reactions. 3. Expose a small portion of the chloride to direct sunlight. 
 Any change ? What practical application is made of this reaction ? 
 (Read Richter, p. 340.)* 
 
 (5) Place strips of the metals Zn, Fe, Sn, Pb, and Cd in a solution of 
 silver nitrate. What is the result in each case? Explain. 
 
 GOLD. Au. 
 
 (1) How could you distinguish the metal Au from Hg, Ag, and Cu? 
 
 Reactions. 
 
 (2) i. Dissolve a small piece of gold (or of a substance containing 
 gold) in aqua regia, concentrate the solution at a gentle heat and pour it 
 into a porcelain dish. Add a solution of FeCl 3 to an SnCl 2 solution until 
 the latter is permanently yellow. After diluting, dip a glass rod into this 
 and then into the gold solution. A purple streak (purple of Cassius) is 
 formed. 2. Add ferrous sulphate to some of the AuCl 3 solution (?). 
 
 In what respects do the members of this group differ from each other, 
 and how can they be distinguished from the metals of the preceding 
 groups ? 
 
 Problems. 1.5 grms. of HgO gave on ignition with carbon 4.63 grms. 
 of metallic mercury; the specific gravity of the vapor of HgCl 2 referred 
 to H, was found to be 135.5. What is the atomic weight of Hg ? 2. 
 The molecule of Hg contains how many atoms, if the vapor density 
 equals- TOO? 3. On analysis a chalcocite was found to contain 20.15 per 
 cent, of S and 79.85 per cent, of Cu. Deduce the molecular formula 
 of the mineral. 4. What quantities of Ag, Au, and Hg can be precipi- 
 tated from their respective solutions by i grm. of Cu ? 
 
 CHAPTER XI. 
 
 ALUMINIUM, TIN, LEAD, BISMUTH. 
 
 ALUMINIUM. Al. 
 
 (i) By what methods is this metal obtained on a large scale? What 
 are its properties? Try the action of the following reagents upon Al : 
 HC1, HNO 3 , and NaOH solution. Write the reactions. 
 
 * If practicable, the instructor should here show and explain the preparation of a 
 photographic negative. 
 
TIN. 47 
 
 Sa/ts. 
 
 (2) Sulphate of Aluminium and Potassium. KA1(SO 4 ) 2 -j- i2H 2 O. 
 
 Prepare saturated solutions of A1 2 (SO 4 ) 3 and K 2 SO 4 ; mix these so that 
 the resulting liquid contains the two sulphates approximately in the pro- 
 portion of their molecular weights. The double sulphate crystallizes on 
 standing. Why? Recrystallize it from water. What is the form of the 
 crystals ? 
 
 What is an alum! (See Richter, p. 351.) 
 
 Reactions. 
 
 (3) Use alum. i. Heat a little of the salt on a platinum wire in the 
 oxidizing flame, moisten with Co(NO 3 ) 2 , and heat again. A blue mass 
 (?) is the product. 2. To an aqueous solution add ammonia (?). Add 
 (NH 4 ) 2 S to another portion of the solution. What do you observe ? 4. 
 To the diluted solution add NaOH, drop by drop. Note the successive 
 changes (?). 
 
 TIN. Sn. 
 
 (1) Examine a bar of this metal, i. Note the sound it emits on 
 bending (?). 2. Etch a smooth surface with HC1 (?). 3. Try the solu- 
 bility of Sn in hot HC1. 4. What action have moderately dilute, and 
 concentrated, HNO 3 upon it? Write the reactions. 
 
 (2) Determine the specific heat of Tin. 
 
 A thin glass beaker of about 200 cc. capacity is carefully covered on 
 the outside with a moderately thin layer of cotton wool. This may' be 
 called the calorimeter. Pour 100 cc. of distilled water into the beaker. 
 Suspend a thermometer in the water. Place 25 grms. of granulated tin 
 into a test-tube, close the mouth of the latter with a plug of cotton. 
 Introduce the test-tube with its contents into a beaker glass containing 
 boiling water. A stout copper wire will serve as a handle. After ten or 
 fifteen minutes the tin will have acquired the temperature of the boiling 
 water 100. The tube is then rapidly removed from the latter and its 
 outer surface freed from moisture by quickly passing a towel over it. 
 Remove the cotton from the mouth, and transfer the tin to the calori- 
 meter. While the metal is being introduced raise the thermometer from 
 the water, and replace it as soon as all the metal has been added ; stir the 
 liquid well and observe, as accurately as possible, the highest point 
 
48 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 reached by the mercury column. Approximate results can be obtained 
 from these data. Calculate as follows : 
 
 Let y temperature of water before introducing the tin. 
 " z " " " after " 
 
 " w = weight of the water. 
 " v = " of the metal. 
 " x = sp. heat then 
 
 x __ loo(z-y) 
 "25 (loo-z) 
 
 (Study Richter pp. 256-259.) Would the specific heat found for tin, 
 when multiplied by the constant 6.4 give the same value as that found in 
 experiment (3) for the equivalent of tin ? Explain. How many series 
 of tin compounds are there ? 
 
 (3) Determine the equivalent weight of Tin. 
 
 Place about 3 grms. of tin in a porcelain crucible that has been pre- 
 viously weighed. Cover the metal with 5-10 cc. of concentrated HNO 3 . 
 Then carefully apply heat by means of an iron plate. The tin is dis- 
 solved, while fumes of NO 2 are set free. When the acid has been entirely 
 expelled, heat the crucible with the white stannic oxide over a Bunsen 
 burner ; allow to cool and weigh. 
 
 Let w = weight of crucible and SnO 2 
 " v = " " " " metallic tin. 
 a. v _ (( (t 
 y 
 
 Then w - v = weight of O, 
 andv - y = " " Sn. 
 
 Equiv. ofSn=i^-^ 8 
 
 W V 
 
 Softs. 
 
 (4) Stannous Chloride. SnCl 2 . 
 
 Dissolve 10 grms. of granulated Sn in warm cone. HC1 with the addition 
 of a few drops of PtCl 4 (?). Put the solution into a well-stoppered bottle. 
 
 Reactions. 
 
 (5) Stannous Compounds. Use the chloride solution. 
 
 i. Conduct H 2 S through a portion of the diluted liquid. A brown 
 precipitate (?) is thrown down. Is it soluble in yellow ammonium sul- 
 phide? What does HC1 precipitate from the sulphide solution? 2. 
 What is the action of HgCl 2 upon SnCl 2 (see p. 44). 
 
 (6) Stannic Compounds. Add a few drops of Br to a portion of the 
 SnCl 2 solution, and boil (?). Use the diluted liquid for the tests. 
 
 1. Pass H 2 S into a portion of the solution. What is the color of the 
 precipitate? Is it soluble in HC1? in (NH 4 ) 2 S? 
 
 2. Add Cu-turnings, boil, decant the liquid, and add HgCl 2 . What 
 happens? Explain.. 
 
BISMUTH. 49 
 
 LEAD. Pb. 
 
 (i) How can this metal be obtained from the oxide? By what physi- 
 cal properties can it be distinguished from other metals ? Is it soluble in 
 the mineral acids? (2) In a solution of 5 grms. of lead nitrate in 
 about 50 cc. of water, suspend a strip of metallic Zn and let stand for a 
 few days (?). 
 
 Salts. 
 
 (3) Dissolve 5 grms. of granulated lead (test-lead) by warming with 
 dilute HNO 3 . Concentrate by evaporation and allow to crystallize. 
 
 Reactions. 
 
 (4) i. Before the blowpipe, on charcoal, lead compounds are reduced 
 to metallic beads, which are sectile with the knife. 2. Add HC1 to a 
 solution of the nitrate. Boil the precipitate with water. (?) What takes 
 place on cooling? 3. To another portion add dilute H 2 SO 4 (?). 4. 
 Pass H 2 S into a third portion (?). 
 
 BISMUTH. Bi. 
 Reactions. 
 
 (i) i. Mix a little of the oxide or nitrate of Bi with Na 2 CO 3 and heat 
 in the reducing flame on charcoal. Does the resulting metallic globule 
 resemble lead ? Is it sectile? 2. Pass H 2 S into a solution of the chloride 
 or nitrate in HC1 (?). 3. Add a large volume of water to a bismuth 
 solution. What occurs? What reactions distinguish Al, Sn, Pb and Bi 
 from each other, and from the metals previously studied ? 
 
 Problems. i. What is the molecular formula of a mineral containing 
 
 SiO 2 = 43.08 
 A1 2 3 = 36.82 
 CaO = 20.10 
 
 100.00 
 
 2. A compound of tin and chlorine yielded on analysis 29.42 parts of 
 Sn and 35.40 parts of Cl ; its vapor density was ascertained to be 132.85. 
 What is the atomic weight of tin ? 3. Deduce the formula of Cosalite 
 from the following analysis : 
 
 S = 15.27 
 
 Bi = 41.76 
 
 Pb = 40.32 
 
 Ag = 2.65 
 
 IOO.OO 
 
50 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 CHAPTER XII. 
 
 CHROMIUM, MANGANESE, IRON, NICKEL/ COBALT. 
 CHROMIUM. Cr. 
 
 CHROMIUM AND OXYGEN. 
 
 (1) Chromic oxide. Cr. 2 O 3 . 
 
 i. Preparation. Mix intimately 20 grms. of potassium dichromate 
 and 4 grms. of sulphur. Heat the mixture in a porcelain crucible over 
 the blast lamp for about 20 minutes. Cool, extract the residue with 
 boiling water and dry it at a gentle heat. What is its color ; is it soluble 
 in dilute HC1? 2. Fuse a portion of it with six times its weight of 
 NaHSO 4 in a platinum crucible. What takes place? 3. Repeat this 
 experiment with some finely powdered chromite. (?) 
 
 Salts. 
 
 (2) Chromic Chloride. CrCl 3 . 
 
 Prepare the anhydrous salt. Intimately mix 10 grams of Cr 2 O 3 , pre- 
 pared as described, and 3 grms. of powdered charcoal, and convert this 
 into a dough with a little starch paste. Form the product into balls of the 
 size of a pea ; dry, and then ignite these (covered with charcoal powder) 
 in a Hessian crucible, provided with well-fitting lid. Place the residue 
 into a tube of hard glass, and heat it in a current of CO. 2 to expel every 
 trace of moisture. With the aid of a blast lamp increase the tempera- 
 ture and replace the CO 2 by a current of Cl. The excess of Cl should 
 be absorbed by conducting it into a bottle filled with caustic soda. (?) 
 The resulting CrCl 3 sublimes to the cooler portions of the tube. Describe 
 its appearance. Is it soluble in water? 
 
 What other chlorides are prepared in a similar way ? Write the equa- 
 tion, expressing the reaction. 
 
 (3) Chrome Alum. Cr 2 (SO 4 ) 3 .K 2 SO 4 + 24 H 2 O. 
 
 Dissolve 10 grms. of K 2 Cr 2 O 7 in a little water ; acidify with H 2 SO 4 , pass 
 SO 2 into the liquid until the latter is saturated with the gas. Allow to 
 stand ; the double salt crystallizes. What is its crystalline form ? Dis- 
 solve some of it in cold water and note the color of the solution ; now 
 warm it. What takes place (see Richter, p. 374)? 
 
 (4) i. Examine crystals of potassium dichromate, K 2 Cr 2 O 7 . How is it 
 obtained? 2. Dissolve 10 grms. of this salt in water, and from a burette 
 carefully add a caustic soda solution until the color is changed to yellow (?). 
 
MANGANESE AND OXYGEN. 51 
 
 What crystallizes from the solution on evaporation ? How can you re- 
 convert the product into the dichromate ? 
 
 Reactions. 
 
 (5) i. Dissolve a minute quantity of a chromium compound in a borax 
 bead. Heat in the oxidizing and in the reducing flame. Results? 
 2. Heat a little of the compound with KNO 3 on a platinum foil (?) 
 
 (6) Chromic compounds. Use chrome alum for the tests. i. Add 
 caustic soda, drop by drop, to a little of the solution. (?) Continue the 
 addition of the reagent till the precipitate is redissolved. What takes 
 place on boiling the solution ? 2. What is the action of ammonia on the 
 solution of the chromium salt ? 
 
 (7) Chromates. Use a solution of potassium chromate. i. Add lead 
 acetate solution. Note the color of the precipitate. Is it soluble in 
 acetic acid? 2. Substitute BaCl 2 for the lead salt in the preceding ex- 
 periment. (?) 3. Acidify the chromate solution with H 2 SO 4 and add 
 H 2 O 2 to the liquid. What happens? 4. To some of the chromate 
 solution add a few drops of HC1 and about i cc. of alcohol. What occurs 
 when the mixture is heated to boiling ? 
 
 MANGANESE. Mn. 
 MANGANESE AND OXYGEN. 
 
 (1) In what proportions do these two elements unite with each other? 
 Enumerate the oxides which occur in nature. What is formed when the 
 oxides of manganese are heated in H ? When they are ignited in the 
 air? 
 
 Salts. 
 
 (2) Manganous Chloride. MnCl 2 -f- 4H 2 O. 
 
 Evaporate in a porcelain dish the solution obtained in the preparation 
 of Cl from MnO 2 and HC1. Heat the dry residue over a small flame for 
 some time. Add much water and boil. Filter, and to ^ of the filtrate 
 add a solution of Na 2 CO 3 in excess. Allow the precipitate (?) to settle, 
 draw off the supernatant liquid with a siphon, and wash the remaining 
 precipitate several times with water by decantation. Add the precipitate 
 then to the principal solution and digest at a gentle heat until a small 
 filtered sample mixed with (NH 4 ) 2 S gives a flesh-colored precipitate which 
 is completely dissolved by dilute acetic acid. Now filter and evaporate 
 to crystallization. 
 
 (3) Potassium Manganate K 2 MnO 4 and 
 
 Potassium Permanganate K 2 Mn 2 O 8 . 
 
52 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 In a porcelain crucible fuse a mixture of 5 grms. KOH and 2.5 grms. 
 KC1O 3 \ gradually add 5 grms. finely powdered MnO 2 . Maintain a 
 moderate red heat for 15 minutes. Dissolve the dark-green residue in a 
 little water. Observe the color of the solution. What does it contain? 
 Then dilute with much water and conduct CO 2 into the liquid. Is there 
 any change ? If so, write the equation expressing it. 
 
 K 2 Mn 2 O 8 as well as K 2 MnO 4 are powerful oxidizing agents. Warm a 
 little of the alkaline K 2 MnO 4 solution with a few drops of alcohol (?). 
 To a little of the permanganate solution, acidified with H 2 SO 4 , add sul- 
 phurous acid (?). Treat the acidified solution also with solutions of ferrous 
 sulphate and oxalic acid (?). 
 
 Reactions. 
 
 (4) i. What color do Mn-compounds impart to a borax bead in the 
 oxidizing flame? What is the effect of the reducing flame? 2. Heat a 
 little of an Mn-compound with Na 2 CO 3 and KNO 3 on a platinum foil. 
 What does the resulting mass contain ? 3. To a little of the solution of 
 the chloride in water add (NH 4 ) 2 S. What is the color of the precipitate. 
 Test its solubility in acids (including acetic acid). 4. Add caustic soda 
 to another portion of the chloride solution. Is the precipitate soluble in 
 an excess of the reagent ? Is its color affected by exposure to the air ? 
 Explain. 
 
 IRON. Fe. 
 
 (1) Preparation. Into a tube of Bohemian glass place a porcelain 
 boat filled with the finely powdered oxide. Pass a current of dry H 
 through the tube, and when all the air is expelled (how could you test 
 it?), apply heat to that part of the tube which contains the boat. What 
 is formed in the anterior portion of the tube ? After a red heat has been 
 maintained for 10 minutes allow the boat to cool in H, and examine its 
 contents. Are they attracted by the magnet ? Expose the product to 
 air (?). 
 
 How is iron obtained from its ores on a large scale ? What are its 
 properties? (see Richter, pp. 393 and 394). Distinguish between cast- 
 iron, steel, and wrought-iron. 
 
 Salts. 
 
 (2) Ferrous Sulphate. -Fe SO 4 -f 7H 2 O. 
 
 To 25 grms. of Fe in the form of nails or wire, free from rust, contained 
 in a flask, add 200 cc. of dilute (i : 4) sulphuric acid. When the evolu- 
 tion of the gas (? Note its odor !) is no longer violent, warm, and finally 
 boil until the liberation of gas ceases. A sample of the solution poured 
 
IRON. 53 
 
 into a test-tube should, on cooling, give a copious separation of crystals. 
 Filter into a casserole containing 2-3 cc. of cone, H 2 SO 4 , and let stand 
 for 8 hours. Collect the crystallized product in a funnel the stem of 
 which is closed with a loose plug of glass wool,* allow the mother 
 liquor to drain off, wash with very little cold water (?), and dry between 
 sheets of filter paper. Examine the product carefully. Note its color, 
 taste, solubility in water and crystal form. What other salts of analo- 
 gous composition are isomorphous with it ? 
 
 What is observed when some of the salt is heated, first moderately, then 
 strongly, in a tube of hard glass ? 
 
 Expose the aqueous solution of the salt to the air for several hours (?). 
 
 (3) Ferrous Ammonium Sulphate. Fe (NH 4 ) 2 (SO 4 ) 2 -f- 6H 2 O. 
 
 In 100 cc. of dilute sulphuric acid dissolve clean iron wire till no more 
 hydrogen is given off; neutralize a like quantity of the acid exactly with 
 ammonia water, and add to it a few drops of dilute sulphuric acid. Filter 
 the iron solution into that of the ammonium salt. Let the salt crystallize, 
 drain it on a funnel provided with a perforated platinum cone, wash and 
 dry as described under (2). Preserve in a well-stoppered bottle. What 
 metals can replace the iron in this salt without altering its crystalline 
 form? 
 
 (4) Ferric Ammonium Sulphate. Fe 2 (SO 4 ) 3 .(NH 4 ) 2 SO 4 -f 24H 2 O. 
 Place 20 grms. of crystallized ferrous sulphate into a porcelain dish 
 
 together with a few cc. of water and 3.5 grms. of oil of vitriol. Warm 
 on an asbestos plate, adding nitric acid, drop by drop, until no further 
 change of color (?) is observed. Evaporate the excess of HNO 3 , dissolve 
 the residue in hot water and add 3.5 grms. of (NH 4 ) 2 SO 4 ; filter, and set 
 the solution aside for crystallization. Separate the crystals from the mother 
 liquor, and wash and dry them as under (2). To what class of substances 
 does this salt belong ? Why ? 
 
 Reactions. 
 
 (5) In a borax bead dissolve a small quantity of an iron compound, 
 and treat it successively in the oxidizing and reducing flames. What 
 changes do you observe ? 
 
 (6) Ferrous Compounds. Use a freshly prepared solution of ferrous 
 sulphate for the following tests: i. To a few drops of it, diluted with 
 water, add ammonia. Note the color of the precipitate, and the changes 
 which occur on exposure to the air (?). 2. Add (NH 4 ) 2 S to another 
 portion (?). Is the resulting precipitate soluble in HC1 ? 3. In a porce- 
 lain capsule bring together a little of the ferrous solution and a drop of 
 
 * It is better to use a perforated platinum cone, and to remove the adhering solution 
 with the aid of a filter pump. 
 
54 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 a potassium ferrocyanide solution. Result ? 4. In a similar manner test a 
 drop of the iron solution with ferricyanide of potassium. 
 
 (7) Ferric compounds. In the presence of free acids, oxidizing agents 
 convert iron compounds from the ferrous into the ferric condition, i. 
 Acidify the ferrous sulphate solution with sulphuric acid, warm, and add 
 cone. HNO 3 until it fails to produce a change in color ; the iron is then 
 in the ferric state. 2. Dilute a few drops of the yellow liquid with 
 several cc. of water and add ammonia (?). 3. Test a drop of the ferric 
 solution with potassium ferrocyanide (?). 4. Treat a. second drop with 
 ferricyanide of potassium (?). 5. Mix another drop with a solution of 
 potassium sulphocyanate (?). 6. Conduct H 2 S into some of the ferric 
 sulphate solution. What do you observe? Explain the reaction, and 
 write the equation expressing it. 7. Place a piece of metallic Zn in a 
 test-tube containing a solution of the ferric salt. What takes place ? 
 
 (8) Quantitative estimation of iron. Under manganese it was observed 
 that the salt potassium permanganate is an oxidizing agent. To show 
 how this salt acts with iron in its lower form of oxidation, fill a burette 
 with an aqueous solution of it ; allow it to drop slowly into the solution 
 of a ferrous salt acidulated with H 2 SO 4 . The pink color of the perman- 
 ganate immediately disappears on stirring with a glass rod. This con- 
 tinues until the ferrous salt is completely oxidized to the ferric state. A 
 drop of permanganate added in excess will then impart a faint pink color 
 to the liquid. This indicates that the reaction is ended. Write the 
 equation. 
 
 This behavior may be utilized for determining the quantity of iron in a 
 solution. That this may be done, it is first necessary to standardize the 
 FIG. 37- permanganate solution. Proceed as follows: Dissolve about 2 
 grms. of the permanganate in 1000 cc. of H 2 O. Fill a burette 
 'with this solution. Weigh out .2 grm. of clean piano wire. 
 Place this into a small flask (Fig. 37) provided with a cork 
 and valve.* Cover the iron wire with dilute sulphuric acid. 
 Warm. When the iron is completely dissolved, remove the 
 cork, add cold water to the solution, and slowly admit the per- 
 manganate until the final pink coloration appears. Note the volume of 
 the K 2 Mn 2 O 8 required to produce this effect. Suppose 30 cc. had been 
 consumed, then : 
 
 30 cc. K 2 Mn 2 O 8 = .2000 grm. metallic iron. 
 I " = .00666 " " 
 
 This is then the standard of the permanganate in iron units. 
 
 * With a sharp knife make a longitudinal incision of about I cm. length, in a rubber 
 tube, and close one end by means of a glass rod. 
 
COBALT AND NICKEL. 55 
 
 Next, dissolve i grm. of ferrous ammonium sulphate in 100 cc. distilled 
 water, add 5 cc. H 2 SO 4 , and then introduce the permanganate until the 
 final reaction is observed. Calculate the percentage of iron in this salt 
 and compare the experimental result with the theoretical value. 
 
 How much oxygen will each molecule of K 2 Mn. 2 O 8 give up in oxidizing ? 
 How many molecules of FeO can be changed to Fe 2 O 3 by a molecule of 
 K 2 Mn 2 8 ? 
 
 COBALT. Co AND NICKEL. Ni. 
 Reactions. 
 
 i. Dissolve a minute quantity of a cobalt compound in a borax bead. 
 Heat first in the oxidizing, then in the reducing flame (?). 2. What is the 
 behavior of nickel compounds under like conditions? 3. Add caustic 
 alkali to a solution of Co(NO 3 ) 2 , warm the mixture (?). What action 
 have caustic alkalies on solutions of nickel salts? 5. To the cobalt solu- 
 tion cautiously add ammonia. After a precipitate (?) has formed, add 
 more of the reagent. What takes place ? Expose the resulting solution 
 to the air in a shallow dish (?). 6. Treat a nickel solution in an analo- 
 gous manner (?). 7. To the solutions of Co and Ni, each in a separate 
 test-tube, add (NH 4 ) 2 S. Filter and wash the precipitated sulphides, and 
 test their solubility in acids (?). 
 
 Note the colors of cobalt and nickel salts, in the hydrated as well as in 
 the anhydrous state. 
 
 Is there any marked difference between Co and Ni in respect to their 
 chemical deportment? 
 
 Point out the differences in the reactions of Cr, Mn, Fe, Co, and Ni. 
 
 How may ferrous compounds be distinguished from ferric ? What con- 
 ditions are favorable to the conversion of the former into the latter? 
 The latter into the former ? 
 
 By what means may chromic salts be changed into compounds of 
 chromic acid? How may the reverse change be effected? 
 
 Devise a method for separating the elements treated in this chapter. 
 
 Problems. i. How much K 2 Cr 2 O T can be obtained theoretically from 
 100 kilos of a chromite containing 58.6 per cent, of Cr 2 O 3 ? 2. 100 
 grms. of a pyrolusite which was found to contain 4 per cent, of impuri- 
 
56 EXPERIMENTS IN GENERAL CHEMISTRY. 
 
 ties, will give what volume of O, measured at 20 C and 745 mm., when 
 strongly ignited ? What is the weight of the residue, assuming that one- 
 half of the impurities was moisture, the other half quartz ? 3. How 
 many grms. of HNO 3 are required to oxidize 12 grms. of crystallized ferrous 
 sulphate ? 4. What percentage of metallic iron is contained in a salt, of 
 which .7 grm. are exactly oxidized by 17.8 cc. of permanganate solu- 
 tion (standard : i cc. = .0056 grm. Fe) ? 
 
APPENDIX. 
 
 TABLE OF METRIC WEIGHTS AND MEASURES, 
 
 MEASURES OF LENGTH. 
 
 I metre = 10 decimetres 100 centimetres = 1000 millimetres, 
 i metre 1.09363 yards = 3.2809 feet = 39.3709 inches. 
 
 MEASURES OF CAPACITY. 
 
 i cubic metre 1000 litres 1,000,000 cubic centimetres 1,000,000,000 cubic 
 millimetres. 
 
 i litre = 61.02705 cubic inches = .035317 cubic foot 3= 1.76077 pints = .22097 gallon. 
 
 MEASURES OF WEIGHT. 
 I gram = weight of i cc. of water at 4 C. 
 
 I Kilogram = 1000 grams = 100.000 centigrams = 1,000,000 milligrams. 
 I Kilogram = 2.20462 Ibs. = 35-2739 ounces = 15432.35 grains. 
 
 TABLE OF ATOMIC WEIGHTS OF ELEMENTS. 
 
 Aluminium . 
 
 . Al . . 
 
 27.0 
 
 Lead . . . 
 
 . Pb . . 
 
 .... 207.0 
 
 Antimony 
 
 . Sb . . 
 
 I2O.O 
 
 Magnesium 
 
 Mg 
 
 24.0 
 
 Arsenic 
 
 As 
 
 7S.O 
 
 Manganese 
 
 Mn 
 
 cc o 
 
 Barium . . 
 
 . Ba . . 
 
 137.0 
 
 Mercury 
 
 H g 
 
 2OO.O 
 
 Bismuth 
 
 Bi 
 
 . . . 208.0 
 
 Molybdenum 
 
 Mo . 
 
 Q6.O 
 
 Boron . . . 
 
 . B . . 
 
 II.O 
 
 Nickel . . . 
 
 . Ni . . 
 
 C.O.O 
 
 Bromine . . 
 
 . Br . . 
 
 . . . 80.0 
 
 Nitrogen 
 
 N 
 
 14.0 
 
 Cadmium 
 
 Cd 
 
 II2.O 
 
 Oxygen 
 
 o 
 
 16.0 
 
 Calcium . . 
 
 . Ca . . 
 
 4O.O 
 
 Phosphorus 
 
 . P . . 
 
 31.0 
 
 Carbon 
 
 C 
 
 I2.O 
 
 Platinum 
 
 Pt 
 
 195.0 
 
 Chlorine . . 
 
 . Cl . . 
 
 35.5 
 
 Potassium . . 
 
 . K . . 
 
 .... 39.0 
 
 Chromium 
 
 . Cr . . 
 
 C2.C, 
 
 Silicon 
 
 . Si 
 
 . . . 28.0 
 
 Cobalt 
 
 Co 
 
 SO.O 
 
 Silver 
 
 Ag 
 
 108.0 
 
 Copper . 
 
 . Cu . . 
 
 63.3 
 
 Sodium . . . 
 
 . Na . 
 
 23.0 
 
 Fluorine 
 
 Fl 
 
 IQ.O 
 
 Strontium 
 
 . Sr . 
 
 87.5 
 
 Gold . . . . 
 
 . Au . 
 
 ..... 197.0 
 
 Sulphur . . . 
 
 . S . . 
 
 .... 32.0 
 
 Hydrogen 
 
 . H . 
 
 i.o 
 
 Tin ... 
 
 . Sn . 
 
 118.0 
 
 Iodine . . . 
 
 . I . 
 
 . . 127.0 
 
 Zinc 
 
 . Zn . 
 
 .... 65.0 
 
 Iron 
 
 Fe 
 
 c6.o 
 
 
 
 
 
 
 
 
 
 
 57 
 
UNIVERSITY OF CALIFORNIA LIBRARY 
 BERKELEY 
 
 Return to desk from which borrowed. 
 This book is DUE on the last date stamped below. 
 
 DEC 3 19*' 
 
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