QD 
 
 45- 
 
 SB flOM 
 
 A Laboratory Manual 
 of General Chemistry 
 
A LABORATORY MANUAL 
 
 OF 
 
 GENERAL CHEMISTRY 
 
 BY 
 
 W. S. HENDRIXSON, 
 
 Professor of Chemistry in Grinnell College. 
 
 GRINNELL, IOWA, 
 
 1918 
 
ALL RIGHTS RESERVED 
 
PREFACE. 
 
 This the fourth edition of the author's Experiments in General 
 Chemistry is printed, as its predecessors have been, primarily for the 
 use of students in general chemistry in Grinnell College. This fact 
 may explain certain departures from custom in the preparation of 
 such books, such as suggestions to teachers and detailed descriptions 
 of apparatus and its manipulation. As a matter of fact apparatus at all 
 complicated is not only shown by cuts, but it is set up on the lecture 
 table and many experiments for any period are there carried through 
 before the students enter the laboratory. Some apparatus is even set 
 up in the laboratory and left there for inspection during the laboratory 
 perio'd. 
 
 Not satisfied to use the same set of even his own experiments year 
 after year and wishing to provide new laboratory work for classes of 
 students who have taken chemistry in the high school, the writer has 
 provided for more laboratory work than can be done in a three- or 
 four : hour course of one year. 
 
 In this book an attempt is made to connect rationally general 
 chemistry and qualitative analysis. Students who complete a first 
 year course in chemistry should have some knowledge of qualitative 
 analysis, but it should not be permitted to take the place of general 
 chemistry in the second half year, which is usually devoted to study of 
 the metals. Qualitative analysis ought to be a development from the 
 general chemistry to which it gives point, and its introduction as an 
 outgrowth of the general chemistry greatly stimulates the student's 
 interest in both subjects. In this book tests for acids and other com- 
 pounds are given in the study of the non-metals, and a system for the 
 detection of acids is given after the study of the non-metals has been 
 completed. In the study of the metals emphasis is placed on properties 
 that are of analytical significance, though other facts are not neg- 
 lected. After each group of metals has been studied their separation 
 is taken up, and the work is extended as rapidly as the student's ex- 
 perience justifies it, to the detection of both metals and acid radicals 
 in "unknowns." The scheme of qualitative analysis as outlined is not 
 supposed to be complete but is meant to serve as an introduction to 
 the subject and a preparation to the more rigorous course in quali- 
 tative analysis the following year. 
 
 387354 
 

 TO THE STUDENT. 
 
 I. On coming to the laboratory the first day find the number of 
 your desk, get the key or combination and a printed list of the appa- 
 ratus the desk should contain. Verify the apparatus, asking the names 
 of the things you do not know. Make sure that apparatus is wanting 
 before calling for it. Present broken or faulty apparatus for exchange. 
 
 II. After using apparatus, clean it and return it to the desk. 
 Clean the top of the desk and lock it before leaving. Have old cloths 
 for cleaning and a towel for the hands. In grading notes you will be 
 held responsible for the condition of the desk without and within anfl 
 for the bottles on its shelves, which should always be kept in the 
 same order. 
 
 III. Provide an approved note book of the sort shown in the lec- 
 ture room. Leave the first leaf blank and enough margins to permit 
 corrections by the instructor. 
 
 Begin every experiment on a new page. Write notes in the labor- 
 atory in full and do not copy them. Only original notes made while 
 doing the laboratory work are of value. The notes on each experiment 
 should have a suitable heading and should have the same number as 
 the experiment. Paragraph to suit the subject matter. Present the 
 note book for grading and criticism before leaving the laboratory. 
 
 IV. Ascertain in advance what the laboratory work is to be for 
 each period and prepare for it before coming to the laboratory by use 
 of the text or reference book. The subject will be announced in the 
 class room, and references to books will be given. Bring the text- 
 book to the laboratory. 
 
 V. Every experiment is intended to teach something of value. 
 Try to find out what it is by yourself, using observation and the text- 
 book, then ask the instructor. State results in the notes. Do not con- 
 fine these efforts merely to answering the questions of the laboratory 
 book. For example after equations have been studied in class all 
 equations for reactions in the laboratory work should be written 
 whether asked for or not. Write them where they belong scattered 
 through the notes and do not write them in mass at the end of the 
 notes. 
 
 VI. When using chemicals replace the covers or stoppers of the 
 containing vessels. Do not throw stoppers upon the desk. Return all 
 bottles to their places on the shelves. Replace all weights in the 
 boxes. 
 
 VII. Do not wash pieces of apparatus with distilled water but. 
 with hydrant water. If they are well drained it will not be necessary 
 in most cases even to rinse them with distilled water. On the other 
 
LABORATORY MANUAL OP GENERAL CHEMISTRY 3 
 
 hand always have distilled water in your wash bottle and use it in 
 making all solutions. 
 
 VIII. Throw no solid matter into the sinks but into the jars pro- 
 vided for that purpose. On the other hand wash apparatus at the 
 sinks and do not throw large amounts of water into the jars. 
 
 DIRECTIONS FOE LABORATORY WORK. 
 GENERAL MANIPULATION. 
 
 1. The Bunsen Burner: Take apart the burner and study its con- 
 struction. Determine how to regulate the supplies of gas and air. Put 
 it together and turn on the gas at the cock and regulate the supply by 
 the screw on the burner. Regulate the gas and air so as to secure a 
 non-luminous flame. Too much air may cause the flame to blow out, 
 "snap back" or burn with much noise. A long green hissing flame in- 
 dicates that the gas is burning also at the base. Turn out the flame, 
 reduce the amount of air, light again. A flame about 3 inches long is 
 usually sufficient. 
 
 The ordinary flame is used to heat test tubes directly, or to heat 
 such as flasks and beakers placed upon wire gauze, of which nichrome 
 gauze is best. The crown top is used to heat beakers and flasks with- 
 out protection of gauze. 
 
 The Wing Top : Put the wing top on the burner, turn on the gas 
 and regulate gas and air so as to secure a just non-luminous flame 
 about as high as broad. Too much air is objectionable. Such a flame 
 should be used exclusively for bending tubing. 
 
 '2. Breaking and Bending Glass Tubing: Draw once a sharp, three- 
 cornered file across the piece of tubing where you wish to break it. 
 Hold the tube with both hands, with thumbs together and opposite the 
 scratch. A slight pull will break the tube squarely at the scratch. 
 
 To bend glass tubing always use the wing top. No one but an ex- 
 pert can make a gcod bend with the flame of the ordinary burner. A 
 bend so made will be uneven, crinkled and the tube is likely to break 
 at that point. 
 
 Hold the tube above the middle of the wing flame, having its 
 length in direction of the breadth of the flame. Revolve the tube and 
 move it back and forth in the direction of its length until it softens 
 and begins to yield to a slight pressure, then bend it slowly as de- 
 sired. If it shows a tendency to collapse or flatten at any point, stop 
 bending at that point, heat at one side or the other and there complete 
 the bend. The tendency to collapse shows that the bend is being made 
 too short. Practice with scraps of tubing until you can make a good 
 bend, and then make the tubes like the samples shown in fig. 1. 
 They represent the tubes most used in this course, and should be kept 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Fig. 1 
 
 throughout the year and not cut up to make other things. Heat the 
 ends of the tubes to redness in the ordinary flame to fuse down the 
 sharp edges. Other glass working will be shown in the class room as 
 occasions arise. 
 
 If there is no wash bottle in the desk construct one as shown in 
 fig. 1. This wash bottle should always contain a supply of distilled 
 water and should not be torn down to use the tubes or flask for any 
 other purpose. 
 
 3. Weighing. The following may be applied to all balances and 
 weighings, but will be supplemented by the instructors when weigh- 
 ings of great accuracy are required. 
 
 Never place chemicals, other than pieces of metals, directly upon 
 the pans of the balance, but in some suitable containing vessel. When 
 only moderate accuracy is required, balanced papers may be used, but 
 not where accurate weighings are called for. In the latter case two 
 methods are good. Place a dish on the left hand pan and weigh it ac- 
 curately. Place in it about the desired amount of substance and weigh 
 again. Of course the difference between the weights gives the weight 
 of the substance. The scales rarely "balance" or the pointer rarely 
 stands at the zero point. This error, however, is eliminated by the 
 above method. Instead of the dish, the substance may be weighed in 
 a corked tube, then the desired amount may be taken out and the tube 
 weighed again. Of course in both weighings the tube must be in the 
 same pan, always the left. 
 
 When equilibrium has been attained, that is, when the pointer 
 makes excursions equidistant right and left from the center, count the 
 weights without removing them from the pan, record the weight in the 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 5 
 
 note bcok and verify by counting the weights again as they are re- 
 moved to the box. Write the weight in the book as one number. If 
 the weights are 5 g., 2 g., 200 mg., 50 mg., write the total at once 7.25 
 grains. Unless the nature of the experiment requires it, it is a mis- 
 take to try to weigh a definite amount of a substance. For example, if 
 the directions say "Weigh accurately about 5 grams of the substance," 
 do not try to get just five grams, but take about five grams and weigh 
 it accurately. 
 
 4. Pleasuring Volumes of Liquids: Study a pipet and a gradu- 
 ated cylinder. Why is the cylinder graduated from the top downward 
 and from the bottom upward? 
 
 Weigh accurately a small flask on the horn-pan scales observing 
 the directions given under weighing. Now put a small piece of rub- 
 ber tubing on the larger end of the pipet, by suction fill it with dis- 
 tilled water to above the mark and pinch the rubber tube. Carefully 
 let out the water to the mark and then run the remainder into the 
 flask and weigh again. Find the weight of the water, the weight of 
 1 cc. and the error in the pipet according to your work. Why in this 
 experiment does it make little difference whether the balance with 
 pans empty was in exact equilibrium or not? 
 
 OXYGEX. 
 
 5. Preparation of Oxygen: (Study text-book in advance). Many 
 substances containing it give off all or a part of their oxygen when 
 heated. The following are illustrations: 
 
 In a small test tube heat about half a gram, estimated, of potas- 
 sium chlorate. It melts and then seems to boil owing to the evolution 
 of oxygen. Light a splinter or tooth pick, blow out flame, lower glow- 
 ing end into the test tube. It should rekindle. 
 
 In another small tube, preferably one made of "hard glass" tubing 
 heat a little mercury oxide persistently and test with glowing splinter. 
 Note sublimed mercury on wall of tube. By persistent heating all the 
 mercury oxide may be decomposed into mercury and oxygen. 
 
 As in the case of mercury oxide heat a little manganese dioxide 
 and test for oxygen. Only one-third of the oxygen is given off. 
 
 Oxygen may be obtained also by heating lead dioxide, potassium 
 permanganate, potassium perchlorate, barium dioxide to redness. 
 
 All these instances illustrate chemical change, and the nature of 
 the change is that of decomposition. 
 
 ({. Preparation of Oxygen, Laboratory Method: (a) Read through 
 the experiment and have everything required within reach. Set up 
 the apparatus as shown in fig. 2, having trough filled with water to 
 about 2 inches above the shelf. Place the delivery tube well to the 
 farther side of the trough so as not to be obliged to reach over it. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Fig. 2. 
 
 Fill two jars* with water by immersing in the trough with mouth 
 slightly upward, inverting them and placing upon the shelf as shown. 
 Direct the base of the retort stand backward so that it will not be in 
 the way of the burner. (See fig. 2). There it no need of clamping the 
 test tube tightly so as to endanger breaking it. Have the clamp near 
 its mouth. Why? 
 
 Weigh approximately on platform scales 8 grams potassium chlor- 
 ate and 4 grams powdered manganese dioxide, mix them on paper with 
 a spatula and from the paper slide the mixture into the test-tube. 
 
 Now heat the substance in the tube slowly and evenly, best by 
 holding the burner in the hand and moving it back and forth along the 
 tube, so as to avoid over-heating and softening the tube at any one 
 point. So regulate the heat that the gas shall be coming off slowly 
 when the jar is about full. When the jar is full slide it to one side 
 and place another jar over the mouth of the tube, then place the cover 
 securely on the full jar keeping its mouth under water, remove it *o 
 the desk and clamp on the cover. Fill five jars or bottles. When the 
 gas ceases to come off remove the stopper of the tube to prevent a 
 back flow of the water. Allow the tube to cool and meantime use the 
 oxygen in 7. Return then to (b). 
 
 (b) Prepare a funnel with filter paper. Fold a filter in halves, 
 then in fourths, open out one thickness making a cone, press down 
 evenly into a funnel and wet it with water to make it hold its shape. 
 Fill half full of distilled water the tube used to heat the potassium 
 chlorate and manganese dioxide, and when the residue is wet through- 
 out by shaking heat to boiling and filter by pouring the contents into 
 the filter, being careful rot to run it over, and catch the clear liquid, 
 filtrate, in a large test tube or beaker. If not clear run it through 
 again. To a little of the clear filtrate in a clean test tube add a few 
 
 * Rubber seal, pint fruit jar. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 7 
 
 drops of solution of silver nitrate. Now add a few drops of dilute ni- 
 tric acid. The white precipitate is silver chloride. The formation of 
 this white precipitate by silver nitrate in the presence of nitric acid is 
 a much used test for a chloride. Make solutions of potassium chloride 
 and potassium chlorate from the shelf bottles, add a little dilute ni- 
 tric acid to each, then silver nitrate. Which gives silver chloride? 
 What does your clear filtrate seem to contain? Evaporate a little of 
 it to dryness in a porcelain dish. Compare the appearance and taste 
 of residue with potassium chloride and chlorate from shelf. What 
 chemical change took place on heating potassium chlorate? Was the 
 manganese dioxide changed chemically on heating? For the use of 
 manganese dioxide as a "catalyzer" see (8). 
 
 7. Properties of Oxygen: Place about half a gram of sulfur in 
 a deflagration spoon, best covered with a bit of asbestos paper, ignite 
 the sulfur and lower it into a jar of oxygen. When the combustion 
 ceases, add a little distilled water to the jar and shake. Test the water 
 with blue litmus paper. The change to red shows acid formed by the 
 union of water with sulfur dioxide produced by the combustion. 
 
 Ignite a piece of charcoal on a clean spoon and lower into a jar 
 of oxygen containing a little water. Cover jar, shake and test the 
 water with blue litmus paper. A faint red should be given the paper 
 showing the formation of carbonic acid. Add to the jar clear lime 
 water* and shake. The milky appearance is due to calcium carbon- 
 ate, produced by calcium hydroxide in the lime water and carbon di- 
 oxide. 
 
 Place a little red phosphorus on asbestos paper in the spoon, light 
 it and quickly lower into oxygen. Test the water in the jar with blue 
 litmus paper. Phosphorus pentoxide is 'first formed and this unites 
 with the water forming phosphoric acid. What is an element? All 
 three of the substances burned are elements. These are non-metals, 
 and such on combustion usually form oxides that are -acidic or unite 
 with water to form acids. 
 
 Heat the end of a piece of picture wire, dip into powdered sulfur 
 for an instant. A portion of the S will adhere and burn. Lower it 
 into a jar of O. The sulfur burns and then the iron burns with a 
 shower of sparks. 
 
 Without removing the kerosene which adheres to it place a bit 
 of sodium as large as a grain of wheat in a spoon, ignite the kerosene 
 and lower into a jar of oxygen. The kerosene should burn and ignite 
 
 *If lime water is not present in the laboratory in quantity, it may 
 easily be prepared by shaking a little slaked lime in a jar of water 
 for some time, letting the lime settle and filtering. The funnel should 
 be placed in a flask to protect the filtrate from carbon dioxide in the 
 air. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 the sodium. Shake the bottle and test the water with red litmus pa- 
 per. The so-called alkali metals such as sodium burn and form oxides 
 which unite with water forming an alkali. 
 
 The burning of these elements forming compounds, oxides, illus- 
 trates chemical change, and chemical reactions of the type called di- 
 rect combination or synthesis. 
 
 8. Manganese Dioxide as a Catalyzer: In a test tube heat about 
 1 gram of potassium chlorate till it melts and begins to give off oxy- 
 gen. Without cooling drop a very little manganese dioxide upon the 
 fused salt. Is there much change in the rate of giving off oxygen? 
 Compare temperature required when oxygen is prepared from a mix- 
 ture of potassium chlorate and manganese dioxide, and from potassium 
 chlorate alone. The next experiment illustrates further the influence 
 of catalyzers, and also a very good method of preparing oxygen. 
 
 9. Oxygen from Sodium Dioxide (Peroxide) : In each of three dry 
 test tubes place about a gram of sodium dioxide. In one place also a 
 very little fine copper oxide and in another powdered manganese di- 
 oxide. In each of the three pour about 5c.c. of water and note rates at 
 which oxygen is given off. 
 
 To prepare a larger volume of oxygen by this method set up the 
 apparatus as in fig. 10, but use a dry flask instead of the test tuba. 
 Place in flask about 10 grams of sodium dioxide, then about half a 
 gram of powdered copper oxide or manganese dioxide, shake to mix 
 and spread the mixture evenly over the bottom of the flask. Fill the 
 funnel with water and by means of the pinch cock let it run in slowly 
 so as to maintain a suitable flow of oxygen. No heating is necessary. 
 The oxygen may be tested with a glowing splinter. Sodium hydroxide, 
 NaOH, is formed at the same time. Test a little of the solution in 
 the flask with red litmus or turmeric paper. Test sodium hydroxide 
 from the shelf with the same paper. 
 
 HYDROGEX. 
 
 11. (a). In as many test tubes place bits of magnesium turnings, 
 aluminium turnings or wire, zinc, iron filings, or other finely divided 
 iron such as steel wool; tin, copper, lead. To each tube add dilute 
 hydrochloric acid. In what cases is hydrogen, a colorless gas, given 
 off? In cases where it is not, add a few drops of concentrated hydro- 
 chloric acid. If some metals still resist warm the acid. 
 
 (b). In clean tubes and fresh metals try the same series with di- 
 lute sulfuric acid warming where necessary. In text-book refer to 
 "electro-motive series." From the more electro-positive end of the 
 series the metals give off H with decreasing readiness till hydrogen 
 itself is reached. The metals beyond or below H do not give H with 
 acids. That is, the metals that are more electro-positive than H dis- 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 place it from solution. There are secondary reasons why some met- 
 als above H do not evidently replace hydrogen. Lead is an example, 
 (c). For the action cf nitric acid on metals see 52. 
 12. Preparation of Hydrogen. Caution: Mixtures of hydrogen 
 and oxygen and hydrogen and air are dangerously explosive. In the 
 
 preparation of hydrogen in 
 this experiment and else- 
 where have no flame near the 
 delivery tube until all air is 
 expelled from the flask. . The 
 gas should be lighted the first 
 time with a test tube of burn- 
 ing gas. Should the stopper 
 of the flask be removed so as 
 to permit the entrance of air. 
 Fig. 3. the same precaution must be 
 
 used in lighting the gas the second time. 
 
 In a small flask fitted as in Fig. 3 place about 25 grams of granu- 
 lated zinc and about 50c.c. of dilute sulphuric acid. Add a few drops 
 of a solution of copper sulphate if the action proceeds too slowly when 
 sulfuric acid is used. To get the crystals mentioned below it is neces- 
 sary to u~e rrulphuric acid. In performing the experiment subsequent- 
 ly, inrtead of dilute sulphuric acid the zinc may be covered with water 
 and then strong hydrochloric acid cautiously added through the fun- 
 nel tube until the gas comes off freely. Collect the gas in jars reject- 
 ing the first jar full. Why? Collect two jars and then light the gas 
 with a test tube of burning gas. That is, hold the delivery tube up- 
 ward, and place over its end a small test or specimen tube. The hy- 
 drogen will rise and force the air downward. When the tube has had 
 time to fill move it mouth downward to a flame, which should be a 
 foot or two removed, and the H will light with a slight noise if pure. 
 Carry the tube back and lower it for an instant over the delivery tube. 
 Hold the flame inside the mouth of a dry bottle and continue until the 
 liquid deposited collects in drops. Taste it and test it with litmus pa- 
 per. What is it? In case the evolution of gas becomes too slow it is 
 due to exhaustion of the acid and the accumulation of zinc sulfate, or 
 chloride if hydrochloric acid was used. Merely adding more acid 
 does not suffice.' It is better to pour off the solution and replace it 
 with fresh acid as in the beginning. 
 
 Fill a jar by pouring hydrogen upward into it from another jar 
 and prove the presence of H in the jar by applying a flame to its 
 mouth. 
 
 Hold a jar of H mouth downward, remove cover and 
 insert a bit of lighted candle held on the end of a file or wire. Remove 
 
10 
 
 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 candle when it will light again. Repeat several times. Does a candle 
 burn in hydrogen? 
 
 When the evolution of gas in the flask has nearly ceased filter 
 the liquid into a dish. If sulfuric acid was used crystals of zinc" sul- 
 fate will appear when it becomes cold, or after it has been evaporated 
 down one half and cooled again. If hydrochloric acid was used eva- 
 porate to dryness to obtain zinc chloride. Place some of it on a glass 
 plate and observe at the next period. Is it hygroscopic? 
 
 13. Hydrogen from Water: (a) Sodium should be handled with the 
 rorceps or dry hands. Press a piece of the metal tightly into a 22 car- 
 cridge shell. Fill a test tube with water, drop the shell into the trough 
 d,nd collect the gas by displacement of water. Test the gas for hydro- 
 gen. 
 
 (b) Set up the apparatus as shown in fig. 4 using a half inch gas 
 pipe at least a foot long. Put near the middle of tube about 10 grams 
 iron turnings held in place by loose plugs of steel wool. To protect 
 the stoppers wrap cotton around the ends of the tube and keep it sat- 
 urated with water. Use a very small flask and only lOc.c. of water in 
 it. Heat strongly the middle of the iron tube and after several min- 
 utes boil the water in flask. After the air has been driven out of the 
 tube collect and test the hydrogen. Prevent a back flow by removing 
 delivery tube from the trough before removing burner from the flask. 
 
 Fig. 4. 
 
 The experiment may be continued till a large amount of the iron is 
 oxidized when it may be used in 14b; or, when cold more of the iron 
 may be "rusted" by wetting it and putting aside for a day or two. It 
 may then be heated and dried with a current of air before use in 14b. 
 
 In this experiment the reaction is shown from left to right: 
 
 3Fe-f4H 2 O=(reversibly)Fe 3 O4+4H 2 
 
 while in 14b it goes from right to left, if the contents of this tube or 
 magnetite is there used as provided for. 
 
 14. Hydrogen as a Reducing Agent: The experiment may be per- 
 formed as in (a) or (b). 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 II 
 
 (a) Place in a piece of hard glass 
 tubing a column of granulated cop- 
 per oxide using loose plugs of as- 
 bestos to hold the oxide in place. 
 The tube may be drawn out and 
 bent at right angle as shown in 
 Fig. 5 or a small right angled tube 
 may be connected with its outer 
 end by rubber. Connect with the 
 flask which contains zinc, and add 
 dilute sulphuric acid. Turn the 
 small tube upward and determine 
 when the air is wholly expelled by 
 lighting the hydrogen with a test 
 
 tube as in Ex. 12. Now turn the tube downward, place it in a test tube 
 surrounded with water and heat the copper oxide, beginning at the end 
 of the column nearer the flask. Continue till the copper oxide is all 
 reduced, that is, all changed to red copper. Examine the liquid that 
 has collected in the test tube. What is it? 
 
 (b) For the glass tube and copper oxide in (a) substitute the 
 iron tube and contents from (13b) or use the iron tube and about 5 
 grams of ferric oxide. First fill the tube completely with hydrogen as 
 proved by lighting it with a test tube, heat the oxide strongly, protect- 
 ing the stoppers with the wet cotton as stated in 13b. 
 
 With a moderate flow of hydrogen continue the reduction 10 to 15 
 minutes. What is the source of the water in the test tube? When 
 cold and if ferric oxide was used test some of the contents of the tube 
 in con. HC1 warming. What gas is evolved? What evidence have you 
 that the reaction in 13b is reversed? What is a reversible reaction? 
 What governs the direction of this one? How could you produce a 
 state of equilibrium in it? What is the meaning of equilibrium? 
 
 WATER. 
 
 15. Solubility of Solids: Some substances are very soluble in 
 cold, still more soluble in hot water. Shake 3 grams of powdered so- 
 dium nitrate or ammonium chloride with 5c.c. of water in a test tube 
 till all is dissolved. Is there any change in temperature of the water? 
 Now add 3 grams more and shake. Shake and determine whether it 
 all dissolves. If not, heat till dissolved. What occurs on cooling to 
 room temperature? 
 
 Some substances are very soluble in hot water, slightly in cold 
 water, and such may easily be purified by crystallization from hot so- 
 lution. By heating dissolve 4 grams of ordinary potassium chlorate 
 in lOc.c. distilled water. Filter boiling hot if there is any turbidity or 
 
12 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 solid matter suspended in the liquid. When the solution is cold filter 
 off the crystals, setting filtrate aside and wash crystals with a little 
 cold water, dissolve a portion of them in distilled water and add sil- 
 ver nitrate to the solution and to filtrate. Compare the amount of 
 precipitate obtained with silver nitrate in the two tubes. It is due to 
 a chloride, commonly found in ordinary potassium chlorate. 
 
 Some substances are little more soluble in hot water than in cold 
 wau-r. Shake 5 grams of common salt with lOc.c. of water in a test 
 tube. Note the amount of salt undissolved. Boil the solution for a few 
 moments to saturate the water with salt. Note again, the amount of 
 salt. Filter the solution boiling hot and cool to room temperature. 
 Does much salt crystallize out? Why not? Compare the solubility 01 
 salt in hot and cold water with that of potassium chlorate. A few sub- 
 stances are even less soluble in hot water than in cold. Examples are 
 calcium sulfate (gypsum) and slaked lime. 
 
 16. Solubility of Liquids in Liquids, and the Separation of Solutes 
 between two non-miseible Solvents: Measure accurately in a cylinder 
 about 50c.c. of water reading at the lower surface of the meniscus. Now 
 carefully pipet into the cylinder 25c.c. of common alcohol. Mix thor- 
 oughly and read the volume. Is the volume now the sum of those of 
 the alcohol and water? Do water and alcohol dissolve each other 
 completely? 
 
 In a test tube place about lOc.c. of water and about 2c.c. of chloro- 
 form and shake. Let stand and note whether they are mixed. Add a 
 crystal of iodine and shake for some time and let stand. Which li- 
 quid takes up most of the iodine? 
 
 As above try to mix lOc.c. of \vater and 2c.c. of carbon disulfide. 
 Add about oc.c. of bromine water and shake. Which liquid takes up 
 most of the bromine? 
 
 To another tube add water and a little benzene, and a small crystal 
 of potassium permanganate and shake. Which liquid takes up the 
 permanganate? 
 
 Determine whether alcohol and chloroform, benzene and chloro- 
 form, carbon xdisulfide and benzene will mutually dissolve each other. 
 
 17 Solubility of Gases: Gases also vary widely in their solu- 
 bility in water. In text-book see solubility of oxygen, nitrogen and hy- 
 drogen sulfide which do not act chemically with the water, dissolve 
 only moderately and obey Henry's law, (which see). Others which 
 generally act chemically with water dissolve in very large amounts. 
 
 In a test tube or flask heat hydrant water and observe gas bub- 
 bles given off before the water begins to boil. Try the same with dis- 
 tilled water. Does it contain dissolved gases. Why does water when 
 drawn from the hydrant sometimes look milky and quickly become 
 clear on standing? 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 13 
 
 Note carefully the odor from a solution of ammonia due to am- 
 monia given off. Pour about Ic.c. of the solution into lOc.c. of water 
 in a dish and test with turmeric paper. Boil till the water is half 
 evaporated and test again with the paper. What is the effect of heat 
 on the solubility of gases. To a degree the following shows one of 
 many exceptions to the general rule: 
 
 Boil lOc.c. of dilute hydrochloric acid till three-fourths cf it has 
 evaporated. Boil the same volume of concentrated hydrochloric acid 
 till one half has evaporated. Add to each sample of boiled acid wheji 
 cold, the same amount of granulated zinc. Do they seem to act on zinc 
 at about the same rate? A more accurate determination would show 
 that they have the same concentration. By boiling long enough 
 samples of dilute and concentrated hydrochloric acids one arrives at 
 the same result; namely acids of concentration 20.2 per cent. 
 
 18. Chemical Action of Water: Refer back to 13 , for the action 
 of water on sodium and on iron, and to 9 for its action on sodium di- 
 oxide. 
 
 Upon a piece of quick lime drop water slowly until the water is no 
 longer absorbed and the piece of lime looks wet. Place it in a dish 
 and note what occurs. This is the familiar "slaking" of lime. 
 
 10. Water in Combination: In a test tube heat a small amount of 
 copper sulfate, observing water given off and changed appearance. In 
 the same way try borax, alum, sodium phosphate. These are "hyd- 
 rates" and the water they contain is called "water of crystallization." 
 For contrast try potassium sulfate and common salt. 
 
 In a weighed porcelain crucible with lid weigh accurately about 4 
 g. of barium chloride, place the crucible on a triangle over a burner 
 and heat ten minutes. When cold weigh the crucible and contents 
 again and from the first weight of the barium chloride and the loss on 
 heating, calculate the per cent, of water. 
 
 Efflorescence: On a glass plate or watch-glass place a few crys- 
 tals of sodic sulphate, expose to air till the next laboratory period. 
 
 Deliquescence: Expose to air in dishes small pieces of calcium 
 chloride and caustic potash till the next laboratory period and record 
 results. 
 
 20. Electrolytic Decomposition of Water: Support the U tube 
 shown in fig. 14 with a clamp and use the current described in 53. 
 Connect the side-arms of the tube with the trough of water by means 
 of short delivery tubes. Fill the U tube nearly to the side arms with 
 a 5th normal solution of sodium sulfate already made up and add a few 
 drops of litmus solution to each side of U tube. Start the current and 
 at the same time place over the ends of the delivery tubes two test 
 tubes the same size and full of water, thus collecting the hydrogen and 
 oxygen set free. Continue till the smaller volume equals about 5c.c. 
 
14 
 
 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 and compare volumes. Note color of the solution at each electrode and 
 state cause. Name electrodes. From which comes the 0, and the H? 
 
 21. Purification of Water by Distillation: To a small volume of 
 hydrant water add a few drops of barium chloride solution. The white 
 precipitate shows carbonate and sulfate radicals present. Now add 
 dilute hydrochloric acid. The precipitate remaining shows sulfate 
 radical. To another portion add a few drops of dilute nitric acid and 
 a few of silver nitrate. A white precipitate shows chloride present. 
 
 Set up apparatus shown in fig. 6, fill flask one-third full of water, 
 distill enough to clean the tubes and reject it. Distill about lOc.c. of 
 the water and test it for sulfate and chloride radicals. None should 
 be obtained, all ordinary mineral matter being left in the boiler in dis- 
 tillation. 
 
 Try the action of ammonia and so- 
 dium hydroxide on turmeric paper 
 and dilute siilfurio aoi^ on blue lit- 
 mus papei To one-fourth- of 
 flask full of water add about 5c.c. of 
 ammonia, distill and determine wheth- 
 er any ammonia distills over. Clean 
 apparatus carefully and try sodium 
 hydroxide, adding a few c.c. to one- 
 fourth flask of water. Distill from 
 a fresh portion of water containing 
 a little sulfuric acid and a little 
 potassium permanganate. Do they go 
 Fig. 6. over? 
 
 HYDROGEN DIOXIDE (PEROXIDE). 
 
 22. Preparation: Place a beaker with lOOc.c. of water in water, 
 preferably ice cold. Stir in little by little 5 grams of sodium dioxide, 
 NaO 2 the operation taking about 5 minutes. Even then some of the 
 peroxide will be decomposed giving off oxygen. Now add gradually 
 in the same way dilute hydrochloric acid till a drop of the liquid taken 
 out with the stirring rod just turns blue litmus paper red. This gives 
 a solution of hydrogen dioxide, but it contains also common salt. It 
 may be used where hydrogen peroxide is required below save where it 
 is used with silver and lead. There the commercial peroxide should be 
 used. 
 
 Rub in a mortar with a little water about half a gram of starch, 
 transfer to a dish or beaker, add about lOOc.c. of water and a few crys- 
 tals of potassium iodide and heat to boiling. This is known as "starch- 
 iodide solution" and filter paper wet with it is called "starch-iodide pa- 
 per." Each is frequently required. To a part of the solution add a 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 15 
 
 few drops of the solution of hydrogen peroxide which sets free iodine 
 and this colors the starch blue. To see the color by transmitted light 
 dilute with much water. This is used as a test for iodine or hydrogen 
 peroxide or starch. That is, two being known to be present the pres- 
 ence or absence of the third can be determined by the test. 
 
 23. Oxidation with Hydrogen Peroxide: Moisten a strip of filter 
 paper with very dilute lead acetate and expose it to hydrogen sulfide 
 a little of which may be made in a test tube as in 68. The black sub- 
 stance is lead sulfide, PbS. Pour upon the paper a few drops -of hy- 
 drogen dioxide, which will change the lead sulfide to white lead sul- 
 fate, PbSO*. Why is this called oxidation? 
 
 To a solution of silver nitrate add NaOH and then carefully add 
 just enough of a solution of ammonium hydroxide to dissolve the pre- 
 cipitate at first formed. Now add commercial hydrogen peroxide. The 
 gray precipitate is finely divided metallic silver, and the escaping gas 
 is oxygen which may be tested by trying in the tube a glowing splint- 
 er. This action appears to be one of reduction, but it is not primarily. 
 Probably a higher oxide of silver is formed and at once decomposes 
 into silver and oxygen. The next two cases are of the same sort. 
 
 To about a gram of manganese dioxide in a test tube add hydrogen 
 dioxide and test the gas with a glowing splinter. Repeat using a con- 
 centrated solution or a few crystals of potassium permanganate in- 
 stead of the manganese dioxide. 
 
 CHLORIDE. 
 
 24. Preparation of Chlorine: Chlorine is dangerous if breathed. 
 Experiments with it should be conducted in hoods. If carried out on the 
 students' desks only half the usual number should work in the room at 
 one time and windows should be freely opened. They should stand to 
 windward of the apparatus when collecting the gas. When through 
 collecting at once place the delivery tube in a test tube nearly full of 
 concentrated sodium hydroxide, and remove the source of heat. Be- 
 fore beginning the experiments students should read them through, 
 the right plan in any case, provide everything necessary and thus re- 
 duce the time to the minimum. Make a solution of starch-iodide as di- 
 directed in 22. , 
 
 Chlorine is made commercially by the electrolysis of fused com- 
 mon salt or a solution of common salt. All other methods depend up- 
 on the oxidation of the hydrogen of hydrochloric acid. Several 
 methods may be illustrated on a very small scale. 
 
16 
 
 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Upon a few crystals of potas- 
 sium permanganate in a test tube 
 pour about Ic.c. of con. hydrochlo- 
 ric acid and pour a little of the 
 heavy, greenish yellow gas into a 
 little of the starch solution in an- 
 other tube. In the same way treat 
 a little potassium dichromate heat 
 and test for chlorine. Try in the 
 same way lead dioxide and con. 
 hydrochloric acid. Use a very 
 little potassium chlorate and con- 
 centrated hydrochlpric acid, HC1; 
 also, about 1 cc. con. HC1 and a 
 few drops of con. nitric acid. 
 
 Fig. 7. 
 
 25. Set up the apparatus as in fig. 7, with the water bath one- 
 fourth full of water. A copper can is best, but a tin can or a beaker will 
 serve for a water bath. See that the thistle tube reaches nearly to the 
 bottom of the flask. In the flask place 25 grams of manganese diox- 
 ide MnO 2 , and add 40c.c. of concentrated HC1 diluted with lOc.C. of 
 water. Heat the flask and collect five jars or bottles of chlorine. Tlie 
 green color will show when they are full. They should be well filled, 
 but a large excess should not over-flow into the room. One jar should 
 have the bottom wet with con. sulfuric acid to dry the gas for use in 
 bleaching. 
 
 When through collecting place the delivery tube in NaOH in a 
 test tube. Let the flask cool, removing water bath, and proceed with the 
 next experiment. 
 
 26. Properties of Chlorine: To show bleaching action and the 
 need of water, suspend in the jar of dry chlorine strips of colored 
 cotton cloth and litmus paper. After a few moments note any fading 
 of the colors, then moisten the strips and suspend again in the jar, 
 and note effect. Refer to a text-book for information on hypochlorous 
 acid and bleaching with chlorine. What does the bleaching? 
 
 Into a jar of chlorine pour successively with shaking, small vol- 
 umes of solutions of litmus, cochineal, much diluted ink. 
 
 From a piece of antimony scrape with a knife a very little of the 
 metal letting it fall into a jar of chlorine. Using forceps or tongs heat 
 to redness a strip of copper foil and lower it into the same jar of chlo- 
 rine. If "Dutch metal" is used it need not be heated. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 17 
 
 Burn the laboratory gas at the end of a glass tube and lower the 
 small flame into a jar of chlorine. What is the black substance? When 
 the green color has disappeared blow breath over mouth of jar, which 
 will give a fog, consisting of droplets of water containing HC1. Will 
 carbon burn in Cl? Try an ignited piece of charcoal. 
 
 On a deflagration spoon lower a small bit of white phosphorus in- 
 to a jar of Cl and avoid inhaling the Cl or fumes of PC1 3 . It should 
 soon melt, then take fire. 
 
 Now take the flask and all jars to the sink best under hood and 
 standing well back fill them with water. Wash well any remaining 
 MnOi. and place it in a vessel provided for that purpose. 
 
 HYDROCHLORIC ACID. 
 
 27. Preparation : Read the experiment through and have all the 
 necessary materials at hand so that once begun the experiment may 
 be carried through rapidly. ^ 
 
 All soluble chlorides give hydrochloric acid when treated with 
 concentrated sulfuric acid. For many reasons sodium chloride, com- 
 mon salt, is to be preferred. 
 
 Set up the apparatus as in fig. 7, omitting the water bath. In the 
 flask place 20 grams of sodium chloride. Dilute 35 grams (20c.c.) of 
 concentrated sulfuric acid by pouring it slowly with stirring into 8c.c, 
 of water in a beaker. Pour slowly into the flask and let stand a few 
 moments till acid and salt are in contact throughout then apply a low 
 heat, best using a burner with crown top. Collect the gas in dry jars 
 or bottles in the same way as chlorine. Abundant fumes will indicate 
 when the jars are full. After collecting two jars, fill a dry bottle 
 which has been fitted with a stopper and a short piece of tubing one 
 end flush with the large end of stopper and the other drawn out and 
 cut off so as to leave a small orifice. It should reach at least to the 
 middle of the bottle's length. Insert stopper when full and place the 
 bottle mouth downward in water. Press down the bottle and pour cold 
 water over it to start the absorption. A fountain will result. 
 
 Into a test-tube of water insert the delivery- tube so that it reaches 
 a very little below the surface of the water. Note ready absorption of 
 the gas. Is the solution lighter or heavier than water? Lower the 
 tube as necessary to absorb all the gas. Why does the water become 
 warm? Continue till the water is nearly saturated, then try the action 
 of small portions of the solution on a little zinc, marble, sodium car- 
 bonate. Compare its action with that of the dilute HC1 from the shelf. 
 Test its action on blue litmus paper. To a little of the solution add a 
 few drops of silver nitrate then a few drops of dilute nitric acid. The 
 precipitate is silver chloride. This is a test for hydrochloric acid or 
 a chloride. Repeat using instead of HC1 a few drops of a solution of 
 
18 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 salt or of any other chloride. With a known chloride could it be used 
 as a test for silver? Save flask and its contents for 29. 
 
 28. Acid, Base, Salt, Neutralization: Wet the inside of a jar with 
 a solution of concentrated ammonia and pour out excess of liquid. 
 Cover jar with a glass plate. Place^ it over a jar of the HC1 gas from 
 the last experiment, bringing the jars mouth to mouth and remove 
 plate. The ammonia unites with acid forming ammonium chloride, a 
 salt. 
 
 Standing well away drop one or two small bits of sodium into a 
 little distilled water in a bottle. When the action is over test the wa- 
 ter with turmeric paper. The action of sodium on water gives sodium 
 hydroxide, a base, and what gas? (see 13). 
 
 Test sodkim hydroxide from shelf bottle with the papers. Pour 
 about 5c.c. of the alkali into a dish and add dilute hydrochloric acid till 
 the solution turns litmus paj>er red, testing by taking out a drop of the 
 solution with a stirring rod and touching the paper ; never place pa- 
 pers in the solution or dip them into it. Now add a few drops of NaOH 
 or dilute HC1 as may be necessary with stirring, till the solution 
 changes neither turmeric nor blue litmus paper. It is now neutral. 
 Evaporate to dryness, taste the residue. Place upon it a few drops of 
 con. sulfuric acid and note odor of the gas. What was the solid resi- 
 due? 
 
 29. The preparation of HC1 from salt and sulfuric acid gives a 
 good example of a reversible reaction: 
 
 NaCl+H 2 SO 4 = (reversibly) HNaSO 4 +HCl. 
 
 HC1 is a stronger acid than H 2 SO4. In the cold or in a small closed 
 space even on heating the reaction would not complete itself to the 
 right. But, the HC1 is easily volatile and heating drives it out of 
 solution and away so that it cannot react with HNaSO 4 to the left. Sul- 
 furic acid is volatile only at very high temperature. Try restoring 
 HC1 thus: Dissolve with the least volume of water and heating, the 
 contents of the flask used in preparing HC1. Cool some of the solution 
 of HNaSO4 in a test tube and add a few c.c. of con. HC1 from shelf, 
 which will precipitate sodium chloride. Why is sulfuric acid used to 
 prepare easily volatile acids from their salts? Why are the reactions 
 completed when heat is used? 
 
 30. Preparation of Bleaching Powder, Potassium Hypochlorite 
 and Potassium Chlorate: Study these subjects in a text-book and read 
 the experiment through. 
 
 Arrange the apparatus as shown in fig. 8. The wash bottle con- 
 tains diluted sulfuric acid, made by adding 40c.c. of the con. acid with 
 stirring, to lOc.c. of water. Its purpose is to remove the greater part 
 of the HC1 and water from the Cl. The horizontal test tube contains 
 about 2 grams of dry slaked lime spread evenly throughout i r s 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 19 
 
 O 
 
 length. The upright test tube should contain a solution of 4 grams of 
 potassium hydroxide dissolved in 12c.c. of water. It must be cooled 
 and kept cold by surrounding with water as shown. 
 
 Charge the flask with MnO 2 and diluted HC1 as in 25. Maintain a 
 moderate stream of chlorine for about 15 minutes or until the larger 
 
 portion of it seems to pass 
 through the solution in the second 
 test tube. Now preserve half the 
 contents of this tube as potassium 
 hypochlorite, and heat the re- 
 
 j, |>| mainder to boiling and without 
 
 ( cooling continue to pass chlorine 
 
 into it for about five minutes or 
 till a drop taken out with a stir- 
 ring rod does not feel soapy to the 
 fingers. Now heat to boiling and 
 filter the solution. Cool by stand- 
 ing the tube in water, when crys- 
 tals of KC1O 3 form. When quite 
 cold filter them off and wash with 
 a little cold water. Test the ni- 
 trate with silver nitrate. What 
 was formed * besides potassium 
 chlorate? Dissolve the chlorate 
 Fig. 8. by passing about 3 cc. of boil- 
 
 ing water through the filter several times, cool, let crystallize, pour 
 off the water from the crystals, dissolve them in water and add 
 silver nitrate. Compare the first and second precipitate formed by sil- 
 ver nitrate. Pure chlorate would give no precipitate. 
 
 Into a little of the hypochlorite solution place a drop of diluted 
 ink, into another .portion a bit of colored cloth. After noting any 
 bleaching add a little dilute acid to each solution and note effect. To 
 a third portion add a few drops of strong solution of ammonium hyd- 
 roxide. What gas is given off? 
 
 Try the action of dilute acid on a little of the dry bleaching pow- 
 der from the horizontal tube. Dissolve a portion of it so far as pos- 
 sible, filter and try the action of the filtrate on ink, colored cloth, lit- 
 mus paper, before and after adding dilute acid. 
 
 To show the instability and oxidizing power of potassium chlorate 
 mix on paper 5 grams of the salt and 5 grams of powdered sugar, 
 but do not grind them together in a mortar. Place the mixture on an 
 iron plate, take out a drop of con. H 2 SO 4 with the stirring rod, stand 
 well away and drop the acid upon the mixture. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 BKOMENE 
 
 IODOE. 
 
 31. Preparation of Bromine: Place 5 grams of manganese diox- 
 ide and 5 grams of sodium bromide on paper, hold the neck of the re- 
 tort pointing slightly upward and slide the mixture in at the tubulus 
 without letting it fall into the neck of the retort. Support the retort 
 as in fig. 9, add 50c.c. of dilute sulfuric acid through a funnel. The 
 flask should contain about lOOc.c. of water and the tip of the retort 
 neck should dip under its surface. Apply heat and continue till all 
 the bromine has distilled over. Move the flask away so as to bring 
 the neck of the retort above the surface and then remove the burner. 
 Is bromine soluble in water? Is it heavier than water? Set aside the 
 
 flask containing bromine for lat- 
 er use. 
 
 Give two reasons why we do 
 not prepare Br in the same way 
 as Cl; that is, by the action of 
 hydrobromic acid on manganese 
 dioxide? 
 
 May Cl be prepared by the 
 same method used for Br; that 
 is, by use of NaCl. MnO 2 and di- 
 luted H 2 SO 4 ? Try it in a small 
 way in a test tube, but strength- 
 en the acid by adding about Ic.c. 
 con. H,S0 4 to 2c.c. of dilute acid. 
 Write equations for preparation 
 of eaph. 
 
 ~~ Why is it not best to use con. 
 
 Fig. 9. sulfuric acid in preparing either 
 
 Br or Cl by this method? 
 
 31. By the same method prepare a little iodine, using about O.r> 
 gram of sodium iodide, a gram of MnO 2 and 5c.c. of a mixture of di- 
 lute and con. suifuric acid. Note sublimed iodine near the mouth 
 of the test tube. Compare the reaction with that in the preparation 
 of Cl and Br by the same method. 
 
 Heat a few crystals of iodine in a dry test tube. Does it form 
 a liquid before subliming? What is sublimation? Note crystals 
 higher up on walls of the tube. 
 
 33. Hydriodic Acid and Comparison of HCI, HHr, HI: Treat 
 very small amounts of sodium chloride, sodium bromide and sodium 
 iodide with a few dropi of con. sulfuric acid. How did you prepare 
 HCI? Could HBr be prepared in the same way? What is the black 
 substance set free by the action of the acid on sodium iodide? 
 
LABORATORY MANUx\L OF GENERAL CHEMISTRY 21 
 
 When salt and sulfuric acid are heated together hydrochloric 
 acid is set free, but no chlorine. Though H 2 SO4 may be reduced and 
 HC1 oxidized, they are too stable to act on each other. 
 
 Though statements to the contrary are common, hydrobromic 
 acid may be prepared in precisely the same way as you prepared 
 HC1, remembering that the sulfuric acid there used was somewhat 
 diluted. A very little HBr is oxidized by the H 2 SO 4 giving a trace 
 of free bromine and sulfur dioxide. Write the equation. But when 
 the gas is absorbed with water and the solution is distilled, the 
 trace of Br quickly passes off, and the sulfur and its compounds 
 that may be liberated are oxidized to sulfuric acid which remains 
 to the last in the distilling vess.el. 
 
 Hydriodic acid cannot be prepared by this method, since the 
 HI is almost completely oxidized to free iodine and water, and the 
 sulfuric acid is reduced to sulfurous acid and even to hydrogen sul- 
 fide. 
 
 Compare the degrees of stability of HC1, HBr, HI, and compare 
 these with their heats of formation by consulting a reference book. 
 
 34. Preparation of Hydriodic Acid: Arrange a test tube as in 
 fig. 10 high enough to permit heating and having a right angled de- 
 livery tube. Place in the dry test tube 5 grams powdered iodine and 
 on top of it 0.5 red phosphorous, shake to mix, put tube in place and 
 warm till the P and I react. When cool place delivery tube in 
 small flask or bottle and by means of the pinch cock drop into test 
 tube about 20 drops of water. Warm gently and collect the receiver 
 full of HI and stopper it. Now place the delivery tube in water in 
 test tube to absorb the gas, but have the tip of delivery tube just 
 above the surface of the water. 
 
 Prepare a little Cl in a test tube by warming a few crystals of 
 KC1O:; and a little con. HC1. Pour some of the Cl into the flask con- 
 taining the HI. What is set free? 
 
 Test the solution of HI in water with silver nitrate. What was 
 the action of the P and I? What was the action of this compound 
 and the water? What is hydrolysis. Write all equations. 
 
 This same method with some modifications is often used for 
 the preparation of hydrobromic acid. (See the text book for de- 
 scription.) 
 
 35. Powder a few crystals of iodine in a mortar, place it 
 in a test tube half full of water, and pass into the tube hydrogen 
 sulfide gas prepared as in 68. Have the delivery tube reach quite 
 to the bottom of the test tube. When all the iodine has disappeared 
 boil the solution for a time to expel the excess of hydrogen sulfide 
 and filter, several times if necessary to get rid of all the sulfur. 
 
22 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Test the filtrate with blue litmus paper. Set the solution aside f.o 
 be used in the next experiment. Hydrogen sulfide here acts as a 
 reducing agent. Write the equation. 
 
 :W. Properties of Chlorine, Bromine, Iodine and Their Com- 
 pounds: Try to dissolve a little powdered iodine in water. Now 
 pour off most of the water, add a few crystals of sodium iodide and 
 shake. It is supposed the solution now- contains NaI 3 . 
 
 Try the solubility of I by pouring about Ic.c. of carbon disul- 
 fide upon a crystal of I in a test tube and shaking. Now add water 
 to the tube, shake and let carbon disulfide settle. In which liquid 
 is most of the iodine? Try the solubility of iodine also in alcohol 
 and in chloroform. Save the solutions of iodine. 
 
 Make a solution of starch by boiling about half a gram of 
 starch in a dish or beaker half full of water and with stirring. 
 Pour a little of the starch solution into a beaker of water, add a 
 little iodine solution. This is a good test for free iodine. Add a 
 little sodium iodide solution and a little of the starch solution to a 
 bleaker of water. To one portion add a little bromine water and to 
 the other chlorine water. Is iodine set free by the Cl and the Br? 
 To a solution of sodium bromide add chlorine water. Is bromine 
 set free? To concentrate the bromine and also to show its solu- 
 bility, to the liquid in the test tube add a little carbon disulfide, 
 shake and let the latter settle. 
 
 Arrange the three halogens in the order of the ability of each 
 to replace the others, and compare this order with that of the sta- 
 bility of their compounds with hydrogen. 
 
 To a little solution of iodine and to blue starch-iodine solution 
 add a solution of sodium thiosulfate till the colors disappear. 
 
 To show the reducing power of hydriodic acid make very di- 
 lute solutions of potassium dichromate and potassium permangan- 
 ate, add a little dilute sulfuric acid and then some of your solution 
 of hydriodic acid. The chromate and permanganate are reduced 
 and the solution is colored brown by iodine. For the same purpose 
 treat a little of a solution of sodium iodate, NaIO 3 , with dilute H 2 S0 4 and 
 add some of your solution of HI. Refer to a text book and write 
 the equations. 
 
 37. Tests for the Halogens: How may each of the halogens in 
 the free condition be recognized and tested for? (See 36). The 
 following applies to them in the form of soluble halides : 
 
 To three tubes containing respectively a little dilute HC1 or so- 
 lution of any chloride, solution of any bromide, solution of any 
 iodide, add a few drops of silver nitrate. Compare the colors of 
 the three silver halides. Try to dissolve a little of each with am- 
 thiosulfate. Try other portions with dilute nitric acid. Dissolve a 
 
LABORATORY MANUAL OP GENERAL CHEMISTRY 23 
 
 aionium hydroxide. Try other portions with a solution of sodium 
 little of each with, potassium cyanide, KCN, being careful not to get it 
 on the hands. 
 
 38. Hydrofluoric Acid: Slowly and evenly heat a glass plate 
 over a burner flame. Rub over the surface a piece of bees-wax so 
 as to make a continuous film of the melted wax. When the wax has 
 hardened, write on the plate with a pointed file or knife, cutting 
 4uite through the wax. In a lead dish place a little ammonium 
 fluoride, or powdered calcium fluoride, moisten with con. sulfuric 
 acid. Cover the dish with plate, wax side down, and apply a low 
 heat to the dish so as not to melt the lead. After a few moments, 
 remove the wax and examine the glass. 
 
 To test for a fluoride place a little of it mixed with a little 
 sand, in a dry test tube, drop upon the mixture a little con. H 2 SO 4 
 and heat. While heating hold in the tube a wet stirring rod. Where 
 wet the rod will be covered with a jelly-like layer of silicic acid 
 Why cannot a solution of H 2 F 2 be kept in a glass bottle? 
 
 EQUIVALENT WEIGHTS. 
 
 39. Equivalent of Zinc: Set up the apparatus as shown in fig. 
 tO and prove that all joints are tight by placing water in the funnel 
 covering the end of the delivery tube with the wet finger and open- 
 ing the pinch cock. The water should not run down. The end of 
 the delivery tube should be bent upward and securely placed under 
 the jar. Use pure, bright granulated zinc. 
 
 Fig. 10. 
 
 Weigh with great care not less than 0.2 grams nor more than 
 0.25 grams of the zinc for every lOOc.c. that the jar or bottle will 
 hold. If an analytical balance is used one weighing will suffice. 
 
^ I LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Horn pan balances are not very accurate and rarely in exact equi- 
 librium. In this case weigh the zinc on one pan, then on the other 
 and take half the sum of the weights. What errors in the balance 
 will this practically eliminate? Place zinc in the test tube, fill 
 collecting jar completely with water, and add 15c.c. of con. HC1 to 
 the funnel. Let in all the acid and close the clamp. When all the 
 Zn is dissolved bring the level of the water in the jar to that in the 
 trough and securely place on the cover. Dry the outside of the jar 
 and weigh on the platform balance. Pill it completely with water 
 and weigh again. The difference in grams less the volume of acid 
 gives the volume of the hydrogen in cc. (V). Find the temperature of 
 the water in the trough (t), and the reading of the barometer (P). Cal- 
 culate the volume (V) to normal conditions by use of the formula, 
 V'=VX273X(P aq. tens, at t)-^(273-H) X760. Why? The weight 
 of H=V'X. 00009. The equivalent of zinc equals its weight divided by 
 the weight of HX 1.008, and its atomic weight equals twice this value. 
 
 40. Equivalents of Other Metals: Find the equivalent of one or 
 more of the following, using not more than the weights given for each 
 100 c.c. that the receiving vessel holds. Aluminium 0.07 gram; mag- 
 nesium, 0.08; .iron, 0.15. In the cases of aluminium and magnesium 
 it is better to place the test tube in a beaker of water to keep the tem- 
 perature down, and to use dilute acid. In the case of iron it may be 
 necessary to warm the con. HC1 used. Collect the hydrogen and pro- 
 ceed precisely as in the previous experiment. 
 
 41, Equivalent of Chlorine (a) : In normal times porcelain Gooch 
 crucibles are cheap, and if they are available this method should be 
 used; otherwise use (b). 
 
 Weigh accurately about 0.5 gram pure silver, preferably foil, place 
 in a beaker, add lOc.c. water and 5c.c. pure nitric acid, and cover with 
 clock glass. Warm if necessary, and after all is dissolved boil gently. 
 Wash under side of glass and the inner surface of beaker till beaker 
 contains about 75c.c. liquid, then add 20c.c. dilute hydrochloric acid 
 and stir. Set in dark place till ready to filter. In a clean Gooch cru- 
 cible make a mat of asbestos with aid of suction of the filtering pump 
 as shown by the Instructor. It should be thick enough so that you 
 cannot see light through it, dry for half an hour in oven at about 140 
 degrees, cool, in desiccator or on a clean surface and covered, twenty 
 minutes and weigh. 
 
 With suction pump filter off silver chloride into crucible as shown, 
 heat in oven one hour at about 140 degrees. Desiccate as before and 
 weigh. Find weight of silver chloride and thence chlorine, and cal- 
 culate the equivalent of chlorine if the equivalent of silver is 107.9. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 25 
 
 The equivalents of bromine and iodine may be determined in the 
 same way. Indirectly many other equivalents may be determined by 
 the method as the following problem illustrates: 
 
 If 3 grams potassium chloride be treated with an excess of silver 
 nitrate and the silver chloride weighs 5.773 grams, find the equivalent 
 of potassium, if that of chlorine is 35.5. 
 
 (b) Weigh accurately a small porcelain dish, place in it about 
 0.5 gram of silver foil and weigh again. Add to dish 5c.c. con. nitric 
 acid diluted with lOc.c. of water, cover with a watch glass. Give it 
 time and do not heat unless necessary. When all the metal is dis- 
 solved, with a small amount of water in a fine stream from the wash 
 bottle wash any spattered liquid from the under side of the watch 
 glass, into the dish. Add to the dish 20c.c> pure dilute HC1 and eva- 
 porate the liquid to complete dryness on a water bath. Heat the dish 
 in an oven for half an hour at about 125, or heat some distance above 
 the burner flame till the silver chloride begins to melt. When fully 
 cold weigh accurately the dish and contents, subtract from the weight 
 of the silver chloride the weight of the silver and find the equivalent 
 cf chlorine calling that of silver 107.9. 
 
 42. Equivalents of Copper and Other Heavy Metals : The equiva- 
 lents of copper and several other metals may be determined Toy con- 
 verting the weighed metal into nitrate with nitric acid, and decom- 
 posing the nitrate by heat, leaving the oxide. 
 
 Use accurately weighed copper foil or clean copper turnings, and 
 proceed the same as in 41 (b), but use no hydrochloric acid. Evapor- 
 ate to complete dryness, and heat high over a flame' with the watch 
 glass on the dish. Gradually lower the dish and when nearly all blue 
 color has disappeared, or all evidence of steam in the case of other 
 metals, remove the watch glass and apply the full capacity of the 
 burner for half an hour when there should remain a layer of black 
 copper oxide. Material spattered upon the glass must be washed back 
 into the cooled dish and the liquid must be evaporated again on the 
 water bath, and the dish strongly heated. Weigh when quite cold. 
 
 Prom the weight of the copper oxide subtract the weight of the 
 copper, and find the equivalent of copper calling oxygen 8. If 16 is 
 used for oxygen the number obtained for copper is its atomic weight. 
 
 MTROGEtf A1VD ITS COMPOUNDS. 
 
 43. Preparation of Mtrogen From the Air: Fill a pneumatic 
 trough with water till it rises, about % inch over the shelf. Place up- 
 on the shelf a small crucible or cupel and place in it as much red 
 phosphorus as will lie on a half inch of the end of a knife blade or 
 spatula. Ignite the P and at once place over the vessel and upon the 
 shelf a jar or wide mouth bottle. Let the jar remain till it is cool and 
 
26 
 
 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 (he white fumes have been absorbed. Note the height to which the 
 water has risen and estimate the ratio of the volume of oxygen which 
 has been consumed to the nitrogen which remains. Slip cover on the 
 jar, place it upright and test the gas remaining with a burning splint- 
 er. What other gases are mixed with the N thus obtained? Why does 
 not this method give accurate results as to the volume of oxygen in 
 air? 
 
 4-1. More Accurate Determination of Oxygen in Air: Arrange the 
 apparatus as shown in fig. 11. When all is adjusted, remove ihe test 
 tube, mix about 7 cc. of a solution of pyrogallol 
 with about the same volume of sodium hydroxide. 
 At once pour the liquid into the funnel. By open- 
 ing the clamp let the liquid down till it just 
 reaches the lower end of the tube. Attach the test 
 tube by inserting the stopper firmly, and opening 
 the clamp move it till it closes out of the way upon 
 the rubber and glass connecting tube. Move the 
 test tube to a horizontal position to expose the air 
 to more of the liquid surface. If the surface of 
 the liquid in the funnel threatens to lower into 
 the neck, add a little water. When you are sure 
 the liquid has ceased to enter the test tube, invert 
 the test tube and bring it up till the surfaces of the 
 liquid in test tube and funnel are at the same 
 Fig. 11. height, as shown by the dotted lines. Estimate, or 
 
 better, measure the lengths of the columns of the liquid in the test 
 tube and the gas which remains, Which represents oxygen and which 
 nitrogen? Find their ratio by volume in air. 
 
 15. Preparation of jVitrogen from Chemicals: Set up the appar- 
 atus as shown in fig. 12. The nitrogen is obtained by heating a solu- 
 tion of ammonium nitrite but since this substance is difficult to make 
 and to preserve, the same results may be obtained by heating a solu- 
 tion of ammonium chloride, NH 4 C1, and sodium nitrite, NaNO 2 . These 
 give in the solution the same groups, NH 4 and NO 2 , which form am- 
 monium nitrite, NH 4 NO 2 . 
 
 Put in the flask 10 grams of ammonium chloride, 10 of sodium ni- 
 trite and lOOc.c. of water. Heat and after the air has been expelled col- 
 lect the nitrogen in jars. 
 
 Light a bit of candle, place it upon the deflagration spoon and 
 lower it into a jar of the nitrogen. Lower burning phosphorus into 
 another jar. 
 
 Note the curious fact in this experiment that the oxygen of am- 
 monium nitrite oxidizes the hydrogen of the same compound; that is, 
 one part cf it acts as a reducing agent and the other as an oxidizing 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Fig. 12. 
 
 agent. Also, hereafter notice that whenever oxidation takes place witti 
 respect to one substance, reduction takes place at the same time with 
 regard to some other substance. Find illustrations of this under prep- 
 aration of chlorine, 24. 
 
 46. An analogous method may be used for the preparation of N by 
 heating an intimate mixture of ammonium chloride, 2 grams, and po- 
 tassium dichromate, 5 grams, instead of ammonium dichromate. The 
 method of heating and collecting is quite the same as that for oxygen 
 in 6. 
 
 4". Preparation of Ammonia, Place about a half gram of ammon- 
 ium chloride, ammonium sulfate, ammonium nitrate in three test 
 tubes. Add to each about 2c.c. of sodium hydroxide solution and 
 warm. Note odor in each case, and hold in each tube a strip of moist 
 turmeric paper. Repeat using about the same amounts of the am- 
 monium compounds, but mix each with about its own weight of slaked 
 lime. Formulate a general method of preparing ammonia. 
 
 Refer to a text-book for an account of calcium cyanamide, its 
 manufacture, hydrolysis, use, importance. In a test tube place about 
 a gram of the substance, barely moisten with water. Suspend in the 
 tube a strip of moist turmeric paper, corking the tube and thus hold- 
 ing the strip in place. Observe evidence of ammonia after an hour. 
 
 Set up the apparatus as in fig. 13 and provide the dry fountain 
 bottle as in 27. The collecting jars or bottles must be dry. Weigh and 
 
28 
 
 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 mix in mortar 10 grams each of 
 ammonium chloride and slaked 
 lime. Place on paper and slide 
 into iron tube. Protect the stop- 
 per of tube with wet cotton. 
 Heat the tube and collect two 
 bottles or jars of gas by the 
 downward displacement of air. 
 When full moist turmeric paper 
 held at the mouth of a jar will 
 be instantly turned brown. 
 While it is bottom upward 
 clamp cover on jar or cover 
 bottle with glass plate ani 
 place on desk bottom upward. 
 Fill the fountain bottle and pro- 
 duce a fountain as with HC1 in 
 Fig. 13. 27. Turn the delivery tube down 
 
 and place it in a test tube containing lOc.c. pure water so that the end 
 of the delivery tube shall reach just under the water. Observe ab- 
 ( sorption of the ammonia and lower the delivery tube in the water as 
 necessary to absorb all the gas. Heat as long as ammonia seems to 
 come off, remove delivery tube from test tube then remove the lamp. 
 Observe odor of the solution of ammonia in the test tube. There 
 is evidently gaseous ammonia above the liquid. There is also am- 
 monia in solution, some ammonium hydroxide and some ionized am- 
 monium hydroxide. 
 
 48. Action of Acids and Ammonia; Neutralization: In a jar pour 
 a few c.c of con. HC1, wet the sides of the jar and pour out excess oi 7 
 liquid. Place over this the jar of ammonia gas, bottom upward, remove 
 cover, placing jars mouth to mouth. Mix by reversing the pair, placing 
 the jar with HC1 above. The fumes and the solid on walls of jars con- 
 sist of ammonium chloride, NH 4 C1, made by direct union of what? Try 
 a burning splinter in the remaining bottle of ammonia gas. 
 
 In this paragraph and in all other such cases do not dip test papers 
 into liquids, but take out a drop of the liquid with a stirring rod and 
 touch it to the paper. In a porcelain dish place lOc.c. of ammonium 
 hydroxide and neutralize with dilute nitric acid in just the same way 
 as sodium hydroxide was neutralized with HC1 in 28. Evaporate about 
 two-thirds of the liquid over a flame and complete the evaporation on 
 a water bath. Why on a water bath? Press dry the ammonium nitrate 
 between folds of filter paper. 
 
 Small quantities of ammonia may be tested for as directed in 
 Qualitative Analysis, Group V. Very small quantities are tested for 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 29 
 
 with Nessler solution, which is a strongly alkaline solution of mercury 
 and potassium iodides. If at hand and to show the sensitiveness of the 
 test, fill a clean test tube nearly full of distilled water, and in another 
 tube of distilled water dissolve a granule of ammonium chloride as 
 large as a pin head. To each tube add about 2c.c. of Nessler solution. 
 In very dilute solutions the amounts of ammonia are in proportion to 
 the depths of color, and the test is much used to determine very small 
 amounts of ammonia in water. 
 
 49. Preparation of Nitrogen Monoxide (Nitrous Oxide): Set up 
 the apparatus as in fig. 12, but omit the thistle tube. Heat in the dry 
 flask 15 grams of ammonium nitrate, and regulate the heat so as to 
 control the flow of gas, which should be collected in three jars or bot- 
 tles. When through collecting remove the delivery tube from the 
 water, then remove the burner. 
 
 In one jar thrust a glowing splinter. In another lower burning 
 phosforus. Burn a piece of picture cord in the same way as in oxygen. 
 Write the equation and compare' it with the self-oxidation of ammonium 
 nitrite to form nitrogen. 
 
 50. Preparation of Xitrie Acid: Arrange the retort and receiving 
 flask as in the preparation of bromine, fig. 9, but retort and flask should 
 preferably be dry. Put into the retort as there described 20 grams of 
 sodium or potassium nitrate. Add through the tubulus of the retort 
 by means of a funnel 25c.c. of con. sulfuric acid. At once wash the 
 funnel and measuring cylinder. Let stand till acid and solid are in 
 contact throughout and then heat with a small flame, preferably of the 
 crown top. There is no need of cooling the receiving flask if the heat 
 applied is properly regulated. Turn down the flame if fumes escape 
 in considerable amount from the flask. Stop heating when the acid 
 comes over slowly and the liquid in the retort is clear and seems to 
 be viscid. What is the substance left in the retort? To remove it when 
 cool fill the retort nearly full of water and heat gently, best on a water 
 bath. Pour off solution, add fresh water and heat again. Repeat till 
 all is dissolved. If the cake of solid comes loose do not shake it about, 
 since the thin retort will be broken. 
 
 51. Oxidizing Action of Mtric Acid: Nitric acid, after free oxy- 
 gen, is the most important of all oxidizing agents and its oxidizing ac- 
 tion should be well understood. 
 
 When the very concentrated acid is heated a part of it breaks down 
 thus: (1) 2HNO 3 =H 2 O+2NO 2 +O. Whenever the more dilute acid is in 
 contact with something easily oxidized the reaction is likely to be: 
 (2) 2HNO 3 =H 2 O+2NO+3O. To illustrate, (a) heat a few drops of your 
 acid and note the red-brown gas given off. (b) Place a loose plug of 
 woolen yarn in the mouth of a test tube which contains about 2c.c. of 
 your acid, and boil the acid till the vapors ,set the wool on fire, (c) 
 
30 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Heat a little sawdust in a dish till it begins to char and pour upon it 
 about Ic.c. of your acid, (d) Place about 3 grams of sugar in a flask 
 and 25 c.c. of con. nitric acid from the shelf. Heat till copious fumes 
 of N0 2 are formed. The sugar is mainly oxidized to oxalic acid. The 
 next experiment contains a good illustration of oxidation of metals 
 by' HNO :i . 
 
 In the most common case of oxidation with nitric acid the acid 
 itself is reduced to water and nitric oxide, NO. Oxidation and reduc- 
 tion, it will be remembered, go on together usually. In some cases the 
 reduction does not go so far. An instance is the reduction of a nitrate 
 to nitrite by heating with lead (see 54). In some cases the reduction 
 proceeds to the formation of hydroxyl-amine, HONH 2 , but more often 
 to ammonia. To illustrate this treat a few bits of aluminum in a test 
 tube with about 3 c.c. of sodium hydroxide, and add 2 drops of con. 
 nitric acid. Warm the tube and when the action becomes rapid note 
 odor of ammonia. Incline the tube and hdld in it moistened turmeric 
 paper without touching the glass. Here the hydrogen from the sodium 
 hydroxide and Al reduces the nitric acid to ammonia. Arrange in a 
 horizontal line in note book the successive reduction products of ni- 
 tric acid. 
 
 Aqua regia is a mixture of 1 part of con. nitric with 3 parts of con. 
 hydrochloric acid. Make about 1 cc. of the mixture, heat it and notice 
 chlorine given off. It will dissolve many substances that are attacked 
 by neither of the acids alone. 
 
 52. Action of Nitric Acid on Metals, Nitric Oxide: See in a text 
 book the electro-motive series, and the action of nitric acid on metals. 
 
 The very electro-positive metals will give some hydrogen with di- 
 lute nitric acid. Set up the apparatus to prepare H, but use a test 
 tube instead of the flask. Place in tube magnesium turnings. Dilute 
 con. HN0 3 by mixing 5 c.c. with 30 c.c. of water. Fill two test tubes 
 with water and have them ready to collect the gas. Put in the dilute 
 acid, about 5 c.c. at a time and after the air has been expelled collect 
 two test tubes of the gas. Test the gas for H. Note red gas in test 
 tube when air enters, using the second tube. Save the liquid in the 
 tube in which H was evolved. Repeat the experiment, using granulated 
 zinc instead of magnesium. Can you detect any H? Note red furnes of 
 nitrogen dioxide, N02, always formed when nitric oxide, NO, comes in 
 contact with air or oxygen. When the action is over pour off a little 
 liquid from the Zn and test it and the solution saved from the mag- 
 nesium for ammonia: To the liquid add NaOH in excess, hold a strip 
 of moist turmeric paper in the tube not letting it touch the glass and 
 warm gently. Also note odor of ammonia. 
 
 The metals lower than zinc in the electro-motive series give no H 
 with nitric acid, and those below hydrogen do not give H with any 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 31 
 
 acid. Try the action of con, nitric acid diluted with an equal volume 
 of water, on a little tin, copper, powdered antimony. Oxides are ob- 
 tained with the first and third metals. Would HC1 act on the copper 
 and antimony? Could HNO 3 and metals be used to prepare pure H? 
 
 53. Nitric Oxide: Use apparatus as in fig. 3. Put in flask 25 
 grams granulated copper, about 10 cc. water, see that thistle tube 
 reaches into the water, then add con. nitric acid as necessary to secure 
 a moderate flow of gas. Reject the first half jar full of gas. Fill three 
 jars, leaving one on the shelf of the trough. Pass some of the gas 
 through a solution of ferrous sulfate in a tube. The dark color is due 
 to FeSOiNO. What does the blue solution in the flask contain? Pour 
 it into a vessel provided for the purpose. Wash the ^remaining copper 
 and put into the vessel provided for it. 
 
 Prepare oxygen as in 6 or 9 and pass a little into jar on shelf. 
 The red gas is nitrogen dioxide, NO 2 . Let it become absorbed then pass 
 in more oxygen. Thus continue till the water ceases to rise in the 
 jar. Write the volume equation for the union of NO and 2 . What is 
 formed when NO 2 dissolves in oold water; in hot water? 
 
 In one jar of NO burn red phosphorus. Try a candle in the other. 
 
 Express the action of nitric acid on copper in two equations, the 
 first of which shows the decomposition of the acid in oxidation as in 
 equation (2), 51, forming 3CuO; the second the dissolving of this oxide 
 in the acid. 
 
 54. Reduction of a Nitrate to Nitrite, Nitrous Acid: (a) In an 
 iron crucible heat and stir with an iron rod, spike nail, 5 grams of 
 potassium nitrate and 20 grams granulated lead till the mass glows. 
 When cold add water and boil for some time. Filter off the liquid into 
 a tube. Add an excess of dil. H 2 SO 4 . Test the gas with moist starch 
 iodide paper. The sulfuric acid sets free nitrous acid which breaks 
 up into water, NO and NO-; equations. Note that NO+NO 2 =N 2 O 3 , the 
 anhydride of nitrous acid. It exists only as a liquid. 
 
 (b) Read (a) carefully. In a flask with thistle tube and right-an- 
 gled delivery tube place 10 -grams sodium nitrite, add enough water to 
 cover end of thistle tube, place delivery tube in 5 cc. NaOH in a wide 
 test tube. Add about 25 c.c. dil. sulfuric acid to the flask. Nitrous acid is 
 set free, but breaks up as in (a) . The gases are absorbed by the NaOH 
 making NaNO 2 . Add dilute acid to this solution and test the gas as 
 in (a). 
 
 55, Nitrogen Dioxide, Nitrogen Tetroxide: The lead nitrate used in 
 this experiment should be prepared in quantity before the laboratory 
 period by crushing to moderate fineness in a mortar and heating for 
 three hours in an oven at about 125. 
 
 Heat 10 grams of the dried lead nitrate in iron tube shown in fig. 
 13, turning the delivery tube down, and placing it in a small dry flask. 
 
32 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Heat until the flask is filled with the red-brown gas, NO 2 mixed with 
 N 2 Oi. Stopper the flask and heat it high over a flame, but do not heat 
 so much as to burst the flask. A temperature of 142 is sufficient to 
 change all the gas to NO 2 . Note depth of color. Let the flask cool a 
 few moments in air, then under running water, and observe the depth 
 of color again, when 80% of the gas will consist of colorless N 2 0u 
 Write the equation for the production of the gas and the reversible 
 reaction of the change from one to the other constituent of the mixture. 
 
 Heating lead nitrate again pass the gas into about 3 cc. of sodium 
 hydroxide for a few moments, then add dilute sulfuric acid to the 
 solution. What evidence have you that a nitrite was formed by the 
 action of the gas on sodium hydroxide? What else was formed? See 
 text book for the action of this gas with cold water and with hot water. 
 
 56. Tests for Nitrogen and the Nitrate Radical: (a) Tests for ni- 
 trogen in organic matter: Heat in a test tube with soda-lime or a mix- 
 ture of lime and powdered sodium hydroxide, bits of woolen material, 
 such as yarn, or bits of leather, dry albumin or corn meal, and hold in 
 the mouth of the tube wet turmeric paper. Is ammonia given off? 
 
 Secure an imperfect test tube and heat in it any of the organic 
 substances mentioned, with a bit of sodium not larger than half a pea. 
 (Do not touch sodium with the wet hands). Heat strongly till the 
 reaction is complete, and standing well back place the hot end of tube 
 in a little water in a beaker. The end cracks off. Boil the water and 
 filter. To a portion of the filtrate add a little solid ferrous sulfate, 
 then a few drops of ferric chloride, and boil. Now make acid wiih 
 dilute HC1, when a deep blue color will result. To'the rest of the solu- 
 tion add a crystal of potassium nitro-ferri-cyanide, which will give a 
 red color, showing sulfur in the organic matter. 
 
 (b) Tests for the nitrate radical, or nitric acid: To a few cc. of a 
 solution of ferrous sulfate add a very little of any nitrate, and when 
 the substance is dissolved and the solution is cool incline the tube and 
 pour in steadily con. H 2 SO 4 , so that it will run down and collect at the 
 bottom of the tube. Where it and the solution meet will form a dark 
 ring of FeSO. (NO). 
 
 The following test should be used only for excessively small 
 amounts of nitrate: Upon a bulk of any nitrate not larger that a pin 
 head and in a dish, drop about ten drops of phenol-disulphonic acid. 
 Warm the dish gently, best on a water bath for a few moments. Add 
 about 25 cc. of water and make alkaline with ammonia. A yellow 
 solution shows nitrate radical is present. 
 
 ELECTRO-CHEMISTRYTHEORY OF IONIZATION. 
 
 57. Electrical Terms and Fundamental Laws: Using good text 
 books on chemistry and on physics study, till clear to you, the meaning 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 33 
 
 of these terms: anode, cathode, Volt, Ampere, Ohm, Coulomb, Watt, 
 electro-chemical equivalent, electro-motive series of the elements. 
 
 What is Ohm's Law? On a 110-Volt circuit find the current 
 strength or amperage transmitted by a 25-Watt lamp. Find the resis- 
 ance of this lamp in Ohms. 
 
 What is Faraday's Law? What are the electro-chemical equiva- 
 lents of H, O, Cl, SO 4 , NO 3 , Cu, Ag, Sn". Sn""? Give relations of 
 electro-chemical equivalents, atomic weights (or in the case of radicals 
 the sum of the atomic weights), and valences. 
 
 >*. Electrical Coiuluctiyity : The current used is the direct light- 
 ing circuit, 110 volts, cut down by a 25-Watt lamp. This would give 
 what amperage through a wire without resistance connecting th,e two 
 binding posts? Calculate and read the ammeter. Do the results agree? 
 Determine the comparative conductivity of the following 5th normal 
 solutions, already made by the instructor: hydrochloric acid, sulfuric 
 acid, acetic acid, sodium acetate, sodium hydroxide and ammonium hy- 
 dioxicle. Since they are all N/5 they have the same number of chemical 
 equivalents per unit volume. According to whose law might they be 
 expected to conduct the same? 
 
 To determine the conductivity secure one of the conductivity tubes 
 
 as shown in fig. 14, wash, rinse 
 with distilled water and drain it. 
 In each case be careful thus to 
 clean the tube. Paste a strip 
 of gummed paper on it and 
 in each case fill to this 
 mark. Connect the tube thus 
 filled with the binding posts, 
 note the intensity of the glow of 
 the lamp and read the ammeter. 
 Thus determine the conductivity 
 of all. Why should they conduct 
 so differently? Knowing the re- 
 sistance of the lamp and the cur- 
 rent it alone transmits, calculate 
 from one of your readings of the 
 ammeter the resistance of the 
 Fig. 14. solution? 
 
 Fill the conductivity tube with a neutral solution of sodium or pot- 
 assium sulfate, add a few drops of litmus solution and mix well, and 
 subject the solution to the action of the current. Which electrode 
 gives off oxygen and shows the formation of acid around it; which 
 shows hydrogen evolved and alkali formed around it? Try a solution 
 of common salt without litmus solution. Prove with starch-iodide pa- 
 
34 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 per that Cl is given off at the anode. Test the solution at the cathode 
 with turmeric paper. By what secondary reaction is alkali formed in 
 the last two instances? What is the action of sodium on water? (13) 
 
 In electrolysis many of the heavy metals instead of acting upon 
 water at the cathode are there deposited. Subject a solution of cop- 
 per sulfate to electrolysis for several minutes. What is deposited on 
 the cathode? t What gas is given off at the anode? Copper sulfate 
 itself reacts slightly acid, but if the solution at the anode be tested it 
 will be found much more acid. How do hydrogen and the metals ac- 
 cumulate at the cathode and the acid radicals, with which they were 
 associated, at the anode? 
 
 59. Chemical Facts Best Explained on the Theory of lonization: 
 The Mature of Acids: Test with blue litmus paper dilute HC1, dilute 
 HNO 3 , dilute H 2 S0 4 , acetic acid and any other acids that may be avail- 
 able. Why should substances of such different composition all turn 
 the paper red? Why should they all taste sour? Try the action of 
 each on bits of zinc. Why should all give hydrogen? Why should 
 they all neutralize bases giving salts and water? Why should they all 
 give hydrogen at the cathode when electrolyzed? 
 
 It would seem that H is the one component of all acids, and that 
 to it their acid properties are due. In electrolysis H accumulates at 
 the cathode and the rest of the acid molecule at the anode. It is not 
 far to the thought that the H is, in the solution of an acid, compara- 
 tively free from the rest of the molecule. Since unlike electric 
 charges attract each other and like charges repel, the inference is that 
 the H atom is charged positively thus, H + , and the remainder of the 
 molecule, for example, Cl is charged negatively, thus, Cl". In this con- 
 dition the H and Cl atoms are called ions. Ions are formed when acids, 
 bases and salts are dissolved in water and without regard to any in- 
 fluence of the electric current. 
 
 60. The Character of Bases: Test with turmeric paper solutions 
 of sodium hydroxide, potassium hydroxide, ammonium hydroxide, cal- 
 cium hydroxide, barium hydroxide. Again, why should such different 
 substances all turn the paper brown? They also all neutralize acids in 
 the same way forming salts and water. Their common constituent is 
 hydroxyl OH. Since all bases have OH it is inferred that the alka- 
 linity is due to this group. It seems to be easily separable from the 
 metal of such group as NH 4 . In electrolysis it travels to the anode 
 where it breaks up into water and oxygen while the metal goes to the 
 cathode. For reasons given under acids, it is believed that a base in 
 solution is more or less ionized. That is, sodium hydroxide consists 
 largely of the ions Na + and OH". 
 
 61. Ready Cleavage or lonization of Salts: In as many test tubes 
 place a few drops of solutions of the following: NaCl, KC1, NH 4 C1, CaCla 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 35 
 
 BaCl 2 and other chlorides that may be at hand. Add to each a few 
 drops of silver nitrate. Why should all these different chlorides gives 
 the same precipitate of silver chloride? Silver sulfate or silver ace- 
 tate might have been used instead of silver nitrate and precisely the 
 same precipitate of silver chloride would have resulted. In the solu- 
 tions of chlorides it is evident that chlorine is very slightly held by the 
 metals if at all, and the same is true of the silver in the solutions of the 
 silver salts. All these salts are electrolytes, the metals going to the ca- 
 thode and the non-metals to the anode. The inference is that they are 
 largely ionized in solution. Thus, sodium chloride is largely Na + and 
 d~, and silver nitrate is Ag + and NO 3 ~. 
 
 Any element as Cl is not always an ion in solution. To solutions of 
 pure potassium chlorate, chloral and chloracetic acid add a little silver 
 nitrate solution. No silver chloride is obtained, since KC1O 3 ionizes into 
 K + and C10 3 ", chloracetic acid into H + and CH 2 C1CO 2 ~, while chloral 
 hydrate gives no ions. 
 
 62. Degree of lonization: Refer to your experiment on conduc- 
 tivity (58). Did all the acids conduct equally well and were they 
 equally ionized? Compare the conductivity and ionization of sodium 
 hydroxide and ammonium hydroxide ; of acetic acid and sodyAm acetate. 
 
 The ionization of the same substance may be increased by diluting 
 and decreased by concentrating its solution. 
 
 (a) Compare the colors of 2-normal solutions of copper sulfate, 
 nitrate and chloride. The color of each is supposed to be due to the 
 color of the undisscciated salt and to the copper ion. Now dilute a 
 small portion of each solution with 10 times its volume of water. Why 
 are the solutions now more nearly the same color? To a portion of 
 each solution add an excess of ammonia. The same color is due to the 
 same ion, Cu(NH 3 ) 4 ++ , 
 
 (b) The concentration of a solution may be in effect increased 
 and the dissociation of the solute decreased by adding a substance hav- 
 ing an ion in common with the solute. Solid copper bromide is black, 
 its concentrated solution is brown due to CuBr 2 , while its dilute solu- 
 tion as that of every other cupric salt shows blue due to the Cu ion. In 
 a dish dilute about 2 c.c. of the brown solution till it becomes blue. To 
 one half of the blue solution add a few drops of con. HBr, which has 
 the common ion Br, till brown. To the other half add solid copper 
 chloride, having the common ion Cu, till brown. Explain how the ad- 
 dition of the common ion effects these changes. 
 
 (c) Dilute a little acetic acid from shelf bottle with 20 times its 
 volume of water. Add a few drops of methyl orange and divide into 
 two portions. To one add one half its volume of con. sodium acetate 
 and compare its color with that of the other half. What common ion 
 was added and what was the effect on the acidity of the acid? 
 
36 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 (d) To a saturated solution of salt add concentrated HC1 till a 
 large precipitate of salt is obtained, and account for its formation. 
 
 63. Acidity Due to Hydrolysis: Many substances containing no 
 hydrogen give an acid solution. They are mostly salts of metals which 
 form weak bases and radicals of stronger acids. With blue litmus pa- 
 per test solutions of salts of copper, iron, aluminium. In these cases 
 the hydrolysis goes only a little way, for example thus: 
 
 FeCl 3 +H 2 0=(reversibly)FeCl,OH+HCl. 
 
 Add a little antimony chloride and bismuth chloride to five times 
 their volumes of water. Here the reactions go far to the right precipi- 
 tating SbOCl and BiOCl. Now reverse the reactions by adding con. 
 HC1. See the hydrolysis of the halides of phosphorous, 34 and 79. 
 
 64. Alkalinity Due to Hydrolysis: Some salts of strongly basic 
 metals and weak acids give alkaline solutions due to hydrolysis. Test 
 with turmeric paper or a few drops of phenol-phthalein solutions of bor- 
 ax, sodium carbonate, Na 2 CO3, and sodium phosphate, HNa 2 POi. In the 
 case of sodium carbonate, Na 2 CO 3 +H 2 O=HNaC0 3 4-NaOH, and of 
 course the last is highly ionized, and the alkalinity is due to the ion 
 OH~. For olher examples see 77 and 95. 
 
 65. The Electro-Motive Series of Elements: In the following any 
 metal which replaces another from solution is said to have a greater 
 solution tension or to be more electro-positive. In a little zinc sulfare 
 solution place bits of magnesium turnings and let stand. When the 
 action has nearly ceased note gray deposit of zinc. Place zinc in a so- 
 lution of cadmium chloride and later note cadmium deposited on the 
 remaining zinc. Try cadmium or zinc in solutions of copper sulfate 
 and lead acetate and note copper and lead deposited. Try iron in a 
 copper solution. Place a strip of copper in a solution of mercuric 
 chloride and after a few minutes remove, rub and note mercury coat- 
 ing. Put a copper wire in a silver solution and note deposit of silver, 
 Try silver in a solution of gold chloride and note deposit of gold on 
 the silver. All these cases are practically alike. One metal goes into 
 solution as ions and the other is as it were forced out as neutral metal; 
 for example, Cu ++ +SO4- + Zn=(reversibly)Zn ++ +SO 4 "+Cu. Write a 
 similar ionic equation for the preparation of H with zinc and sulfuric 
 acid. Has the SO 4 much to do with either case? 
 
 Arrange the above metals in the order of their capacities to replace 
 other metals, or their ionizing tendency. Compare the result with the 
 arrangement of the Electro-Motive Series in a text-book. Wh'at two 
 metals above if used for the plates in an electric battery cell would 
 give the greatest electromotive force? 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 37 
 
 SULFUK. 
 
 <;#. Solubility and Crystallization: Do not heat carbon disulfide 
 or have a flame near it, since it is very volatile and easily inflam- 
 mable. 
 
 Place in a dry test tube about 2 grams of powdered sulfur, add 
 not more than one-fourth test tube full of carbon disulfide, shake and 
 filter through a dry filter into a crystallizing dish or small beaker. 
 Let two or three drops fall into a clean watch glass. After the liquid^ 
 has evaporated at room temperature, examine the crystals in each ves- 
 sel, and those in the watch glass with a microscope. The crystals be- 
 long to the rhombic variety of sulfur. 
 
 <J7. Effects of Heat, Allotropic Forms: Since the test tube can- 
 not be cleaned fill an imperfect one three-fourths full of lumps of 
 sulfur, slowly and evenly heat till it is all changed to a yellow, mobile 
 liquid. Now increase the temperature and note that it becomes darker. 
 Find a temperature at which the tube may be inverted for a moment 
 without the sulfur's running out Increase the temperature till it 
 shows signs of boiling and becomes quite fluid again. Pour in a steady 
 small stream into a beaker or other 'vessel full of water. Examine 
 this "plastic sulfur", place aside and examine at the end of the lab- 
 oratory period and at the next period. It gradually changes to mono- 
 clinic sulfur, and after a long time to rhombic, the permanent form. 
 Consult text-book on the forms of sulfur, and their properties. 
 
 <js. Hydrogen Sulfide: Fit a small bottle or flask, the former pre- 
 ferred, with a delivery tube consisting of two right angled pieces joined 
 with rubber tubing, the outer piece long enough to reach to the bot- 
 tom of the collecting vessel or test tube when resting, on the desk. 
 Place in the generating vessel 20 grams of ferrous sulfide in small 
 bits, sliding in the larger pieces to avoid breakage if a flask is used, 
 add about 40 c.c. of water and drop directly upon it a little at a time 
 con. sulfuric till a suitable flow of gas is secured. If the flow of gas 
 becomes too slow at any time do not add more acid, but pour off the 
 liquid which is nearly exhausted of acid but contains much iron salt, 
 and add water and acid as in the beginning. 
 
 Collect a bottle of the gas by displacement of air, in the same 
 way as chlorine, and burn the gas in the bottle. What is the deposit 
 on the sides of the bottle? Collect gas in a test-tube one- fourth full 
 of water, cover mouth of tube with the thumb and shake. Is the gas 
 soluble in water? Turn upward the outer section of the delivery tube 
 and light the gas at the orifice. Hold cold porcelain #s a dish or cru- 
 cible lid in the flame at about its middle. What is the yellow deposit? 
 Insert the flame into a bottle with wide mouth as far as you can with- 
 
38 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 out extinguishing the flame. Why does sulfur deposit on the dish and 
 bottle instead of burning? 
 
 69. Precipitation with Hydrogen Sulfide: Hydrogen sulfide is 
 much used in separating the metals into groups in analytical chemis- 
 try. In test tubes in the test tube holder place about 5 c.c. of solu- 
 tions of salts of the following metals, using chlorides, nitrates, ace- 
 tates as may be convenient: silver, lead, mercury, bismuth, copper, 
 cadmium, antimony, arsenic, tin; also, zinc, cobalt, nickel, manganese, 
 iron; also, barium, calcium, strontium, sodium, add a drop or two of 
 dilute HC1 or HNOs to each. Pass the gas flowing with moderate ra- 
 pidity into each of these solutions for about a minute, washing off the 
 delivery tube by dipping into water before inserting it into a new so- 
 lution. Note carefully what metals can be precipitated as sulfides 
 with H 2 S from slightly acid solutions remembering the acid added and 
 the fact that when precipitation occurs acid is necessarily set free. 
 The precipitation of zinc is not complete. Pass an excess of gas into 
 the tube containing zinc, filter off the precipitate and add to the fil- 
 trate and to each succeeding solution, beginning with cobalt, ammon- 
 ium hydroxide, which should give precipitates of sulfides in all until 
 barium is reached. That is, several metals not precipitated as sul- 
 fides in acid solution are precipitated with hydrogen sulfide and an 
 alkali and they form another group. Name the metals not precipi- 
 tated as sulfides. 
 
 Return to the sulfides precipitated from acid solution, pour off the 
 liquid from the sulfides of Hg, Cu,- Bi and try to dissolve each by boil- 
 ing 'with dilute nitric acid. How may Hg be separated from the other 
 two? Pour off the liquid from the sulfides of lead, antimony and ar- 
 senic and heat each with about 5 c.c. of ammonium sulfide, but do not 
 boil. How may lead sulfide be separated from arsenic and antimony 
 sulfides? 
 
 These are merely illustrations of the use of hydrogen sulfide and 
 many others will be met. Instead of using hydrogen sulfide and am- 
 monium hydroxide separately to precipitate certain sulfides, it is com- 
 mon practice to use ammonium sulfide which may be easily made by 
 saturating diluted ammonium hydroxide solution with hydrogen sul- 
 fide, and then adding an equal volume of the same ammonia solution 
 which gives essentially (NH 4 ) 2 S. 
 
 70. Reducing Action of Hydrogen Snlfide. Hydrogen sulfide is 
 not a very stable compound and it acts as a reducing agent in much 
 the same way as hydriodic acid. Pass hydrogen sulfide into a dilute 
 acidified solution of potassium dichromate, till green ; into a dilute acid- 
 ified solution of potassium permanganate till colorless; into bromine 
 water and a dilute solution of iodine till colorless. In the last two 
 cases filter off the sulfur and test liquids with litmus paper. What 
 
LABORATORY MANUAL, OF GENERAL CHEMISTRY 
 
 39 
 
 acids were formed? Compare the reducing action of hydrogen sulfide 
 with that of sulfur dioxide in experiment 71. 
 
 71. Sulfur Dioxide: (Read reference book on sulfur dioxide, sul- 
 fur trioxide and sulfuric acid.) In a test tube treat a little sodium 
 sulfite with dilute acid. Note odor and test with a strip of paper mois- 
 tened with mercurous nitrate. It should turn dark. 
 
 Determine whether bits of charcoal and sulfur will give SO 2 when 
 heated with con. sulfuric acid. Explain their action on the acid. Ar- 
 range apparatus as in the preparation of hydrochloric acid. Place in 
 the flask 20 grams of granulated copper or bits of sheet copper or 
 turnings and add 25 c.c. con. sulfuric acid. Before heating read 
 through the experiment and have everything needed at hand so as to 
 continue the flow of gas the minimum of time. 
 
 Heat till gas comes off freely and control the rate by turning 
 down the flame. After acidifying them with acetic acid pass gas into 
 a solution of potassium dichromate till green and into a solution of 
 potassium permanganate till colorless; also, into bromine water till 
 colorless, into dilute nitric acid, and into distilled water. To each so- 
 lution add barium nitrate and then dilute HC1. The formation of a 
 white precipitate by barium nitrate, insoluble in HC1 shows the pres- 
 ence of the sulfate radical, SO 4 . 
 
 The first action of SO 2 on passing it into the above solutions was 
 probably the formation of the unstable sulfurous acid, H 2 SO 3 . All the 
 substances into which the gas was passed were oxidizing agents except 
 of course the distilled water. Show their action on the sulfurous acid. 
 Test for the sulfate radical in the dilute sulfuric acid on the shelf, in 
 solutions of magnesium sulfate, copper sulfate and other sulfates that 
 may be at hand. 
 
 72. Preparation of Sulfuric Acid: Set up the apparatus as shown 
 in fig. 15. The ignition tube may be of hard glass or of half inch gas 
 pipe. If the latter is used it should be 12 to 14 inches long to prevent 
 the overheating of the stoppers. Even then it is well to spray the ends 
 with the wash bottle catching the water as it runs off, or to wrap the 
 ends with cotton or filter paper and keep it saturated with water. 
 
 Fig. 15. 
 
40 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Using loose plugs of asbestos fiber or steel wool ^confine near the 
 middle of the tube about 15 grams of granular iron pyrite. Connect 
 the outer end of the tube with the compressed air system. If this is 
 not available a bulb pump, foot-bellows or even a bicycle pump may be 
 used to force through the air. Instead of forcing through the air the 
 other end of the apparatus may be connected with the exhaust 
 system or a filtering pump. 
 
 Place in the test tube 3 c.c. of con. nitric acid, preferably fuming 
 acid, and to the larger bottle, connected with it add about 5 c.c. of 
 water. The smaller bottle contains a solution of sodium hydroxide to 
 absorb the excess of noxious gases, 
 
 Heat the tube containing the pyrite, preferably with a wing burner 
 and force or draw through a moderate stream of air. When the bot- 
 tles are filled with white fumes it may be known that the pyrite is oxi- 
 dizing. Moderate the heat to avoid distilling off unburned sulfur, and 
 regulate the air so that only a small amount of fumes escape from the 
 smaller bottle. Continue the operation 10 to 20 minutes. To avoid 
 sucking back remove the air connection and then the burner. 
 
 Transfer the liquids in the test tube and larger bottle to a 
 weighed porcelain dish, rinse with a little water and add this to the 
 dish. With a free flame evaporate till the steam and nitric acid va- 
 por give place to the dense gray fumes of sulfuric acid. Let the dish 
 cool and weigh it. Three to four grams of con. acid should be made in 
 a 10 minutes run. Show that the acid will carbonize a splinter dipped 
 into it. Dilute a few drops of it and test with barium chloride. Re- 
 move and examine the contents of the ignition tube. What results 
 from burning iron pyrite in air? 
 
 73. Sodium Sullite ami Thiosulfate: Dissolve 10 grams anhydrous 
 sodium carbonate in 50 c.c. of water with aid of heat if necessary, cool 
 and place just one half of the volume in a large test tube. Pass into 
 the solution in tube sulfur dioxide made by heating in a flask with 
 thistle tube 25 grams of copper turnings with 25 cc. of concentrated 
 sulfuric acid. The delivery tube should reach to the bottom of the 
 test tube. When the flow of gas begins the heat should be reduced so 
 as to avoid a too rapid stream. As the carbon dioxide is set free the 
 liquid shows a tendency to froth. The reaction is completed in about 
 15 minutes when small bubbles disappear, the sulfur dioxide appar- 
 ently is no longer absorbed, and the liquid smells strongly of the gas. 
 The result is acid sodium sulfite. Now add this solution to the re- 
 served half of the sodium carbonate solution which will give normal 
 sodium sulfite. Why? Place the combined solutions in a dish and boil 
 for two or three minutes. If evaporated nearly to dryness on a water 
 bath and allowed to cool crystals of sodium sulfite would be obtained. 
 The whole of the salt may, however, be converted into sodium thiosul- 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 41 
 
 fate. For .this purpose add to the solution three grams of powdered 
 sulfur and boil gently for twenty minutes replacing from time to time 
 the water evaporated. Filter off the sulfur and evaporate the filtrate 
 over a water bath till crystals begin to form and let the solution be- 
 come cold. 
 
 Pour off the solution from the crystals and use it to study the 
 properties of sodium thiosulfate. 
 
 To a small volume of a solution of iodine add the solution of thio- 
 sulfate till the iodine solution is decolorized. 
 
 Make a little silver chloride by adding a few drops of a solution 
 of some chloride to a few drops of silver nitrate. Now add the thio- 
 sulfate solution, a little at a time with shaking, till the silver chloride 
 dissolves. What is the use of sodium thiosulfate, "hypo," in photo- 
 graphy ? 
 
 To the remainder of the solution of thiosulfate add an excess of 
 dilute hydrochloric acid, warm and note the precipitation of sulfur 
 and the odor of sulfur dioxide. Account for the gas and the sulfur. 
 
 Here as elsewhere study the experiment with aid of a reference 
 book till all is clear and write the equations. 
 
 PHOSPHORUS. 
 
 70: ..White and Red Phosphorus, Allotropic Forms: (Caution: 
 Handle white phosphorus only with the forceps. Keep trace of all 
 pieces used and any remaining must not be left about the laboratory 
 desk or thrown in waste jars, but put into a vessel of water provided 
 for that purpose). 
 
 Place a piece of white phosphorus as large as a*grain of wheat in 
 water in a test tube and boil. Phosphorus vapor goes off with the 
 water vapor but burns Jo oxide on reaching the air. Try red phos- 
 phorus in the same way. Does it volatilize? 
 
 Heat a little red P in a dry narrow test tube and note sublimate 
 of white P on wall of tube. Change white P to red P by dropping a 
 small crystal of iodine upon white P in a dry test tube and warming. 
 
 Dissolve a little white P in about 3 c.c. of carbon disulfide in a dry 
 test tube, pour the liquid upon filter paper, place this upon the ring 
 of a retort stand and let dry. Try to dissolve red P in a little carbon 
 disulfide. Do not have a flame near carbon disulfide. What is meant 
 by "allotropic forms"? 
 
 "7. Phosphine: Drop a piece of fresh calcium phosphide into a 
 small bottle or beaker full of water. The gas set free is mainly phos- 
 phine, PH 3 , but it is mixed with another hydride, P 2 H 4 , which takes fire 
 spontaneously. 
 
 78. (Caution: The following experiment is a beautiful one, but 
 it is attended with some danger. It may be undertaken by a small 
 
42 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 class under constant supervision) . Set up apparatus as in the prepara- 
 tion of nitrogen, but place the end of delivery tube in a dish of water, 
 fig. 12. Place in the flask about 20 cc. of sodium hydroxide, made by 
 dissolving the solid in three times its weight of water. Add two or 
 three small bits of white P, then about 1 c.c. of ether having no flame 
 near, stopper the flask and heat. When the gas begins to ignite on 
 reaching the air, moderate the heat so as to maintain a rate of one 
 bubble every two or three seconds. Carefully insert the delivery tube 
 into a test tube of alcohol by which the P 2 H 4 is dissolved. Does the es- 
 caping gas now take fire? 
 
 When cold carefully remove the stopper, fill flask with water, wash 
 the remaining P and place in vessel of water provided. 
 
 79. Chlorides of Phosphorus and Hydrolysis. Refer to 34, 63. In 
 a dish place one drop of water and add one drop of phosphorus tri- 
 chloride, best from a dropping bottle. What acid gas is given off? 
 What acid of P remains? In another dish place about the same amount 
 of water and phosphorus pentachloride of about the bulk of a pea, 
 handling it with the spatula. What gas is here given off? Save the 
 liquid to test for phosphoric acid in the next experiment. 
 
 80. Tests for the Radical P0 4 and Phosphoric Acids: (a) To a 
 
 few c.c. of sodium phosphate add one-fourth its volume of ammonium 
 chloride, make alkaline with ammonia and then add magnesium sul- 
 fate. The slowly forming precipitate is ammonium magnesium phos- 
 phate, (NH 4 ) Mg P0 4 , and this is a good test for phosphoric acid pro- 
 vided arsenic acid is not present. Add water to the second dish in the 
 preceding experiment, boil, transfer to a test tube and test for PO 4 . 
 
 (b) To test for phosphate in a substance insoluble in water dis- 
 solve a small portion in dilute nitric acid, and to a few drops of the 
 solution add twice its volume of ammonium molybdate and warm. A 
 yellow precipitate shows phosphate, provided arsenic acid is known to 
 be absent. If practicable secure a piece of bone, place it in dilute ni- 
 tric acid, and near the end of the period test the solution for phos- 
 phate by this method. 
 
 (c) To solutions of sodium phosphate HNa 2 PO 4 and "microcosmic 
 salt, NH 4 HNaP0 4 , add silver nitrate. Note color of precipitate the 
 same in both cases and try its solubility by adding to one tube dilute 
 nitric acid and to the other an excess of ammonia. 
 
 81. Other Acids of Phosphorus: In a crucible heat a small 
 amount of ordinary solid sodium phosphate and after the water of 
 crystallization has been driven off continue with full flame for five 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 43 
 
 minutes. When cold dissolve some of the fused mass and test with 
 silver nitrate. The ordinary sodium phosphate, Na 2 HPO 4 has been 
 changed to sodium pyrophosphate, NaiPaOi. Repeat this experi- 
 ment using microcosmic salt. While the water is evaporating 
 test for ammonia with turmeric paper, and note odor. This salt is 
 changed by heating to sodium metaphosphate, NaPOs, at high tempera- 
 ture. Let the crucible cool, dissolve most of the metaphosphate and 
 test with silver nitrate. Compare the silver tests for phosphate, pyro- 
 phosphate and meta-phosphate. 
 
 ARSENIC. 
 
 S2. Acidify strongly with dilute HC1, dilute solutions of sodium 
 arsenite NasAsOj and arsenate NasAsO4 and without heating pass into 
 each H 2 S, for a few moments. Do you get a precipitate in the second 
 solution? Now heat each nearly to boiling and pass in the gas again 
 till the precipitation is complete. Filter on 3 the arsenic sulfide, As 2 S^, 
 and wash it on the filter. Try to dissolve a part of the substance in 
 con. HC1 in a dish. Now, boiling only gently add very small bits of 
 potassium chlorate from time to time till nearly all the 1 sulfide is dis- 
 solved. Add a little more con. HC1 to replace that boiled away, if ne- 
 cessary. Make alkaline with ammonia the solution in the dish and fil- 
 ter, and to the filtrate add a solution of magnesium sulfate or chlo- 
 ride. Compare this test for arsenate with that in the next section, and 
 with the test for phosphoric acid. (80a.) 
 
 Remove most of the remaining sulfide to a clean dish, add about o 
 c.c. of ammonium sulfide and heat but do not boil. This is a method of 
 separating arsenic sulfide from many other sulfides. To the solution 
 add an excess of dilute HC1 which will reprecipitate the arsenic sulfide 
 together with some sulfur. 
 
 83. Treat dilute solutions of sodium arsenite and arsenate with 
 silver nitrate and note carefully the colors of the precipitates. Try 
 the solubility of portions of each with dilute nitric acid and with am- 
 monia. Treat a solution of the sodium arsenate with ammonium chlo- 
 ride, make alkaline with ammonia and add a solution of 'a magnesium 
 salt. Compare the precipitate with that obtained in (82) in the same 
 way and also under phosphoric acid. How may an arsenate be dis- 
 tinguished from a phosphate? Given a mixture of arsenate and phos- 
 phate how could you remove the AsO 4 and test for the PO 4 ? 
 
 84. Tests for Small Amounts of Arsenic: One of the following 
 may be used: 
 
44 
 
 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 Fig. 16. 
 
 (a) The Modified Gutzeit Test: Good filter 
 paper should be wet with a solution of mercuric 
 chloride, allowed to dry, and cut into strips 
 narrow enough to slip into the horizontal tube, fig. 
 16. An arsenic solution containing 0.01 mg. per 
 c.c. and another containing 1 mg. per c.c. is made 
 up for class use. On no account should ordinary 
 arsenic solutions be used in this or the next, (b). 
 
 Place a strip of the pap*kr in the tube, and about 
 10 grams of coarsely granulated zinc in the bottle, 
 and add 15 c.c. of dilute HC1 with an equal volume 
 of water. After a little time look to see whether 
 the paper shows by brownish color that the mater- 
 ials used contain arsenic, and if not add 5 c.c. of 
 the arsenic solution of 0.01 mg. of arsenic. Let 
 stand 20 minutes and observe color of paper. 
 
 (b) Set up apparatus as in fig. 17, the outer section of delivery 
 tube being of hard glass. Place in flask 10 grams pure zinc and add 
 
 pure dilute sulfuric acid. 
 When air is expelled light 
 the gas with a test tube. 
 Now heat the tube and add 
 
 f~7 5 c.c. of the solution mark- 
 
 ^ ed .01 mg of arsenic to 1 
 
 c.c. After several minutes 
 examine arsenic mirror 
 beyond the heated por- 
 tion. Move flame toward 
 mirror. Is the latter easily 
 volatilized? Now remove 
 lamp and add to flask 5 
 c.c. of the solution marked 
 1 mg. of arsenic to 1 
 c.c. Hold in the flame a 
 
 porcelain evaporating dish 
 
 Fig. 17. or crucible lid. Examine 
 
 the spots of arsenic form- 
 ed, and try the action of a solution of sodium hypochlorite upon them. 
 Antimony also gives a similar mirror and spots, but it is not soluble in 
 the hypochlorite. 
 
 85. If desired to use either of the tests for arsenic for practical 
 purposes, prepare five standard papers or five standard tubes con- 
 taining .01, .02, .03, .04 .05, mg. of arsenic by using 1, 2, 3, 4 and 5 c.c. 
 of the solution containing .01 mg to 1 c.c. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 45 
 
 Now cut a piece of highly-colored wall-paper for example, one 
 decimeter square, tear into bits, place in a dish, heat with 5 c.c. con. 
 sulfuric acid, until the paper is completely charred. Let cool, treat 
 contents of dish with 25 c.c. of water, stir well and filter. Test the 
 solution for As as above, using only so much as necessary to obtain a 
 mirror or paper containing .01 to .05 mg. of arsenic. Compare with 
 standards, and calculate the amounts of As to one sq. meter of wall 
 paper. A similar proceeding will suffice for cotton fabric. 
 
 AtfTIMOXY. 
 
 86. Heat about 0.5 gram of finely powdered antimony with about 
 10 c.c. con. HC1. Does it dissolve? Now add to the test tube about 1 
 c.c. of con. HNO 3 , being careful that the liquid does not froth over on 
 the hand. When most of the metal is dissolved pour a little of the so- 
 lution into test tube of water, and another portion into a solution of 
 salt. Compare the precipitate of antimony oxy-chloride SbOCl in the 
 two tubes, allowing for salt which may precipitate in the second tube 
 and quickly settle. Add con. HC1 to the first tube till the precipitate 
 just dissolves, and pass an excess of H 2 S into the two solutions. Also 
 pass the gas into solutions of tin chloride and bismuth chloride. Fil- 
 ter off the precipitated antimony sulfide and wash on filter. Remove a 
 little of the precipitate Sb 2 S 3 to a dish and boil with a little con. HC1. 
 How may antimony be separated from arsenic? See 82. 
 
 Dissolve a portion of the sulfides of antimony and tin by warming 
 in a dish with yellow ammonium sulfide, and try to dissolve bismuth 
 sulfide in the same way. To the solutions of Sb and Sn sulfides add an 
 excess of dilute HC1 when the sulfides will be reprecipitated. How 
 may antimony, arsenic and tin in solution be separated from bismuth 
 and most other heavy metals? 
 
 To illustrate the separation of antimony and tin dissolve their sul- 
 fides together in hot con. HC1, boil for a few moments, dilute the so- 
 lution with twice its volume of water, put in "card teeth" or steel wool 
 and boil persistently. Filter off the black scales of Sb and to the fil- 
 trate 'add a solution of mercuric chloride. A white precipitate of mer- 
 curous chloride shows stannous chloride, SnCl 2 , was present. Dissolve 
 the black scales on the filter with a very little aqua regia, dilute with 
 20 times its volume of water and pass in H 2 S which will precipitate 
 antimony sulfide. 
 
 CARBOff. 
 
 87. In a test-tube held in a clamp with its mouth slightly lower 
 than the other end, heat a block of wood. Light the gas given off and 
 let any liquid distilled fall into a vessel not worth cleaning. Examine 
 
46 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 liquid and test for acid. When the volatile products cease to come off 
 examine the charcoal left in tube as to its form, structure, lightness. 
 
 In a small flask boil solutions of litmus, cochineal, and indigo with 
 powdered bone charcoal and filter. Use about 3 g. fresh charcoal with 
 25 c.c. of each solution. If any one of the filtrates is not clear return 
 it to the flask, boil again and filter. 
 
 90. Preparation of Carbon Dioxide: Provide a bottle or flask 
 
 with thistle tube and two-piece delivery tube as in 68. Put in 25-50 
 grams of, calcium carbonate, cover it and the tip of thistle tube with 
 water, and add a little con. HC1 from time to time to maintain a mod- 
 erate flow of COv. Three or four jars of gas may be collected in the 
 same way as chlorine. 
 
 Now pass the gas into a test tube half full of clear lime water till 
 the precipitate of normal calcium carbonate, CaCOs, formed at first 
 completely dissolves as it is changed into the acid carbonate, H 2 Ca 
 (CO 3 ) 2 . Reverse the reaction by boiling one half of the solution. To 
 the other lialf add lime water which will 'precipitate the normal car- 
 bonate. (See 116). 
 
 Pass CO 2 into 5 c.c. 3-normal sodium hydroxide, slowly enough to 
 make absorption apparent. Na 2 CO 3 is formed at first, but later the acid 
 carbonate, HNaCOs, which being sparingly soluble, is precipitated. Re- 
 verse the reaction to the Na 2 CO 3 by boiling. Test a few drops of the 
 solution and also small amounts of other carbonates for carbon diox- 
 ide thus : Place a little of the carbonate in one test tube 'and a few c.c. 
 clear lime water in a larger one. Add an excess of dil. HC1 to the car- 
 bonate a,nd pour the gas set free, but no liquid, into the lime water, 
 shielding the mouths of the tubes from air currents as shown in lec- 
 ture room. A white precipitate in the lime water shows carbon diox- 
 ide and a carbonate. 
 
 91. Introduce into a jar of carbon dioxide rapidly burning phos- 
 phorus. Place a bit of burning candle in a beaker and pour carbon 
 dioxide upon it from a jar. Which is more dense, CO 2 or air? 
 
 Determine whether a candle v/ill burn in the exhaled breath by 
 filling the lungs to capacity, exhaling through a tube into a jar, and 
 without delay lowering into it a bit of. burning candle. 
 
 92. Carbon dioxide may be prepared by heating acid sodium car- 
 bonate, "soda," or magnesite, MgCO* as in Fig. 13 but turning the de- 
 livery tube downward. Soda decomposes into the normal carbonate 
 so easily that a test tube may be used. For magnesite the iron tube 
 is better. In this case the reaction is similar to that in making lime 
 by heating CaCO 3 . 
 
 93. (a) Carbon Monoxide: Set up the apparatus as in Fig 18, 
 having the bottles two-thirds full of NaOH, one part of the solid to 
 three of water. Place in flask 20 grams oxalic acid and 40 c.c. con. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 47 
 
 HiS04. Prepare a tube with copper oxide as in Fig 5, if (b) is assigned. 
 
 Heat the flask so as to maintain a stream of gas slow enough to 
 permit the absorption in the bottles of the CO 2 which comes off with 
 the CO. 
 
 Fig. 18. 
 
 After a jar full of gas has come over, place delivery tube in a test-tube 
 of lime water. It should not become turbid. If turbid, decrease the 
 rate of flow. Collect two jars full of gas. Remove cover of one jar, 
 quickly pour in lime water, seal and shake. Burn the gas in the jar 
 and shake again. Apply a flame to the mouth of the other jar, add 
 lime water, seal and shake. What is formed by burning CO? 
 
 Disconnect flask from bottles and place its delivery tube in a test- 
 tube of lime water. What other gas is mixed with the carbon monox- 
 ide? To determine the amount of this gas fill a slender test-tube with 
 the mixture by displacement of air, slowly withdraw the delivery tube 
 and cover mouth of test tube with the wet finger. Place mouth of tube 
 under sodium hydroxide in porcelain dish and let it remain until the 
 volume of gas no longer decreases. Estimate the volume of remaining 
 gas as compared with the original volume of the mixture and test it 
 with a flame. 
 
 (b) Disconnect the delivery tube and substitute for it the tube pre- 
 pared with copper oxide. Support this tube on a ring, connect with a 
 test tube of lime water. Heat the copper oxide and pass over it a slow 
 stream of carbon monoxide. Use more oxalic acid in the generating 
 flask if necessary. Is the copper oxide reduced? What is formed by 
 its action on carbon monoxide? For the great importance of this pro- 
 
48 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 perty of carbon monoxide see the chemistry of the blast furnace in the 
 text book. 
 
 94. Methane, Marsh Gas: Mix intimately in a mortar 10 grams of 
 fused sodium acetate CH 3 C0 2 Na and 10 grams of soda lime. Place in a 
 retort of gas pipe, provided with a two-piece delivery tube and heat as 
 in the preparation of oxygen, collecting two jars of gas. Light the gas 
 and study the degree of luminosity of the flame. Hold the. flame under 
 the mouth of a dry bottle or jar and note appearance of moisture. Put 
 lime water in the jar and shake. Pass the gas through lime water in 
 a test tube. What evidence have you that methane contains hy- 
 drogen and carbon? 
 
 To the other jar of CH 4 add a little bromine water and shake. Does 
 methane absorb bromine? 
 
 95. Acetylene: Cuprous chloride is required, and should be 
 freshly prepared as in 126. For the prepaartion of acetylene fit up the 
 apparatus in Fig. 10, but substitute a small flask for the test tube. 
 Place in flask about 10 grams of calcium carbide, and by means of the 
 funnel and clamp let in upon the carbide water, a few drops at a time. 
 When the air has been expelled as shown by collecting and burning a 
 test tube of the acetylene, collect three bottles of the gas. Insert the 
 delivery tube into a solution of cuprous chloride made alkaline with 
 ammonium hydroxide. The red precipitate is cuprous acetylide, Cu 2 C:>. 
 In the same way precipitate silver acetylide from a solution of silver 
 nitrate made alkaline with ammonia. These acetylides are explosive 
 and should be washed into the sink before they become dry. 
 
 Insert the delivery tube into a test tube half full of bromine water 
 and let the gas run for several minutes. Also, shake in one of the bot- 
 tles a little bromine water with acetylene. Compare its action on bro- 
 mine to that of methane. 
 
 Uncover one bottle of gas, wait a moment, then apply a flame. Why 
 is there a bright flash followed by quiet combustion giving a very 
 smoky flame? 
 
 96. Flame. Regulate a Bunsen burner so as to produce a small 
 slightly luminous flame. Observe the three parts, lower and inner por- 
 tion which is non-luminous, luminous portion, outer non-luminous por- 
 tion. Compare with a candle flame. With a short piece of glass tub- 
 ing draw off gas from lower part of flame and burn it at the end of 
 the tube. Try the same with candle flame. Hold a piece of wire gauze 
 in candle flame just above the wick and observe interior. Can you light 
 unburned gas above the gauze? Hold a piece of paper for an instant 
 in the same position and examine under surface. Repeat experiments 
 with gauze and paper using the flame of a Bunsen burner. Also turn 
 on gas, hold gauze close to the burner tube and light gas above gauze. 
 Now move gauze to one side. 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 49 
 
 97. Oxidizing and Reducing Flames. Regulate a burner so as to 
 produce a small luminous flame, and find by trial the best position 
 for tip of blowpipe to produce a long slender blowpipe flame with 
 well-marked inner blue reducing and the outer non-luminous, oxidizing 
 portions. 
 
 Make a cavity with butt of pliers or a small coin in a piece of char- 
 coal, place in it a small quantity of lead oxide and heat in the reduc- 
 ing flame of the blowpipe. Continue until a globule of lead remains. 
 Caution: After using charcoal with the blowpipe always extinguish 
 any fire by holding it under the faucet and then return it to the char- 
 coal tray. 
 
 98. Fermentation of Glucose. In a 500 cc. flask dissolve 25 grams 
 of glucose, 1 gram each of sodium-potassium tartrate, ammonium nitrate 
 and soclic phosfate in about 250 cc. of hydrant water. Rub with a little 
 water in a mortar one-tenth of a cake of yeast and wash into flask. 
 Connect the flask with a gas washing bottle nearly full of clear lime 
 water, set in a warm place and let remain three or four days. Observe 
 from time to time the carbon dioxide given off. Take out a drop of the 
 liquid, place on a slide, cover and examine for cells of the yeast plant 
 with a microscope. 
 
 Arrange the flask as in fig. 6 and distill off 20 c.c. keeping dis- 
 tillate well cooled. Clean the flask, put in the 20 c.c. and distill off 
 about 5 c.c. Test this second distillate for alcohol, first testing for it 
 in a known solution as follows: to about 5 c.c. water add about 1 c.c. 
 alcohol and a few crystals of iodine, and shake. Warm the tube and 
 add gradually a solution of sodium carbonate till the iodine disappears, 
 lodoform will appear, at any rate on cooling. Let a few drops cool on 
 a watch glass and examine with the low power of a microscope. . Now 
 examine for alcohol the distillate above. 
 
 99. Reducing Power of Sugars. Invert Sugar. Fehling's solution 
 is a mixture of equal volumes of two solutions, one containing 34.639 
 grams of copper sulfate to 500 c.c. and the other 173 grams sodium pot- 
 assium tartrate and 60 grams sodium hydroxide in 500 c.c. The cop- 
 per in 1 c.c. of the mixture is reduced by 0.005 gram of glucose. 
 
 To about 10 cc. Fehling's solution add enough glucose to precipi- 
 tate all the copper on boiling as cuprous oxide, Cu 2 O, which will settle, 
 leaving the clear liquid above. In a fresh mixture try a little milk 
 sugar. Try cane sugar. 
 
 Now dissolve about a gram of cane sugar in 100 c.c. of water, add a 
 few drops of con. HC1, heat to 70 degrees and let cool. Boil 10 c.c. of 
 this solution of invert sugar with 20 c.c. of Fehling's solution. 
 
 100. Esters, Acetic Ester: Acids act upon alcohols forming esters 
 and water much as they act upon bases forming salts 
 
50 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 and water. Acetic acid and alcohol act thus: CH 3 CO 2 H+C2H 5 OH= 
 CH 3 CO 2 C 2 H S +H 2 O. 
 
 Use the apparatus shown in Fig. 6. With the delivery tube in the 
 test tube pour into test tube 20 c.c. water and mark the level with 
 gummed paper. Remove the water. Place in the flask 15 c.c. glacial 
 acetic acid, 15 c.c. alcohol and 5 c.c. con. sulfuric acid. Warm the flask 
 with a very small flame so that only a few drops shall distill over in 
 10 minutes. Now increase the heat and distill to the 20 clc. mark. In- 
 stead of measuring the distillate a two hole stopper and thermometer 
 may be used, and the distillation continued till it reads 95. Its bulb 
 should be in the vapor above the liquid. 
 
 The ester contains acetic acid and alcohol. To remove most of 
 these shake it persistently with twice its volume of water in a small 
 flask covered with the thumb. On standing the ester rises and forms 
 the top layer. Separate it from the water by using a separatory fun- 
 nel, or the funnel with pinch cock and exit tube as shown in Fig. 10. 
 Measure the purified ester. 
 
 In a test tube shake about 2 c.c. of your ester with an excess of 
 NaOH till it all disappears; that is, till it is all "saponified" forming 
 sodium acetate and alcohol. 
 
 Note odor of acetic ester. Upon this is based a good test for ace- 
 tic acid or its salts. To a little solid sodium acetate in a test tube add 
 a few drops of alcohol and a few of con. sulfuric acid and note odor. 
 Another test is this: To a neutral solution of an acetate add 3 drops 
 of ferric chloride. A red color should be obtained, and on boiling a 
 brown flocculent precipitate. 
 
 101. Saponification, Soap Making: In a dish or beaker place 
 five grams of olive oil or lard or tallow, add 15 c.c. of alcohol and 5 c.c. 
 of sodium hydroxide solution of concentration one to four. Do not use 
 a pipette for NaOH. With occasional stirring heat on a water or steam 
 bath, for at least half an hour, but better, an hour. All the alcohol 
 and most of the water should be evaporated. Examine the soap when 
 cold. If the saponification is complete a small portion should entire- 
 ly dissolve in water, save a slight milkiness. Dissolve more of the 
 soap and determine whether it will form a lather with soft water. Try 
 blowing soap bubbles with the solution. Determine whether the feel- 
 ing is that given by ordinary soap. 
 
 Filter some of the soap solution. Pour a little into distilled water, 
 into hydrant water, into solutions of calcium and magnesium sulfates. 
 Shake each tube and determine in which ones a persistent lather is 
 formed. To those in which it is not formed add more soap solution and 
 shake again, until a lather persists. Calcium and magnesium com- 
 pounds make water "hard." Why do such waters require more soap? 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 51 
 
 To a concentrated solution of soap add dilute hydrochloric acid in 
 excess. The precipitate consists of a mixture of organic acids of 
 which stearic acid, Cn H 33 COOH, is representative. The oils and fars 
 mentioned above consist of mixtures of what are known as esters of 
 these acids. The most common perhaps is stearin, (Ci-HusCOOhCsHs. 
 Write the equation for the saponification of stearin with sodium hy- 
 droxide; also, write equation for action of soap on solution of calcium 
 sulfate. 
 
 SILICON. 
 
 102. Silicic Acid, To 10 c.c. of a solution of sodium silicate in a 
 small beaker add con. HC1 drop by drop stirring for a few moments 
 after each addition. Silicic acid will separate as a jelly. Collect the 
 silicic acid on a filter, wash with water, transfer to a crucible, dry 
 over flame and finally ignite. When cold try to dissolve the residue 
 in water, in hydrochloric acid. 
 
 103. Boric Acid: Place 15 g. borax in 50 c.c. water in a beaker. 
 Heat till dissolved and add 15 c.c. con. HC1 with stirring and allow the 
 liquid to cool thoroughly. Filter off and examine the boric acid. Dis- 
 solve a little of the substance in alcohol in a dish, set fire to alcohol 
 and observe color of the flame. Moisten a strip of turmeric paper with 
 the solution from the boric acid. Observe color, then treat the paper 
 with a solution of sodium carbonate, let dry, best on steam bath, and 
 note color. 
 
 QUALITATIVE TESTS FOR THE COMMON ACIDS 
 
 106. The substance given for analysis may be in solution. If not, 
 dissolve in water, heating if needed. If not soluble in hot water use di- 
 lute nitric acid, and heat. If gas is given off one or more acids of 
 group (1) are present, and the special tests may be applied at once. 
 
 Group (1) : To a small portion of the solution add dilute HNO 3 in 
 excess if not already added, and heat. If gas is set free try to identify 
 it by odor, color anft tests. Only CO 2 , S0 2 , H 2 S, NO-, are likely to occur. 
 Make special tests for the following acids as given in the sections in- 
 dicated by the numbers: H 2 CO 3 (90), H 2 S0 3 (71), H 2 S 2 O 3 (73), H 2 S (68), 
 HN0 2 (54b). 
 
 Group (2) : To another portion of the solution add dilute nitric 
 acid in excess if not already present and if members of group (1) are 
 present boil to expel any gas. To a portion of the boiled solution add 
 barium chloride in excess. A white precipitate shows H 2 SO 4 (71). 
 Filter it off and make nitrate alkaline with ammonia. A yellow pre- 
 cipitate indicates H 2 CrO4 (140). A white precipitate indicates one or 
 more of the following: H 3 PO.,-(80), H 3 AsO 4 . (83), HsAs0 3 (83), 
 H 3 BO 3 (103), HF (38), H 2 C 2 O 4 (see following): To another portion of 
 
52 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 the boiled solution add ammonia in excess, then CaCl 2 . An immediate 
 precipitate may mean any of the acids of this group save sulfuric. Add 
 an excess of acetic acid. If the precipitate is insoluble, oxalic acid or 
 hydrofluoric or both are present. To a third portion of the boiled solu- 
 tion add manganese dioxide, which will give C0 2 if oxalic acid is pres- 
 ent. It is well in any case to make the special test for boric acid, since 
 its barium and calcium salts are distinctly soluble. 
 
 Group (3) To the original solution add silver nitrate. If a pre- 
 cipitate is formed add an excess of dilute nitric acid. If the precipi- 
 tate all dissolves this indicates some acid in previous groups. If in- 
 soluble, one or more of the following are present: HC1 (37), HBr (37), 
 HI (37), H 4 Fe(CN) 6 (142), H s Fe(CN) 6 (142). A pure white precipitate 
 shows only HC1. If the precipitate is colored proceed to the special 
 tests for the others. For halogen acids see also 127b. 
 
 Group (4) Nitric and acetic acids must always be tested for if 
 the substance is soluble in water. For HNO 3 see (56b) , and for acetic, 
 (100). 
 
 SODIUM, POTASSIUM, LITHIUM. 
 
 107. Through a spectroscope suitably adjusted by the instructor 
 examine a flame colored by a sodium salt, and locate the sodium line on 
 the scale. Locate in the same way the red line given by potassium, 
 and that given by lithium. Draw a millimeter scale in your note book 
 and place the lines in their correct positions. If the spectroscope has 
 no scale, estimate as closely as possible the relative positions of the 
 three lines. Note that any chemical will show some sodium. Its line 
 is no evidence that sodium is present in considerable quantity. 
 
 108. What is the action of sodium on water? Standing well back 
 drop a bit of potassium into water in a bottle or beaker. How does the 
 result differ from that given by Na? 
 
 To reduce the violence of the action of the alkali metals on water 
 their alloys with some other metal are often used. Drop into water 
 "hydrone" which is an alloy of sodium and lead. Also try sodium 
 amalgam, an alloy of sodium and mercury. When the action of the 
 latter is over put the mercury in a dish provided for the purpose. 
 Never put mercury into sinks. Why? Test the solution in each case 
 with red litmus or turmeric paper. Test soapy feeling of the solution 
 between thumb and finger. 
 
 109. Preparation of Sodium Hydroxide: Dissolve 5 grams of so- 
 dium carbonate in 75 c.c. of water in a porcelain dish and reserve 5 c.c. 
 Heat the remainder nearly to boiling, and stir in a little at a time 5 
 grams of slaked lime. Boil gently for several minutes, replacing the 
 water evaporated, let settle and filter off the solution. If it destroys 
 the filter, let cool, add a little water and use another filter. To prove 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 53 
 
 the reaction, Na 2 CO 3 +Ca(OH) 2 =CaCO 3 +2NaOH, proceed as follows: 
 Wash the insoluble residue in the dish twice by filling nearly full of 
 water, stirring, letting the substance settle and pouring off the water. 
 Now put a little water on the insoluble substance in dish, and place 
 about the same amounts of lime and water in another dish and pour 
 upon each about 10 c.c. of dilute HC1. Compare the amounts of carbon 
 dioxide given off. Now treat the 5 c.c. reserved solution of sodium car- 
 bonate and the same volume of the filtrate with the same volumes of 
 HC1 and compare the gas given off. You started with a soluble car- 
 bonate and a very sparingly soluble hydroxide, "slaked lime." What 
 did you obtain by the reaction? 
 
 110. Comparative Tests of Crude and Pure Sodium Hydroxides: 
 Dissolve 1-2 grams of crude and pure NaOH, each in 10 c.c. of distilled 
 water. Make each acid throughout with dilute nitric acid and 
 warm. If either gives off gas test for carbon dioxide as in 90. Test 
 small portions of each acidified solution for Cl with silver nitrate, and 
 other portions for SO 4 with BaCl 2 . Test other portions for iron by add- 
 ing to each a few drops of potassium ferrocyanide and potassium ferri- 
 cyanide, which will give a blue color if iron is present. 
 
 HI. Purification of Common Salt: If practicable use the crude 
 rock salt of the feed store, but ordinary salt will do. Dissolve about 5 
 grams in 25 c.c. of water. To one third add dilute HC1 and test for 
 SO 4 , and preserve the tube and contents. To another portion add a 
 little ammonium chloride, make alkaline with ammonium hydroxide 
 and add ammonium carbonate. The precipitate is calcium carbonate. 
 Filter it off and to the filtrate add sodium phosphate which will give a 
 precipitate on standing if magnesium is present. See 80. Save this 
 tube and contents. 
 
 Make a fully saturated solution of the crude salt, first reducing it 
 to powder if rock salt is used, and shaking a long time with water. 
 Filter if necessary and to the solution add an equal volume of pure 
 con. HC1. This will precipitate most of the salt. If gasous HC1 were 
 added more salt would precipitate. Filter off the salt and wash with 
 three small portions of water. Dissolve some of this salt in water, 
 test for SO 4 , calcium, magnesium, as above and compare with the 
 results of these tests with crude salt. 
 
 112. An Acid Salt, Acid Potassium Tartrate, Cream of Tartar: 
 This is one of the few slightly soluble salts of potassium. As the term 
 is commonly used there are no insoluble salts of Na or K. Dissolve 
 about 4 grams of pure, dry potassium carbonate in 25 c.c. of water and 
 10 grams of tartaric acid in 50 c.c. of water, measuring the latter so- 
 lution. To the carbonate solution add one drop of methyl orange, then 
 add cautiously the acid solution till a faint red color is obtained. This 
 gives the soluble normal salt, K 2 C4H 4 O 6 . Note what volume of the acid 
 
54 LABORATORY MANUAL OP GENERAL CHEMISTRY 
 
 solution was added, then add as much more, stir and let stand some 
 time. The precipitate .is the acid tartrate, cream of tartar. Which 
 salt is least soluble? 
 
 Filter off the cream of tartar, let it dry on the filter and weigh 
 with a balanced filter. Calculate the weight of acid sodium carbonate 
 to mix with it to make one kind of baking powder which when wet acts 
 thus: HKC4H 4 O 6 +HNaCO3=NaKC i H 4 O c +H 2 O+CO 2 . What makes the 
 bread rise when baking power is used? Try your baking powder in a 
 tube with water and test for COz. 
 
 113. Potassium Nitrate from Sodium Nitrate: Dissolve in 50 c.c. 
 H 2 O, 25 g. sodium nitrate and the calculated amount of potassium chlo- 
 ride required in the reaction 
 
 KCl+NaN0 3 =KNO 3 +NaCl. 
 
 Evaporate to one-half the volume, let the separated salt settle, and de- 
 cant the clear, hot liquid into a beaker, press solid with spatula and 
 let liquid run into beaker. This solution should turn solid when cold. 
 Transfer this to a filter and let drain. Press solid between folds of 
 filter paper, dissolve in least hot water and let crystallize. Transfer 
 crystals to filter and let dry. 
 
 114. Qualitative Analysis: Determine first the presence of the 
 ammonium radical and the alkali metals in known substances, and 
 then their presence or absence in "unknowns," using the scheme as 
 given 152, Group V. In the initial work fresh substances will be ex- 
 amined, and of course what is said of filtrates from previous groups 
 and their preparation for analysis does not apply. 
 
 BARIUM, STRONTIUM, CALCIUM, (MAGNESIUM). 
 
 115. Upon a piece of quick lime drop water as Icng as it is taken 
 up, place it in a dish and observe from time to time. After it has be- 
 come powdery place some of the "slaked" lime in a jar of water, shake 
 it thoroughly and let settle. Filter a portion of the nearly clear lime 
 water placing the funnel in a flask to protect from the carbon dioxide 
 of the air. Test the clear solution with turmeric paper. To portions 
 of the lime water add one-third of their volumes of ferric chloride and 
 magnesium chloride respectively. The precipitates are hydroxides 
 of the metals. Test the alkalinity of barium hydroxide and its action 
 on solutions of the same metals. 
 
 116. Pass carbon dioxide into 25 c.c. of clear lime water until the 
 precipitate ot calcium carbonate, CaCOs, dissolves, forming the acid 
 carbonate, H.CaCCOaK The latter is the chief substance that gives 
 "temporary hard water." To a small portion of the solution add clear 
 lime water. How may lime soften temporary hard water? Will lime 
 also remove magnesium from water? Boil another portion of the so- 
 
LABORATORY MANUAL, OF GENERAL CHEMISTRY 55 
 
 lution which will reverse to the left the reaction which occurred with 
 CO 2 in excess. 
 
 CaC0 3 +H 2 0+C0 2 --=(reversibly)H 2 Ca(C0 3 ) 2 . 
 
 117. Pure Calcium Chloride: Pour off the liquid from flask in 
 which carbon dioxide was made, make it alkaline with milk of lime ob- 
 tained by shaking slaked lime and water in a jar and pouring at once. 
 Let settle and filter. How does this remove iron and magnesium? 
 Pass into the filtrate C0 2 and boil. Why? Filter again if necessary 
 and evaporate to dryness in a porcelain dish and heat. Expose a little 
 of the calcium chloride to air till next period and observe again. Is it 
 deliquescent? Dissolve the remainder in about 10 times it weight of 
 water, and use as calcium chloride solution. 
 
 118 Comparative Solubilities of Salts of Ba, Sr, Ca, Mg: Carbon- 
 ates: To solutions of Ba, Sr, Ca and Mg chlorides from shelf add equal 
 volumes of water then to each about one-fifth of its volume of ammo- 
 nium chloride, and finally to each, ammonium carbonate. What com- 
 pounds are precipitated? How could magnesium be separated from 
 the other three metals? Add to its solution sodium phosphate, and 
 see the tests for phosphoric acid (80) arid Mg (128). 
 
 119. Chromates: To solutions as under carbonates, but omitting 
 Mg, add a little acetic acid then a solution of pure potassium chro- 
 rhate, or dichromate. How could barium be separated from strontium 
 and calcium? 
 
 120. Suit* ates : To solutions of Ba, Sr, and Ca chlorides add a so- 
 lution of magnesium sulfate and let stand for a few moments. How 
 does this prove that Ba, Sr, and Ca sulfates are less soluble that mag- 
 nesium sulfate? To fresh solutions of Ba and Sr chlorides add a solu- 
 tion of calcium sulfate. How do the results show that Ba and Sr sul- 
 fates are less soluble than calcium sulfate? To a solution of barium 
 chloride add a solution of strontium sulfate and let stand a short time. 
 How do we know that barium sulfate is less soluble than strontium 
 sulfate? Arrange the sulfates in the order of their increasing solu- 
 bility in water. 
 
 121. Calcium Oxalate: To a solution of calcium chloride add an' 
 excess of ammonium carbonate, let stand a few moments and filter. To 
 the filtrate and also to a solution of calcium sulfate add a solution of 
 ammonium oxalate and let stand half an hour. State how you know 
 that calcium oxalate is less soluble than calcium carbonate or calcium 
 sulfate. Devise a scheme for the separation of Ba, Ca, Sr and Mg. 
 
 Calcium sulfate is present in many natural waters and causes 
 "permanent hardness" in the sense that it is not precipitated by boil- 
 ing, though on concentration by evaporation, it forms a hard deposit 
 on the boiler. To calcium sulfate solution add a solution of sodium 
 carbonate. What two chemicals may be added to soften water showing 
 
56 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 both temporary and permanent hardness? 
 
 Test the water of the laboratory for both sorts of hardness. 
 
 122. Make analyses of a solution containing barium, strontium, 
 calcium and magnesium according to the directions of Group IV; also, 
 analyses of unknown solutions or solids which may contain these me- 
 tals, and solutions or solids which may contain also metals of Group 
 V and the ammonium radical. 
 
 DECIHORMAL SOLUTIONS, VOLUMETRIC ANALYSIS. 
 
 123. Weigh accurately a small dish or beaker, add 2.65 grams to 
 the weights and exactly balance with pure sodium carbonate. Dis- 
 solve the carbonate in water, transfer with rinsings of dish to a half 
 liter flask, using a funnel and taking care that none of the solution is 
 lost. Do not even lose some by removal on the stirring rod. Make up 
 the volume to the mark with water, and mix by placing the thumb 
 over mouth of flask and inverting several times. This is a decinormal 
 solution of Na 2 CO 3 . Why? Transfer this solution to a bottle or larger 
 flask and wash the graduated flask. 
 
 Measure in a small cylinder 7.5 c.c. pure con. HC1 and dilute it 10 
 700 c.c. and mix well. Fill a buret with the acid solution to well above 
 the zero, fill the tip of buret and bring the surface to or below the 
 zero. Read accurately at the lowest point of the meniscus. With a pipet 
 (see 4) place 20 or 25 c.c. according to capacity of the pipet, of the so- 
 dium carbonate solution in a dish or beaker, add to it 2 drops of methyl 
 orange. For comparison it is well to place beside it about 50 c.c. of 
 water and add- to it two drops of the indicator. From the buret run in 
 the acid as rapidly as you wish to about 15 c.c. then a few drops at a time 
 with stirring till the solution becomes faintly red as shown by compar- 
 ison with the indicator in water. Make another titration, which should 
 agree within a few tenths with the first. 
 
 Divide the volume of the alkali by that of the acid which gives the 
 decinormal concentration factor of the acid. Measure 500 c.c. of the 
 acid and multiply by the factor, which will give the total volume to 
 which the 500 must be made up with water to become decinormal. 
 Why? Transfer the 500 c.c. to a larger vessel, add the necessary water, 
 mix and titrate again against the alkali. They should neutralize each 
 other volume for volume. 
 
 With these two standard deci-normal solutions the concentration 
 of any other acid or alkali may be determined, or solutions of desired 
 concentration may be made by the same method used in making the 
 acid solution. 
 
 124. In titrating weak acids methyl orange cannot be used ; see 
 next section. One must use a very weakly acid or neutral indica- 
 tor, such as litmus, phenoltalein, congo red. But these are affected by 
 
LABORATORY MANUAL OP GENERAL CHEMISTRY 57 
 
 the carbonic acid from the carbonate and one must titrate the solution 
 at the boiling point or use NaOH free from carbonate. Find the per 
 cent of acid in vinegar by running it from a buret into a measured 
 volume of the carbonate boiling and containing a few drops of phenol- 
 talein, or use cold NaOH supplied by the instructor, instead of Na 2 COs. 
 Methyl orange and phenolphthalein are weak, complicated organic 
 acids. The following will give a correct, general idea of their action: 
 Any strong acid sets free any weak acid from its salt. Let NaR be 
 such a salt where R is the negative radical. Then, 
 
 In the case of phenolphthalein R" gives the red color, while HR is 
 colorless. Phenolphthalein is such a weak acid that even carbonic 
 acid is strong enough to act in the same way as HC1 in the equation 
 and form HR. Hence when the stage in the titration represented by 
 HNaCOs is passed, and H 2 CO a is formed this acts as a relatively strong 
 acid, forms HR and thus destroys the color. On the other hand methyl 
 orange is a stronger acid than carbonic and the yellow color of its 
 negative ion in Na + +R~ persists till there is a slight excess of HC1, 
 when red HR is formed. 
 
 125. The facts stated in 124 are well illustrated by titrating in the 
 same solution both normal and bicarbonate by using different indica- 
 tors, as fallows: 
 
 Dissolve about 0.2 gram of normal sodium carbonate in 25 c.c. of 
 water, without heating, add a few drops of phenoltalein, fill a buret 
 with your deci-normal acid, read and run into the carbonate solution 
 drop by drop near the end, till the pink color just disappears. The 
 carbonate is now all HNaCOs. Add a few drops of methyl orange, read 
 the buret and run in the acid till the solution takes on a faint tinge of 
 red, using the indicator in water for comparison as in 123. Read again 
 and compare the volumes required to change the normal to the bicar- 
 bonate, and to neutralize the latter. ' Why are they approximately 
 equal ? Why are these carbonates alkaline, having no ion OH ? See 6 1. 
 
 COPPER, SILVER. 
 
 Copper. 
 
 126. (a) Dissolve 5 grams copper chloride in 10 c.c. of con. HC1 
 and 10 c.c. water; or (b) prepare a solution of the copper chloride by 
 dissolving 5 grams copper sulfate and 2.5 grams common salt by heat- 
 ing with 10 c.c. of water in a test tube. When dissolved set the tube 
 in cold water for several minutes. Pour off the solution from the sep- 
 arated sodium sulfate. Why is this formed? Now add to the solution 
 10 cc. con. HC1, let stand a few moments and filter off the salt. Why 
 is salt thus formed? Whether (a) or (b) boil very gently the solu- 
 
58 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 tion with 5 grams finely divided copper in a small flask, replacing 
 evaporated liquid if necessary with dil. HC1. Continue till colorless 
 or till a few drops poured into water gives no blue color. You now 
 have HCuCl 2 . Pour a part into water which decomposes it giving in- 
 soluble CuCl. Does this dissolve in ammonia? Compare with silver 
 chloride. Put the ammonia solution in a white dish, stir and note that 
 the cuprous ion is rapidly oxidized to cupric ion as shown by increas- 
 ing blue color. 
 
 To a portion of the liquid from flask add an excess of NaOH. What 
 is the red compound? See Fehling's solution (99). Heat a little of the 
 white precipitate with Br water and give result in terms of ion formed. 
 Expose some of the white CuCl on the filter to sunlight and note result 
 after an hour. 
 
 There is no cupric but only cuprous iodide. To a few c.c. copper 
 sulfate solution add a little potassium iodide solution, and test the so- 
 lution for free iodine with starch paper. Add Ic.c. carbon disulfide 
 shake and let settle. Note color of the carbon disulfide. 
 
 Determine whether dilute sulfuric, hydrochloric and nitric acids 
 give H by their action on copper and cadmium and explain results. 
 
 (c) Tests for Copper: To half a test tube of water add a drop of 
 solution of any cupric salt, and make alkaline with ammonia. The 
 blue color is due to the complex ion Cu(NH 3 )4 ++ . To a like dilute solu- 
 tion of copper add a few drops of acid and a little potassium ferrocy- 
 anide. Dilute equal parts of each solution till the colors are just vis- 
 ible and state which test is the more delicate. 
 
 (d) To dilute solutions of copper, cadmium and zinc preferably 
 chlorides add H 2 S in excess. Let settle and pour off most of the liquid 
 in each case and add equal volumes of dil. HC1. Are the reactions 
 with H 2 S reversible? How may Cu and Zn be separated? To solu- 
 tions of the same metals add an excess of NaOH and heat to boiling. 
 For explanation see 127a and 128c. How may copper be separated 
 from the other two elements? 
 
 To solutions of copper and cadmium add an excess of ammonia. 
 Add to the copper tube KCN (dangerous) drop by drop till the blue 
 color disappears. Add the same volume of KCN to the tube contain- 
 ing Cd. Now pass H 2 S into each tube. How may Cu and Cd be separ- 
 ated? 
 
 Silver. 
 
 127. (a) To a little silver nitrate solution add drop by drop am- 
 monia solution till the small amount of silver oxide at first formed is 
 dissolved, forming the complex ion Ag(NH 3 ) 2 + . Add a little more 
 ammonia and try to precipitate silver chloride with a small amount of 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 59 
 
 NaCl solution. Now add an excess of dilute nitric acid. What is pre- 
 cipitated and why? 
 
 To another portion of silver nitrate solution add sodium hydroxide 
 solution which will precipitate mainly Ag 2 O, but it is alkaline and be- 
 haves as though partially hydrated. A similar copper hydrated oxiclo 
 is formed by adding an excess of NaOH to a solution of copper sulfate 
 and heating to boiling. Try it. The compound is Cu(OH) 2 .2CuO. Add 
 ammonia to the tube containing the silver oxide till it just dissolves 
 and then H 2 O 2 , which will give metallic silver. Prepare the same sort 
 of a solution of silver oxide in ammonia, that is, containing the ion 
 Ag(NH 3 ) 2 + , add 1 gram sodium potassium tartrate dissolved in a little 
 water, warm the test tube and let stand. If silver is not deposited on 
 the tube warm again. Note analogy of the complex silver ammonia ion 
 to that of copper, Cu(NH 3 )4 ++ which is deep blue. 
 
 (b) To four tubes containing silver nitrate add respectively a 
 solution of a chloride, a bromide, an iodide and to the fourth drop by 
 drop a solution of KCN (caution) , till the silver cyanide at first formed 
 is dissolved. Treat a little of each of the halides of silver with an ex- 
 cess of ammonia. Which are dissolved? Treat other small portions 
 with a solution of sodium thiosulfate and shake till dissolved. Treat 
 yet other portions with a solution of KCN till dissolved. How 1 could 
 you distinguish the three halides of silver by their color? How dis- 
 tinguish by their solubility in ammonia? What is formed when sil- 
 ver chloride dissolves in ammonia? What when it dissolves in KCN? 
 Expose a little chloride to sunlight and observe color after a few min- 
 utes. Compare cuprous chloride and silver chloride as to the effects 
 on them of sunlight and ammonia solution. 
 
 (c) To solutions of sodium phosphate, potassium chromate, sodium 
 arsenite and sodium arsenate add silver nitrate. Try to dissolve por- 
 tions of each precipitate in ammonia and in dil. nitric acid. Make 
 careful records of all results for they are to be used in the qualitative 
 testing for acid radicals. 
 
 MAGNESIUM, ZIJC, CADMIUM, MERCURY. 
 
 128. (a) To solutions of magnesium, zinc and cadmium salts, 
 preferably sulfates or chlorides add an excess of ammonium hydrox- 
 ide. Some zinc and cadmium hydroxide are precipitated but in an ex- 
 cess of ammonia they form the complex ions Zn(NH 3 )4 ++ and 
 Cd(NH 3 )4 4+ which are soluble. To show that only a part of the Mg is 
 precipitated as hydroxide, filter it off and add acid sodium phosphate 
 to the filtrate when more magnesium as NH 4 MgPO4 will be precipi- 
 tated. The prevention of complete precipitation as Mg(OH) 2 by am- 
 monium hydroxide is due to the necessary accumulation of highly ion- 
 ized ammonium salt as the reaction progresses. To show this first 
 
60 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 add to a solution of magnesium salt ammonium chloride solution and 
 then ammonium hydroxide. Also, precipitate Mg(OH) 2 with ammonia 
 and then add ammonium chloride. In the first case no magnesium 
 hydroxide was precipitated and in the second it dissolved. 
 
 The action of the ammonium salt which is highly ionized is to 
 supply the common ion NH 4 which forces the already slight dissociation 
 of ammonium hydroxide to the left in NH 4 OH=(reversibly)NHr+OH- 
 till there are not enough OH ions to form sufficient Mg(OH) 2 to exceed 
 its solubility limit. Furthermore, the NH 4 ions from the ammonium 
 chloride or other ammonium salt unite with the OH ions associated 
 with the Mg in solution to form the undissociated NH 4 OH. This latter 
 action is of the same sort as in the neutralization of a base by an acid 
 in which the ion H of the acid unites with the OH of the base to form 
 undissociated water. 
 
 This explanation applies to the solubility of several other hydrox- 
 ides when ammonium salts are added and will be referred to later. 
 
 Prove that acid sodium phosphate alone will not completely pre- 
 cipitate Mg by adding to a magnesium solution an excess of phosphate, 
 filtering and then adding to the filtrate ammonium chloride and am- 
 monia. Make the corresponding salt of zinc, NH 4 ZnPO 4 . 
 
 (b) Burn a little Mg ribbon, place some of the oxide on moist 
 turmeric paper and state whether it is alkaline. Burn a little zinc 
 dust by heating and stirring in an iron crucible, and determine wheth- 
 er it has any alkaline property. 
 
 (c) Precipitate the hydroxides of Mg, Zn and Cd by adding to 
 their solutions NaOH a few drops at a time. Now add more NaOH to 
 half of each and determine which are soluble in an excess. Dissolve 
 the other half of each hydroxide by adding any dilute acid. With so- 
 dium hydroxide in excess zinc and cadmium hydroxides form Na 2 Zn02 
 and Na 2 CdO 2 , which resemble ordinary salts in form and properties. 
 The hydroxide of Zn and Cd and many other elements are "ampho- 
 teric"; that is, they act like bases toward strong acids, and like weak 
 acids toward strong bases. On the basis of the ion theory explain the 
 dissolving of Mg(OH) 2 by HC1. 
 
 (d) To solutions of Mg, Zn, and Cd add H 2 S for several minutes. 
 Which give sulfides? Shake and filter off half of each precipitate 
 and to the filtrates add ammonia. Did the H 2 S completely precipitate 
 both the Zn and Cd as sulfides? To the other half of each precipitate 
 obtained with H 2 S alone add dilute HC1 and finally con. HC1 if needed 
 to dissolve all the sulfides. Are the reactions of Cd and Zn salts with 
 H 2 S both reversible? Which one is most easily reversed? How may 
 Zn, Mg and Cd be separated? 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 61 
 
 MERCURY 
 
 129. Place 10 grams mercury in 4 c.c. con. nitric acid diluted with 
 the same volume of water and let stand a day. Pour off liquid and 
 dissolve the solid in water adding a little dilute nitric acid, letting the 
 metallic mercury remain. This is a solution of mercurous nitrate, (a) . 
 Dissolve about 2 grams mercury in a few c.c. con. nitric acid, heat and 
 if necessary add more nitric acid and heat till a drop of the solution in 
 water gives no precipitate with HC1. Now dilute with about 50 c.c. 
 water. This is mercuric nitrate, (b). 
 
 Treat small portions of (a) and (b) with NaOH in excess which 
 gives Hg 2 O and HgO and not the hydroxides. Compare the result with 
 that obtained with silver nitrate and NaOH, with a cupric compound 
 and NaOH cold and after boiling. Try ammonia on the solutions (a) 
 and (b). No oxides or hydroxides are formed but with (a) Hg and 
 Hg 2 N(NO 3 ). With (b) the same compound but no free Hg. 
 
 Treat a little of (a) with dil. HC1 and determine whether the HgCl 
 is soluble in an excess of HC1 or in nitric acid. Try the action of am- 
 monia which gives black Hg. HgNH 2 Cl. How may mercury in the mer- 
 curous condition be separated from silver? Oxidize a little of solution 
 (a) to (b) by adding Br water till the red color persists, boiling out 
 excess of Br. Prove that only mercuric mercury is present by adding 
 dil. HC1. 
 
 To small portions of (a) and (b) add a solution of KI drop by drop 
 giving green Hgl and red HgL. To the latter add an excess of KI 
 which will dissolve forming a double salt, HgI 2 2KI. 
 
 To a very dilute solution of stannous chloride add a solution of 
 mercuric chloride which gives HgCl while SnCh, stannic chloride, re- 
 mains in solution. This is a good test for either Hg or Sn. What must 
 be the valence of each ion when tested for? 
 
 To (a) and (b), and to solutions of Pb, Cu, Cd add an excess of 
 H 2 S. Let settle, pour off the liquid in each case and try to dissolve the 
 precipitates by boiling with dilute nitric acid. How may mercury be 
 separated from the other metals? Determine whether HgS will dis- 
 solve in ammonium sulfide. How may mercury be separated from As, 
 and Sb. 
 
 TIN. 
 
 130. Dissolve most of 2 grams of tin by heating with 10 c.c. con. 
 HC1, best in a small flask on the water bath. Pour off half of the so- 
 lution into a dish and add 50 c.c. of water to the flask and use the so- 
 lution as stannous chloride, (a). Heat to boiling the solution in the 
 dish and add con. nitric acid a few drops at a time till a drop in a 
 
62 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 little water gives no precipitate with mercuric chloride. Add 50 c.c. of 
 water forming a suitable solution of stannic chloride, (b). 
 
 To small portions of (a) and (b) add a few drops of NaOH, then 
 add in excess. Are these hydroxides "amphoteric" ? To fresh portions 
 of the tin solutions add yellow ammonium sulfide, at first only a few 
 drops, then an excess with heating, but do not boil. The SnS and SnS 2 
 at first formed, should dissolve. Now add HC1 in excess which will 
 reprecipitate the tin sulfide, SnS 2 from each solution. What other sul- 
 fides dissolve in ammonium sulfide? How may the tin sulfides be sep- 
 arated from those of copper, lead, mercury? 
 
 To a little of (a) add a solution of mercuric chloride which will 
 give insoluble HgCl. Try (b) with mercury chloride. Now reduce to 
 the stannous condition the Sn in a portion of (b) by persistent heat- 
 ing with finely divided iron, filter and add mercuric chloride. This is 
 a good test for either Hg or tin. In what condition of oxidation must 
 each be? 
 
 To a portion of (a) add a few drops of gold chloride which will 
 give colloidal gold, the purple of Cassius, a good test for gold. 
 
 LEAD. 
 
 131. To a dilute solution of lead acetate or nitrate add an excess 
 of NaCl solution. Before the lead chloride settles pour off one half 
 and boil, adding more water if necessary and boiling till it all dis- 
 solves. Let the other half stand 5-10 minutes, filter and to the filtrate 
 add dil. sulfuric acid. What proof here that lead sulfate is less soluble 
 than lead chloride? Filter off the lead sulfate and pass into nitrate 
 H 2 S. Which is less soluble, lead sulfate or lead sulfide? From a dilute 
 solution precipitate lead chloride and prove that it is soluble in HC1 if 
 added in large excess. Prove in the same way that lead sulfate is 
 soluble in nitric acid. These facts must be kept in mind in the analysis 
 of Group II. Try to dissolve lead chloride in ammonia. Devise a 
 scheme for the separation of Ag, Hg (mercurous) and Pb. 
 
 132. To a solution of lead acetate add a solution of sodium carbon- 
 ate till alkaline. The precipitate is a basic carbonate similar to "white 
 lead." Shake a solution of lead acetate with PbO, filter and pass 
 through carbon dioxide which will give much the same compound. 
 Mix 5 c.c. con. nitric acid and 5 c.c. water, heat and add a little at 
 a time about 2 grams red lead, PbaO^ The brown product is lead di- 
 oxide, PbO 2 . Filter a few drops of the liquid and add dil. H 2 S0 4 . Is 
 there lead in solution? To show the oxidizing power of lead dioxide 
 add to it and the rest of the liquid a few drops of any salt of manga- 
 nese and boil persistently. Let the solid matter settle and note deep 
 red color due to permanganic acid, HMn0 4 . 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 63 
 
 133. To 20 c.c. of the lead acetate solution of the laboratory add 
 80 c.c. pure water, place in flask and with a thread suspend in it a 
 folded strip of zinc, about 5 grams, and let stand till next period. Ex- 
 amine the "lead tree," and test the solution for lead and zinc. Which 
 metal is most electropositive? 
 
 ALUMINIUM. 
 
 134. Dissolve a few tenths of a gram of Al in con. HC1 and dilute 
 to a test tube full, and use where A1C1 3 is required (a). Dissolve an- 
 other small amount of Al in a few c.c. of NaOH warming, which gives 
 sodium aluminate, Al (ONa) 3 , (b). What gas is given off in each case? 
 
 To a little of (a) add ammonia in excess. To (a) add drop by drop 
 NaOH till a permanent precipitate is obtained. Dissolve one-half of it 
 by adding HC1 and the other half with an excess of NaOH. Is A1(OH) 8 
 amphoteric? In what two ways may it ionize? To (b) add ammonium 
 chloride till the precipitate is permanent, and account for its formation. 
 How could you separate Al and Zn? 
 
 To portions of (a) add in excess sodium carbonate, and ammonium 
 sulfide and account for the precipitation of Al (OH) 8 in each case? 
 
 135. By heating dissolve 15 grams of aluminium sulfate, 
 A1 2 (SOOs 18H 2 O in 50 cc. water, stir in the calculated weight of am- 
 monium sulfate to make ammonium aluminium alum (NH^aAMSOO* 
 and when all is dissolved filter into a crystallizing dish while hot. Let 
 a drop or two of nitrate fall on a watch glass, let the water evaporate 
 and examine the crystals with a microscope. Examine the crystals in 
 dish, dissolve a few and test the solution with blue litmus paper. Why 
 are the solutions of Al salts acid? See 63. 
 
 CHKOMIUM. 
 
 136. Chromate Ion to Chromium Ion: Dissolve in a dish 10 
 grams potassium dichromate in 50 c.c. water with heat and let cool. 
 Add 10 c.c. con. sulfuric acid and then a little at a time 50 per cent 
 alcohol heating if necessary to start the reaction. Continue the addi- 
 tion of alcohol till the solution has a green color, a few drops in much 
 water showing no brown. Avoid a large excess of alcohol. Set aside 
 two-thirds of the solution and examine the crystals of chrome alum 
 at the next laboratory period. Dilute the other third with 20 times its 
 volume of water and use it below. One molecule of the dichromate is 
 reduced by the oxidation of three molecules of alcohol to aldehyde, 
 CH 3 COH. With aid of text write the equation. What was the action 
 of H 2 S and of SO 2 on acidified solutions of K 2 Cr 2 O 7 ? 
 
 Treat a portion of your solution of chrome alum with an excess of 
 ammonia. How separate Cr and Zn? Try to dissolve a little of the 
 remaining precipitate with NH 4 C1. How separate Cr from Mg? To an- 
 
64 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 other portion of chrome alum solution add an excess of NaOH and boil. 
 How separate Cr and Al? Is Cr(OH) 8 amphoteric? Try portions of 
 your chrome alum with solutions of sodium carbonate and ammonium 
 sulfide. In each case the hydroxide is precipitated as in the case of Al. 
 
 137. Chromium Ion to Chromate Ion: To a few c.c. of chromium 
 nitrate or chloride solution in a dish add an equal volume of NaOH. 
 Stir into it gradually 2 grams sodium dioxide and heat to boiling. Fil- 
 ter if necessary, acidify a little of the filtrate with acetic acid being 
 sure of an excess by testing, and add barium chloride. To a little di- 
 lute solution of sodium dichromate add acetic acid and barium chloride 
 and compare precipitates. 
 
 13^. Dichromate to Chromate: (a) Dissolve 5 grams sodium di- 
 chromate in 25 c.c. water in a dish, add slowly with stirring NaOH till 
 the solution becomes yellow. Evaporate till the salt crystallizes on 
 cjoling. 
 
 (fo) Chromate to Dichromate: Dissolve 10 grams sodium chro- 
 mate in 25 cc. water and add the calculated weight of con. sulfuric 
 acid, that is, one molecule of acid to two of the chromate. Evaporate 
 till the dichromate crystallizes out on cooling. 
 
 (c) Chromium Trioxide: Dissolve 2 grams potassium dichromate 
 ii 5 c.c. water by heat, cool till it begins to separate then drop con. 
 dlfuiic acid directly upon the surface of the liquid in test tube till the 
 precipitate formed does not quite dissolve. Heat to dissolve most of 
 it, note the escape of some oxygen, set in rack to cool and observe 
 crystals of CrO 3 in an hour. 
 
 139. Chromium Oxychloride: In a retort place 3 grams dichro- 
 mate, 2 grams NaCl and 15 cc. con. H 2 SO 4 . Heat and collect the oxy- 
 chloride as you did nitric acid. What does it look like? Dissolve some 
 of it in water, add an excess of NaOH, then an excess of acetic acid and 
 finally barium chloride. What is the precipitate? 
 
 140. Test for Chromate Ion : To two small portions of a chromate 
 solution add acetic acid then barium chloride and lead acetate respec- 
 tively. To a neutral solution of a chromate add -silver nitrate. Divide 
 into two portions. In one try the solubility of the silver chromate in 
 ammonia and in the other with dil. nitric acid. A third test is the re- 
 duction of chromate ion to chromium ion with change from yellow or 
 red to green as in 136. 
 
 IROff, NICKEL, COBALT. 
 Iron. 
 
 141. Dissolve most of about 1 gram of card teeth in con. HC1 di- 
 luted with an equal volume of water. Pour one half in a dish. Dilute 
 the other half to a test tube full leaving in it the undissolved iron. This 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 65 
 
 is ferrous chloride, FeCl 2 , solution (a). Heat the half in the dish and 
 with stirring add a little con. nitric acid at a time till the black pre- 
 cipitate at first formed dissolves and a drop of the solution in a little 
 water gives no blue color with potassium ferricyanide. Dilute to a 
 test tube full. It is ferric chloride, FeCls (b). 
 
 Treat small portions of (a) with an excess of ammonia, NaOH, so- 
 dium carbonate, H 2 S, ammonium sulfide. The first and second give 
 ferrous hydroxide Fe (OH) 2 , changing in air to ferric hydroxide Fe (OH) 3 ; 
 the third gives ferrous carbonate. Hydrogen sulfide has no effect, but 
 ammonium sulfide gives black FeS. Treat small portions of (b) with 
 the same reagents in excess. Ammonia, NaOH and sodium carbonate 
 give ferric hydroxide. H 2 S reduces ferric to ferrous iron with the sep- 
 aration of sulfur. Why does it not precipitate iron sulfide? Is either 
 hydroxide amphoteric? Try to dissolve ferrous and ferric hydroxides 
 with ammonium chloride. Compare the action of sodium carbonate 
 and ammonium sulfide on ferric iron, aluminium and chromium solu- 
 tions. Devise ways to separate iron from Al, Cr, Cu, Mg, As. 
 
 142. Treat portions of (a) and (b) with solutions of potassium 
 ferrocyanide, potassium ferricyanide, ammonium sulfocyanide and 
 tabulate results as tests for ferrous and ferric iron. 
 
 The following is a fine example of a reversible reaction and illus- 
 tration of the influence of the common ion: To a test tube nearly full 
 of water add about 5 drops of ammonium sulfocyanide and the same 
 amount of ferric chloride. The reaction is, 
 
 FeCl 3 +3NH 4 CNS= (reversibly) Fe (CNS) 3 +3NH 4 Cl. 
 
 Divide the red solution in four test tubes. To one add more of the 
 ferric chloride, to the second more of the sulfocyanide, to the third 
 ammonium chloride, and compare colors with that of the fourth. Re- 
 fer to 62 and to text book and explain fully. 
 
 143. Double Salts: (a) Dissolve in 50 cc. water by' heat 10 
 grams ferrous sulfate and the calculated amount of ammonium sulfate 
 to make ammonium ferrous sulfate and filter hot into a crystallizing 
 dish. Examine crystals in dish, also let a drop cool on watch glass 
 and use microscope. This is Mohr's salt, (NH 4 ) 2 Fe(SO 4 ) 2 6H 2 O. 
 
 (b) To make iron alum dissolve 15 grams ferrous suifate in 25 c.c. 
 water in dish, also the calculated amounts of ammonium sulfate and 
 con. sulfuric acid. Now heat and add slowly with stirring, con. nitric 
 acid till a drop diluted shows no ferrous iron. Set aside to crystallize, 
 giving (NH 4 ) 2 Fe 2 (S0 4 ) 4 24H 2 0. 
 
 JICKEL AND COBALT. 
 
 144. From nickel and cobalt chlorides or nitrates precipitate their 
 hydroxides with NaOH in excess. Try to dissolve portions of the hy- 
 droxides in excess of NaOH with heating. Are they amphoteric? Try 
 
66 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 to dissolve other portions of the hydroxides with ammonium chloride. 
 How may these metals be separated from ferric iron, aluminium, chro- 
 mium? 
 
 Try to precipitate the sulfides of Ni and Co with H 2 S, then with 
 H 2 S and ammonia or with ammonium sulfide. Try to dissolve the sul- 
 fides with dil. HC1. How may Ni and Co be separated from Zn, Fe, Sb 
 and the metals of groups I and II? 
 
 To 1 to 2 c.c. of solutions of Ni and Co add NaOH with shaking till 
 the hydroxides are just permanent. Add to each tube acetic acid in 
 slight excess, then to each 10 cc. of a solution of potassium nitrite and 
 let stand. The precipitate is K 3 Co(NO 2 )e. How may Co be separated 
 from Ni? 
 
 Again precipitate the hydroxides of the two metals from 1 to 2 cc. 
 of their solutions avoiding a large excess of NaOH. Add to each a so- 
 lution of KCN (dangerous) till the precipitates just dissolve. Now add 
 to each about 1 c.c. of NaOH and bromine water till it colors them per- 
 manently red. A black precipitate of nickelic hydroxide should be ob- 
 tained. 
 
 MANGANESE 
 
 145. To a few cc. of a solution of manganous salt, as MnCl 2 , add 
 ammonia in excess and to another portion NaOH in excess. Does the 
 Mn(OH) 2 redissolve? How may Mn be separated from Zn, Al, Cu, Ag? 
 
 Add to one fresh portion of the solution ammonia in excess then 
 ammonium chloride, and to another portion add ammonium chloride 
 then ammonia. Compare the results with those obtained with Mg and 
 the same reagents. To one of these solutions add ammonium sulfide 
 and to the other hydrogen sulfide. How could you separate Mn and 
 Mg? Try to dissolve the sulfide in dil. HC1. How may Mn be separ- 
 ated from Cu, Hg, As, Sb? 
 
 To a manganese solution add ammonium chloride, ammonia, and 
 sodium phosphate. Compare result with the action of these reagents 
 on solutions of Zn and Mg. 
 
 146. Permanganic Acid and Permanganate: To about 2 c.c. of 
 a solution of any manganous salt add an equal volume of con. nitric 
 acid, then about 1 gram of red lead or lead dioxide, and heat some time 
 at the boiling point. Let the undissolved matter settle and note red 
 color of permanganic acid. This is a good test for Mn. Pour off a 
 little of the clear solution into water to see color better. 
 
 (b) Melt in an iron crucible 5 grams solid KOH and 2.5 grams 
 potassium chlorate, and stir in gradually 2 grams MnO 2 . Heat with 
 stirring till the mass turns solid and raise the temperature with full 
 burner flame, and continue 5 minutes. When the mass is cold dis- 
 solve out by heating with the crucible nearly full of water. Pour the 
 solution into a large test tube and let settle. To one portion of the 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 67 
 
 green solution of potassium manganate, K 2 MnO4, add dilute sulfuric 
 acid till* it just turns red forming potassium permanganate, KMnO<. 
 Try changing the manganate to permanganate in another portion by 
 pa-rsing through it carbon dioxide, and in a third by diluting it with 
 much water. 
 
 147. Dissolve about one-fourth gram of oxalic acid in water, add 
 5 c.c. dil. sulfuric acid, heat to about 80 degrees and add a little at a 
 time a solution of potassium permanganate till the color becomes per- 
 manent. Dissolve about one-half of a gram of ammonium ferrous sul- 
 fate (143a), add sulfuric acid and permanganate as above. These re- 
 actions illustrate oxidation by permanganic acid and are much used 
 in quantitative analysis. Write the equations, assuming that the oxalic 
 acid is oxidized to water and CO- and the FeSO 4 to Fe-CSO^s. 
 QUALITATIVE SEPARATION OF THE METALS 
 
 The following scheme of analysis is prepared for first year stu- 
 dents in chemistry to be used in the separation of the common metals. 
 Provision is not made for every contingency. For example, it is as- 
 sumed that the ions to form insoluble phosphates in Group III are 
 not present. 
 
 148. Group I; Ag, Pb, Hg: Determine with test paper whether 
 the solution is neutral or only slightly acid. If neutral add one-tenth 
 of its volume of HC1, sp. gr. 1.12. If strongly acid, neutralize with 
 ammonium hydroxide and then add the HC1. The purpose is to have 
 enough acid to prevent the precipitation of BiOCl in this group and 
 ZnS in the next, and not enough to prevent the precipitation of SnS 
 and CdS in group II. 
 
 If no precipitate is formed pass to group II. If one is formed let 
 remain a few moments and filter. Set aside the filtrate (1) for group 
 (II). Wash twice the precipitated chlorides of Ag, Pb, Hg 2 with small 
 portions of cold water. Now pour through the filter a half test tube 
 full of boiling water. Boil and pour through again. Add to one-half 
 of filtrate dil. sulfuric acid and let stand. A white precipitate shows 
 lead sulfate. To the other add potassium chromate. A yellow preci- 
 pitate is lead chromate. - 
 
 Treat the remaining precipitate on the filter with about 5 c.c. of 
 NH\OH. Pass it through a second time. Now add to this filtrate an 
 excess of dil. nitric acid, making sure of an excess by mixing and 
 testing. AgCl is obtained if silver is present. The blackening of the 
 residue left on filter when ammonia was added shows Hg 2 present, the 
 black substance being Hg and HgNH 2 Cl. 
 
 Write equations for all reactions that occur in the analysis of 
 Group I. 
 
 149. Group II; As, Sb, Sn, Hg, Pb, Bi, Cu, Cd: Heat filtrate (1) from 
 group (I) nearly to boiling and pass in H 2 S for about 10 minutes keep- 
 
68 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 ing the temperature near the boiling point so as to precipitate arsenic 
 from arsenates if present. Now let cool, add an equal volume of wa- 
 ter and pass in the gas for several minutes to insure the complete 
 precipitation of cadmium and tin. Even then it is well to let the pre- 
 cipitate settle, pour off a few drops of the liquid, add water and more 
 gas. Disregard a white precipitate which might be ZnS. When as- 
 sured that the precipitation is complete let the precipitate settle and 
 pour off the solution through a filter. At once boil it to expel H 2 S and 
 set it aside as filtrate (2) for group III. 
 
 Wash the precipitate four times by decantation; that is, pour upon 
 it hot water, boil, let settle completely and pour off the wash water. 
 Drain carefully the last time since much water present will dilute the 
 nitric acid to be used. 
 
 If it is known that As, Sb and Sn are absent omit the bracketed di- 
 rections, and boil with dil. nitric acid as directed after the bracketed 
 lines. If not known whether they are present or not proceed as di- 
 rected within the brackets. 
 
 [Now heat a small portion of the washed precipitate with about 3 
 c.c. of yellow ammonium sulfide, best in a dish on a water bath with 
 stirring. In any case do not boil. Filter off the solution and add to fil- 
 trate dil. HC1 till acid throughout, and boil. If only white finely di- 
 vided sulfur is precipitated As, Sb and Sn are absent. In this case pro- 
 ceed to boil the remainder of the precipitate with nitric acid as after 
 the brackets. If, however, a flocculent yellowish precipitate is ob- 
 tained one or more of the above metals are present. Treat the re- 
 mainder of the washed precipitate with ammonium sulfide as directed 
 for the small portion, filter and set aside to be examined for As, Sb 
 and Sn, as described within the brackets below. 
 
 The residue consisting of sulfides undissolved by ammonium sul- 
 fide must be well washed on the filter. Punch through the tip of the 
 filter, wash the sulfides into a test tube and let settle.] Pour off as 
 much as possible of the water, add about 10 c.c. dil. nitric acid and 
 boil persistently. All the sulfides except that of Hg are dissolved. A 
 floating residue of S often remains and is to be disregarded. A heavy 
 residue which settles at once to the bottom must be tested for Hg. Fil- 
 ter off the HgS, add to the nitrate 10 cc. dil. sulfuric acid and begin its 
 evaporation in a porcelain dish. To confirm the presence of Hg dis- 
 solve the sulfide on the filter by pouring upon it about 3 c.c. of hot 
 con. HC1 to which a little potassium chlorate has been added. Heat and 
 pass through again if necessary. Add to the solution bromine water 
 till the color of Br persists, boil out the excess and add a few drops of 
 a solution of stannous chloride. A white precipitate turning to gray 
 and perhaps to black shows Hg present. 
 
 Evaporate the nitric acid solution containing the other metals till 
 
LABORATORY MANUAL OF GENERAL CHEMISTRY 69 
 
 heavy white fumes of sulfuric acid freely come off. The evaporation 
 must be carried far enough to expel all the nitric acid. (See 131) Let 
 the dish and contents cool, then add 10 c.c. of water and let stand 5 
 minutes. A white precipitate shows lead sulfate. Filter it off and add 
 to filtrate an excess of ammonium hydroxide, making sure of an ex- 
 cess by shaking to mix and then testing with turmeric paper. A white 
 precipitate shows Bi but its presence should be confirmed thus: Fil- 
 ter it off. Dissolve on the filter in a few drops of con. HC1, evaporate 
 off nearly all of the acid. (Why?) Pour into much water. A white 
 cloud is BiOCl. 
 
 The filtrate from the bismuth hydroxide is blue if copper is pres- 
 ent. If copper is absent add hydrogen sulfide to the filtrate which will 
 give yellow CdS if Cd is present. If the blue color shows Cu present 
 test for Cd by one of the following: 
 
 (a) Add to the blue solution a solution of potassium cyanide (very 
 poisonous) till the blue color disappears, then hydrogen sulfide which 
 will give yellow CdS if Cd is present. 
 
 (b) Add to the blue solution dil. sulfuric acid till colorless then 
 card teeth or steel wool and boil for some time. The iron removes the 
 Cu. Why? Filter, be sure that the filtrate is still acid. If not add a 
 slight excess of sulfuric acid and then H 2 S which will give yellow CdS 
 if Cd is present. 
 
 [To the (NH 4 ) 2 S solution containing As, Sb and Sn, add dilute HC1 in 
 excess, which re-precipitates the sulphides if present. A precipitate 
 is always produced owing to the separation of sulphur. If the sul- 
 phides are present, however, the precipitate is more highly colored and 
 somewhat flocculent. Filter and wash the sulphides, carefully remove 
 to a test tube and heat with strong HC1. Sb 2 S 3 and SnS 2 are dissolved, 
 while AsaSs is not. Dissolve the As 2 S 3 in hot, strong HC1 and KCIO*, 
 boiling if necessary, best in a small flask placed in the evaporating 
 closet. Add to the solution NEUOH in excess then NH 4 C1, filter if not 
 clear and add MgSO 4 , which will give on standing a precipitate of 
 NH 4 MgAsO 4 if As is present. 
 
 Boil persistently the filtrate from the arsenic sulfide with card 
 teeth. The Sb is deposited on the card teeth from which it can be de- 
 tached in scales by stirring with a glass rod, and the tin remains in so- 
 lution as SnCl 2 . Filter off the antimony and test the filtrate for tin 
 with HgCl 2 . Wash the antimony thoroughly, dissolve on the filter with 
 a very little aqua regia, dilute with 10 to 20 times its volume of water 
 and add H 2 S which will give orange yellow Sb 2 S 3 L 
 
 Write equations for all reactions involved in Group II assuming 
 that all metals are present in the beginning as chloride. 
 
 150 Group III: ..Fe, Al, Cr, Co, Hi, Mn, Zn: To a small portion 
 of filtrate 2 from Group II add ammonia in excess as proved by shak- 
 
70 LABORATORY MANUAL OF GENERAL CHEMISTRY 
 
 ing and testing. If a precipitate is formed one or more of the metals, 
 Fe, Al, Cr are present. If no precipitate is formed add H 2 S; or, if a 
 precipitate was formed by ammonia filter it off and add H 2 S to the fil- 
 trate. If H 2 S causes a precipitate which is not black Co, and Ni are 
 absent. A black precipitate means either Co or Ni or both and they 
 may mask Mn and Zn. By making these preliminary tests and atten- 
 tion to the following much work and time may be saved. 
 
 If no precipitate was formed by either ammonia or H 2 S proceed 
 with the rest of the filtrate to Group IV. If a precipitate was formed 
 only with ammonia there is no need to test for Co, Ni, Mn, Zn. In this 
 case add ammonia to all the filtrate, filter off the hydroxides of Fe, 
 Al, Cr, and save the filtrate for Group IV, omitting the addition of H 2 S 
 below and also that part of the directions applying to Co, Ni, Mn, Zn. 
 If H 2 S made a precipitate while ammonia did not Al and Cr are ab- 
 sent but Fe may be present if the precipitate was black. 
 
 The following scheme provides of course for the presence of all 
 these metals: 
 
 To filtrate (2) from Group II add ammonia in slight excess test- 
 ing after shaking, then add about 2 c.c. more. Heat the solution in a 
 flask or large test tube nearly to boiling and pass in H 2 S. Shake and 
 heat frequently to make the precipitate granular and more easily filt- 
 ered and washed. When the precipitation is apparently completed fil- 
 ter a small portion and pass into it more gas. If more precipitate forms 
 add more H 2 S to the whole and thus proceed till complete precipita- 
 tion is attained. 
 
 Filter through a fluted filter and wash with hot water. Boil the 
 filtrate (3) till all hydrogen sulfide is expelled, filter it and set aside 
 for Group IV. As soon as the precipitate produced by H 2 S is suffi- 
 ciently washed, (about two funnels full of hot water), punch through 
 the tip of the filter and wash most of the precipitate into a small 
 flask with a fine stream of water from the wash bottle. At once add 
 considerable excess of dil. HC1. Only NiS and CoS remain undis- 
 solved. Filter, reserve filtrate for B and test for Ni and Co in A. 
 
 A. Dissolve the NiS and CoS in the filter in about 3 c.c. of aqua 
 regia and evaporate the solution nearly to dryness in a dish. Add about 
 4 c.c. water then NaOH drop by drop with shaking till a permanent 
 precipitate is obtained. Avoid a large excess of NaOH. Divide the li- 
 quid and suspended precipitate into two portions, and test for Co and 
 Ni as follows: 
 
 1. To one portion add about 2 cc. acetic acid, 30 per cent, and 15 
 c.c. of a 25 per cent solution of potassium nitrite and let the solution 
 stand half an hour. A yellowish, white granular precipitate is 
 
LABORATORY MANUAL OP GENERAL CHEMISTRY 71 
 
 2. To the other half add a solution of potassium cyanide (dan- 
 gerous) drop by drop till the precipitate just dissolves. It is essen- 
 tial to avoid a large excess. Warm the solution, add about 5c.c. NaOH 
 and then bromine water till the color of Br persists. A black precipi- 
 tate often forming after some time is Ni(OH) 8 . 
 
 B. Evaporate the filtrate from the CoS and MS nearly to dryness, 
 add about 5 c.c. water, make strongly alkaline with NaOH and add 
 gradually with stirring about 2 grams of sodium dioxide. Boil a few 
 minutes, add water and filter off the Fe and Mn hydroxides. Save the 
 nitrate for C and wash the precipitate. To test for Mn make a 
 bead of sodium carbonate, while hot touch to powdered KC1O 3 . Let 
 cool, take up with it a little of the precipitate and fuse with the blow- 
 pipe. A green bead shows Mn present. To confirm the presence of 
 Mn and test for Fe remove a small portion of the precipitate from the 
 filter, dissolve it in dilute nitric acid and add a solution of potassium 
 ferrocyanide which will give a dark blue color if Fe is present. Now 
 pour upon the remaining precipitate about 3 c.c. of hydrogen dioxide 
 and then add 10 c.c. nitric acid, sp. gr. 1.2, which means practically 1 
 vol. of pure concentrated acid diluted with 2 parts of water. Warm 
 and pass through again if necessary. To this solution add a little at a 
 time 2-4 grams of lead dioxide or red lead and heat to gentle boiling 
 for a few moments. Let the suspended matter settle when the liquid 
 above it will be colored red by permanganic acid if Mn is present. 
 
 C. To the nitrate from the Fe and Mn hydroxides which may con- 
 tain Al, Cr and Zn, add con. nitric acid in slight excess testing with 
 litmus paper. Add 5 c.c. ammonium chloride, heat and add ammonia 
 in slight excess as shown by test after stirring. A white flocculent 
 precipitate shows aluminium hydroxide. Filter it off and acidify the 
 filtrate with acetic acid, mixing and testing with litmus paper. The 
 liquid is yellow if chromate radical is present. Add barium chloride 
 which will give yellow BaCrO4. The BaCl 2 must be added in excess to 
 remove all the CrO4 which would interfere with the test for Zn. Filter 
 and pass through several times if necessary. The clear filtrate should 
 not be yellow. If it is add more BaCl 2 and filter again. To the clear, 
 colorless filtrate add H 2 S which will give an evident, white precipi- 
 tate if Zn is present. Hydrogen sulfide always produces an opales- 
 cence at this point and this is to be disregarded. 
 
 GROUP IT. 
 
 151- If the analysis is being carried through all the groups the 
 solution used will be nitrate (3), in which case it should be made 
 slightly acid with dilute HC1, evaporated to two-thirds of its volume, 
 and ammonium hydroxide added till alkaline then ammonium carbon- 
 ate. If the solution is an original one, add ammonium chloride, am- 
 
72 LABORATORY MANUAL OP GENERAL CHEMISTRY 
 
 monium hydroxide till alkaline then ammonium carbonate till the pre- 
 cipitation is complete. To determine this heat the solution nearly to 
 boiling, let settle and add a few drops of the carbonate to the clear so- 
 lution. Filter and to a small portion of the filtrate add sodium phos- 
 phate which will give (NH 4 )MgPO 4 if magnesium is present. Reserve 
 the remainder of the filtrate (4) for Group V. Wash the carbonates 
 on the filter and dissolve by pouring upon the filter not more than 5 
 c.c. of acetic acid, letting it run through into a clean tube. Run through 
 several times if necessary to dissolve all the carbonates. To a small 
 portion of the acetic acid solution add twice its volume of a saturated 
 solution of calcium sulfate. Note carefully whether the precipitate is 
 formed at once or only after a few seconds. Proceed according to one 
 of the following as required, using the rest of the acetic acid solution: 
 
 (a) If no precipitate was formed, to the remainder of the acetic 
 acid solution add ammonia in excess and ammonium oxalate which will 
 precipitate calcium oxalate. 
 
 (b) If the precipitate was slowly formed, only Sr and Ca can be 
 present. To the acetic acid solution add a solution of ammonium sul- 
 fate and let stand. Filter off the strontium sulfate, and to the fil- 
 trate add ammonium hydroxide till alkaline then ammonium oxalate 
 which will precipitate CaC 2 O4, if Ca is present. 
 
 (c) If the precipitate with CaSO 4 was immediate Ba is present, 
 and the other two may be. To the acetic acid solution add an excess 
 of pure potassium chromate solution. Filter off the barium chromate. 
 Make the filtrate alkaline with ammonia and add ammonium carbon- 
 ate, which will precipitate Ca and Sr carbonates if they are present. 
 Filter off the carbonates, wash till most of the excess of chromate is 
 removed, and dissolve on the filter with acetic acid. Test a small por- 
 tion of the filtrate with CaS0 4 . If no precipitate forms on long stand- 
 ing Sr is absent and Ca should be tested for as in (a). If Sr is pres- 
 ent remove it from the remainder of the acetic acid solution with am- 
 monium sulfate, let stand 5 minutes, filter and test filtrate for Ca as 
 in (b). 
 
 GROUP V. 
 
 152. Evaporate a part of filtrate (4) to dryness in a porcelain 
 crucible; heat till all ammonium salts are expelled, and test with the 
 spectroscope as directed in lecture for Na, K and Li. If uncertain as 
 to the spectrum lines and flame colors, compare with spectra of the 
 known substances. 
 
 To test for NH 4 , place in a porcelain crucible a small amount of 
 lime, add enough of the original solution to moisten it, cover with a 
 watch glass with a strip of moist turmeric paper on the under side, 
 and warm gently. If NH 4 be present, NH 8 will be evolved and can be 
 recognized by its action on the turmeric paper and by its odor. 
 
APPENDIX 
 
 153. Atomic Weights of the Common Elements: 
 
 Aluminum 
 
 Al 
 
 27.1 
 
 Lead 
 
 Pb 
 
 207.20 
 
 Antimony 
 
 Sb 
 
 120.2 
 
 Lithium 
 
 Li 
 
 6.94 
 
 Arsenic 
 
 As 
 
 74.96 
 
 Magnesium 
 
 Mg 
 
 24.32 
 
 Barium 
 
 Ba 
 
 137.37 
 
 Manganese 
 
 Mn 
 
 54.93 
 
 Bismuth 
 
 Bi 
 
 208.0 
 
 Mercury 
 
 ffff 
 
 200.6 
 
 Boron 
 
 B 
 
 11.0 
 
 Molybdenum 
 
 Mo 
 
 96.0 
 
 Bromine 
 
 ... Br 
 
 79.92 
 
 Mckel 
 
 IVi 
 
 58.68 
 
 Cadmium.. . 
 
 ....Cd 
 
 112.40 
 
 Nitrogen 
 
 IV 
 
 14.01 
 
 Calcium 
 
 ...Ca 
 
 40.07 
 
 Oxygen 
 
 
 
 16.00 
 
 Carbon 
 
 ...C 
 
 12.00 
 
 Phosphorus 
 
 ....p 
 
 31.04 
 
 Chlorine 
 
 ....Cl 
 
 35.46 
 
 Platinum 
 
 Pt 
 
 195.2 
 
 Chromium 
 
 .... Cr 
 
 52.0 
 
 Potassium 
 
 K 
 
 39.10 
 
 Cobalt 
 
 ...Co 
 
 58.97 
 
 Silicon.... 
 
 Si 
 
 28.3 
 
 Copper 
 
 Cu 
 
 63.57 
 
 Silver 
 
 Ag 
 
 107.88 
 
 Fluorine.. . 
 
 F 
 
 19.0 
 
 Sodium.. 
 
 ...Na 
 
 23.00 
 
 Gold 
 
 Au 
 
 197.2 
 
 Strontium 
 
 Sr 
 
 87.62 
 
 Hydrogen.. 
 
 ..H 
 
 1.008 
 
 Sulphur 
 
 s 
 
 32.06 
 
 Iodine . . 
 
 I 
 
 126.92 
 
 Tin .. . . 
 
 Sn 
 
 118.7 
 
 Iron. . 
 
 Fe 
 
 55.84 
 
 Zinc 
 
 ...Zn 
 
 65.37 
 
 
 
 
 
 
 
 154. Vapor Pressure of Water in Millimeters of Mercury: 
 
 Degrees C. 
 
 Pressure 
 
 Degrees C. 
 
 Pressure 
 
 15 
 
 12.7 
 
 26 
 
 25.0 
 
 16 
 
 13.6 
 
 27 
 
 26.5 
 
 17 
 
 14.5 
 
 28 
 
 28.1 
 
 18 
 
 15.4 
 
 29 
 
 29.8 
 
 19 
 
 16.4 
 
 30 
 
 31.6 
 
 20 
 
 17.4 
 
 31 
 
 33.4 
 
 21 
 
 18.5 
 
 32 
 
 35.4 
 
 22 
 
 20.0 
 
 33 
 
 37.4 
 
 23 
 
 20.9 
 
 34 
 
 39.6 
 
 24 
 
 22.2 
 
 35 
 
 41.9 
 
 25 
 
 23.5 
 
 
 
iTendrixon 
 
 387354 
 
 Lab orate 
 
 ry manual of 
 
 H4 
 
 general che 
 
 387354 
 
 UNIVERSITY OF CALIFORNIA LIBRARY