Laboratory Directions 
 
 IN 
 
 EDITED BY 
 
 WILLIAM C. BRAY, PH. D. 
 ASSISTANT PROFESSOR OF CHEMISTRY 
 
 IN THE 
 
 UNIVERSITY OF CALIFORNIA 
 1915 
 
 Copyright applied for by William C. Bray. 
 
 Lederer, Street & Zeus Co., Printers 
 
c; 
 
CHEMISTRY lA. UNIVERSITY, 0? ' EA-LjIOIOfiA' ::.-' 
 
 LABORATORY DIRECTIONS, 1915 
 
 Provide yourself with a notebook about 8 inches by 5 inches in size and 
 not loose-leaved. 
 
 Notes about recording the results of experiments are given in the first 
 three assignments and these notes apply to all later assignments. 
 
 Bring your lecture notes and textbook of inorganic chemistry into the 
 laboratory and use them for reference. 
 
 Satisfactory progress can be made in this course only when the student 
 accepts the responsibility of mastering, before the next period, the material 
 presented in each lecture and laboratory period. Consult your instructor in the 
 laboratory or at the quiz section about your difficulties. Frequent brief tests 
 will be given, usually at the beginning of the period. 
 
 At the first meeting of your laboratory section go to the room to which 
 you have been asigned (consult posted lists or inquire at office if necessary), 
 apply to your instructor for your key, check your apparatus by means of the 
 list in your locker, sign your name (surname first) in the proper place, and 
 return this list to your instructor. Begin the first assignment. 
 
 If, while checking your apparatus during the first laboratory period, any 
 article is found to be damaged, exchange it at the office, or sign a blue "return 
 slip" in order that you may receive proper credit if the damaged article is 
 broken during the term. 
 
 Fill your wash-bottle with distilled water. Sterilize the mouth-piece by 
 placing it in boiling water for a minute or two. Always use the distilled water 
 in your wash-bottle for the final rinsing of apparatus. 
 
 Rinse out your five reagent bottles, partially fill them with the 6 normal 
 laboratory solutions. Note that, while pouring the reagent from one of these 
 bottles, the stopper of the bottle should be held with the same hand that holds 
 the test tube or beaker. 
 
 In general, avoid contamination of any of the reagents in the laboratory 
 by keeping each stopper clean and returning it at once to the proper bottle. 
 
 Articles to replace broken apparatus, and additional articles for regular 
 use, are obtained at the office in the corridor (on the same floor as your room) 
 by filling out a white order slip and signing your name and desk number. 
 
 A yellow "temporary order slip" is used instead of a white order slip when 
 borrowing articles specified in the last paragraph of the list of apparatus 
 given below. These articles must be returned during the same laboratory period 
 before the office is closed. 
 
 Sign a blue "return slip" whenever any article is returned to the office. 
 The article will not be accepted unless it is clean and in good condition. 
 
 List of Apparatus. 
 
 I. Regular equipment of each locker. At the end of the term the locker 
 must contain the same amount of apparatus, no more and no less ; the locker 
 must be clean; the apparatus must be clean and dry, and in good condition; 
 glass stoppers must fit, and be protected by the insertion of a piece of paper. 
 
 1 Key. 
 
 .5 Beakers, 100 cc., 150 cc., 250 cc., 400 cc., 600 cc. 
 5 Reagent Bottles. 
 
 2 Sample Bottles, 50 cc. 
 
 i Graduated Cylinder, 50 cc. 
 
 4 Flasks, 500 cc., 250 cc., and two 125 cc. 
 
 i Wash-bottle, equipped with glass tubing and rubber stopper. 
 
 314^92 
 
2 Blue glasses. 
 
 2 Glass Rods, 12 cm. and 18 cm. 
 30 cm. Glass Tubing. 
 12 Test-tubes. 
 
 i Watch Glass. 
 
 i Casserole. 
 
 1 Crucible, with cover. 
 
 2 Evaporating Dishes, 6 cm. and 9 cm. 
 15 cm. Rubber Tubing, 3 mm. 
 
 i Bunsen Burner, with rubber tubing. 
 
 i Wire Gauze. 
 
 i Triangle. 
 
 i Test-tube Brush. 
 
 i Test-tube Holder. 
 
 i Test-tube Rack. 
 
 i Package Filter Paper. 
 
 1 Rule. 
 
 2 Towels. 
 
 II. Additional articles may be obtained at the office: 
 
 (a) By signing the regular white order slips. 
 Special apparatus for Assignments III and V, first term. 
 Matches. 
 
 Glass Flasks, 50 cc. 
 Corks. 
 
 Rubber Stoppers. 
 
 Platinum Wire and Platinum Foil. Partial credit will be given at the 
 end of each term for Platinum Wire and Foil returned to the office. 
 
 (b) By signing yellow temporary order slips. 
 Burettes, with clamps and pinchcocks. 
 Graduated Cylinders, 10 cc. and 250 cc. 
 
 Files. 
 Paraffin. 
 
 ASSIGNMENT I. 
 MEASUREMENT OF VOLUMES. 
 
 Notes: Enter the date and the title of the exercise at the top of a page 
 of the notebook. Use one side only of each page for recording experimental 
 results and answering questions. Make calculation^ neatly at the bottom or 
 side of the page, or on the opposite page. Show your notebook to your 
 instructor each day until he is satisfied with your method of making entries. 
 
 Each day, before leaving the laboratory; finish the experiments assigned 
 for the day's work, write out the answers to as many of the questions as pos- 
 sible, and clean the top of your desk. 
 
 Before the next laboratory period finish the questions and problems, 
 review your work and be sure that you understand it perfectly. Read over the 
 laboratory notes on the next laboratory exercise. 
 
 Experiment. Make several measurements of the inside diameter (in 
 centimeters) of your medium-sized beaker. (The instructor will suggest a 
 method.) Record each measurement in your notebook as soon as it is 
 made. The beakers are not perfectly cylindrical, nor is the method very 
 accurate ; the measurements, therefore, will usually show variations of a few 
 tenths of a centimeter. An average of at least four measurements will give the 
 mean diameter to about o.i cm. Questions: What fraction of the diameter 
 is this error? What percent of the diameter is this error? 
 
 Experiment. Fill the beaker with water to just below r the flare. Measure 
 
the depth of the water by dipping the rule into the beaker. Repeat the meas- 
 urement. Mark the level of the water by pasting a piece of gummed paper on 
 the outside of the beaker. Do not pour the water out of the beaker. 
 
 Calculate the area of the cross section of the beaker from your average 
 value of the diameter (assume pi = 3.14). From this result and the depth of 
 the water, calculate the volume of the water in the beaker. 
 
 Make further calculations to determine (i) what the change in the area 
 of the cross section of the beaker would have been (in sq. cm) if the diameter 
 had been o.i cm. greater than the value chosen; and (2) what the change in 
 the calculated volume (in cc.) would have been if the depth of the water had 
 also been o.i cm. greater. Questions: What percentage error in the area of 
 the cross section results from an error of o.i cm. in the diameter? What 
 percentage error in the volume results from errors of o.i cm. in the diameter 
 and o.i in the depth? How many decimal places should be used in expressing 
 the final result for the volume? 
 
 Experiment. Measure the volume of water in the beaker used in the 
 previous experiment by means of a measuring cylinder graduated in cc. 
 
 Compare this measured volume with the volume calculated. If the diff- 
 erence is greater than the estimated error, repeat the calculations and study 
 the following discussion to determine the reason. Repeat the measurements if 
 necessary. 
 
 Note. If you discover an error in your previous work, do not erase your 
 first entry, nor tear out the page. Mark the rejected result and insert the 
 correction near the original entry. 
 
 Note that the central surface of the water in a large beaker is level. Show 
 in a drawing the appearance of a vertical section of the surface of the water 
 (i) where it touches the glass in the large beaker, and (2) in a small glass 
 tube. How great is the difference (in cc.) between the readings at the top and 
 the bottom of the meniscus in your measuring cylinder? Volume readings for 
 water in glass vessels are made at the bottom of the meniscus. 
 
 Compare your notebook with the following sample page. Note whether 
 you have arranged your notes neatly and systematically. 
 
 Measurement of Volumes. Aug 1915 
 
 Used 3rd largest beaker. 
 
 Diameter, 6.3 cm. Depth of water, 7.2 cm. 
 
 6.4 " 7-3 " 
 
 6.3 " 7-2 ' 
 
 6.2 
 
 Average, 7.2 
 
 Average, 6.3 
 
 o.i is 0.1/6.3 = 0.016 of the diameter. This is % 
 
 of the diameter. 
 
 Area of cross section (dia. 6.3 cm.) = 31.2 sq cm. 
 Volume calculated (depth 7.2 cm.) = 225 cc. 
 Area calculated (dia. 6.4 cm.) = 32.2 sq. cm. 
 
 Difference, = i.o sq. cm. i. e %of the area. 
 
 Volume calculated (dia. 6.4 cm, = cc. 
 
 depth 7.3 cm.) 
 
 Difference, = cc., i. e % of the volume. 
 
 Volume measured with graduate = 219 cc. 
 
 Difference, = 6 cc. less. 
 This result .is (or is not?) within the experimental error. 
 
 Experiments. Measure by means of a measuring cylinder the volumes 
 
 5 
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of a second beaker, a test-tube, and a flask. As an exercise in judging vol- 
 umes, estimate the volume of another beaker and then measure it. Record in 
 a table the volumes of your flasks and beakers. What is the volume in liters 
 of your largest flask? 
 
 Experiments. Pour 10 cc. of water into a small test tube. Mark the height 
 of the water by pasting a small label on the side of the tube. Graduate this 
 test tube in the same way for 5 cc. and 2 cc., and set it aside for use in future 
 experiments. Graduate your smallest beaker for volumes of 50 cc. and 100 
 cc. in the same way. 
 
 The paper labels on the reagent bottles in the laboratory are covered with 
 paraffin. Why? Apply to the instructor to have the labels on your graduated 
 test tube and beaker paraffined in a similar manner. 
 
 State how you would use your graduated cylinder to measure out 75 
 grams of water. 
 
 Experiment, to give practice in weighing. Weigh to O.I g. the small beaker 
 which you have graduated. (See Notes on Weighing at the end of Assign- 
 ment). Fill the beaker to the 50 cc. mark, and weigh. What is the weight of 
 water in the beaker? Calculate its volume from this result. 
 
 Problems. (Always answer the problems before the next laboratory 
 period.) 
 
 1. How much water would you weigh out in order to have a volume of 
 10 cc. ? 
 
 2. How much mercury should be weighed out to give a volume of 10 
 cc.? The density of mercury is 13.6. 
 
 3. What is the weight of a liter of air? The density of air is 0.0012 at 
 room temperature and atmospheric pressure. 
 
 4. Mention three methods of determining the volume of water in a 
 cylindrical beaker. 
 
 Bring a table of atomic weights (e. g. Cady, page 96) to the laboratory 
 for use in Assignments II-V. 
 
 Note on Weighing. Follow carefully the directions given by the instruc- 
 tor. Report to him at once if the balance is out of order or if a weight is 
 missing. Enter the results of each weighing in your notebook while the 
 weights are on the scale-pan and check this result while you are returning the 
 weights to their places in the box. Summarize in a table the results of the 
 different weighings in each experiment. 
 
 ASSIGNMENT II. 
 
 THE SYNTHESIS OF COPPER SULFIDE. WEIGHT RELATIONS IN 
 CHEMICAL REACTIONS. 
 
 (Enter the date and title of this experiment on a new page.) 
 
 The purpose of this experiment is to determine the weights of copper and 
 
 sulfur which unite to form a sulfide of copper, and to calculate its formula 
 
 with the aid of the atomic weights of copper and sulfur. 
 
 Experiment. Support a clean porcelain crucible, with a cover, on a tri- 
 angle and heat with the colorless flame of a bunsen burner to low redness. Let 
 the crucible cool about 15 minutes, and weigh it, with the cover, to 10 milli- 
 grams. The weights and the balances are such that weighings cannot be 
 made closer than 10 mg. 
 
 While the crucible is cooling obtain a clean piece of copper wire, weigh- 
 ing about i gram, from the shelf and weigh it to 10 mg. Question: If the 
 weight of the copper is in error by 10 mg., what is the percentage error? 
 
 Place the copper in the weighted crucible and add enough powdered sul- 
 
fur to cover the copper. Place the cover on the crucible and heat gently (with 
 a small flame) until the sulfur ceases to burn at the edges of the cover, but 
 do not remove the cover while the crucible is hot. Then heat more strongly 
 until the bottom of the crucible just becomes dull red. Again allow to cool 
 about 15 minutes and weigh. 
 
 Carefully remove the cover and note the appearance of the contents of 
 the crucible, but do not touch the substance. If there is any free sulfur on the 
 cover or wall of the crucible, replace the cover, heat the crucible and cover, 
 and weigh again. In order to check the final weight, add a small quantity of 
 sulfur and repeat the experiment. (While waiting for the crucible to cool 
 answer the questions in the following paragraph.) At the end of the experi- 
 ment remove the .substance formed, break it, and note the differences between 
 its properties and those of copper and sulfur. To check the weight of the 
 empty crucible, clean the crucible by placing it in a porcelain dish containing 
 a little .nitric acid and warming the mixture, rinse the crucible with distilled 
 water, heat it, let it cool, andjweigh i^_again. By means of a table summarize 
 and compare the results obtained in both parts of the experiment. If there is 
 any discrepancy (beyond 10 mg.) suggest an explanation. 
 
 What does the increase in weight represent? Is any free sulfur left in the 
 crucible? (Give the reasons for your answers.) What would the increase 
 in weight have been if one gram-atom of copper had been used in the experi- 
 ment? How does this weight compare with the atomic weight of sulfur? 
 What, then, is the formula of the sulfide of copper which has been formed? 
 Write the reaction which has taken place. What is the molecular weight of 
 the substance formed? 
 
 Problems. I. (a) From your experimental results alone calculate the 
 percentage composition of the copper sulfide formed, (b) From the atomic 
 weights of copper and sulfur and the molecular weight of the sulfide calculate 
 the percentage composition of pure copper sulfide. Compare the results. 
 
 2. The formula of hydrogen sulfide is H,S. How much sulfur will com- 
 bine with 1.008 grams of hydrogen -to form this compound? What weight of 
 hydrogen sulfide will be formed? 
 
 3. What weight of iron will combine with 32.07 grams of sulfur to form 
 ferrous sulfide, FeS? Write the reaction. What is the percentage composition 
 of ferrous sulfide? How much ferrous sulfide would be needed to make 34-09 
 grams of hydrogen sulfide? 
 
 ASSIGNMENT III. 
 THE WEIGHT OF A LITER OF OXYGEN. 
 
 (Two students work together.) (Enter the date and title of this experi- 
 ment on a new page.) 
 
 When solid potassium chlorate is strongly heated it decomposes with 
 evolution of oxygen. The purpose of this experiment is to measure the volume 
 of a definite weight of oxygen (obtained from potassium chlorate). This may 
 be accomplished by heating potassium chlorate in a hard-glass test-tube, and 
 permitting the oxygen evolved to displace an amount of water equal to its own 
 volume. The loss in weight of the potassium chlorate gives the weight of the 
 oxygen evolved, and the volume of water displaced gives the volume of this 
 amount of oxygen. From these data the weight of i liter of oxygen can be 
 calculated. 
 
 The potassium chlorate decomposes readily, and at a much lower 
 temperature when a small quantity of manganese dioxide is present. The hard- 
 glass test-tube may then be replaced by a heavy-walled test-tube of ordinary 
 easily-fusible glass; but care must be taken not to heat the latter tube to a 
 higher temperature than is necessary for the reaction. The manganese 
 dioxide is a "catalyzer" in this reaction, and all of it may be recovered at 
 
 7 
 
the end of the experiment after the potassium chlorate has been decomposed 
 into potassium chloride, KC1, and oxygen. 
 
 Two students working together obtain from the office a heavy-walled glass 
 test-tube, a rubber stopper, glass tube, rubber tube, pinch cock and clamp, 
 One student signs a white slip for "special apparatus for Assignment III" , 
 but should make sure that the other shares the expense with him if any portion 
 of this apparatus is broken or damaged. The apparatus should be returned as soon 
 as the experiment is finished. Each student must keep a complete record of the 
 experiment in his notebook. 
 
 Experiment. Set up the apparatus according to the diagram on the black- 
 board and follow the directions given by the instructor. Plan a method of test- 
 ing if the apparatus is air-tight, and be sure that the pressure of the air inside 
 the tube at the beginning of the experiment is the same as the atmospheric 
 pressure outside. Do not begin to heat the tube containing the potassium chlo- 
 rate until the instructor has examined your apparatus. 
 
 Warm the heavy-walled test-tube carefully to dry it, holding its mouth 
 lower than its end so that a drop of water may not run down the side on to the 
 hot glass and crack the tube. Transfer at least i gram of powdered potassium 
 chlorate (half the quantity contained in the small tube) to a porcelain dish and 
 dry it by holding it high above a bunsen flame for several minutes ; be careful 
 not to heat the substance too hot. Set the chlorate aside to cool, weigh the test- 
 tube to o.i gram, transfer the potassium chlorate to the test-tube, and weigh to 
 o.oi g. Add a small quantity of manganese dioxide (not more than 50 mg.), mix 
 it with the chlorate by jarring the tube, remove any powder from the outside of 
 the tube and from the inside near its mouth, and weigh the tube carefully. 
 
 Begin the experiment by opening the pinchcock on the siphon tube, and 
 heating the tube containing the potassium chlorate very gently with a small 
 flame. Gradually heat the tube more strongly until the chlorate melts and gas 
 evolution begins. Then continue to heat the tube carefully and not too strongly 
 until enough chlorate has decomposed to force over 250 cc.. to 300 cc. of water 
 into the receiving beaker. Now discontinue the heating and allow the tube to 
 cool to room temperature. By raising the beaker, which holds the displaced 
 water, equalize the level of the water in the beaker and the flask, and then close 
 the pinchcock on the siphon tube. By means of a 250 cc. graduated cylinder 
 measure the amount of water which the oxygen has forced out of the flask. 
 Finally weigh carefully the hard glass tube containing the partially decomposed 
 chlorate. Question. Why is it necessary to cool the test-tube and to have the 
 water in the beaker and flask at the same level before closing the pinch-cock? 
 
 Clean the test-tube by placing water in it and shaking the mixture. The 
 manganese dioxide is difficultly soluble in water, while both potassium chlorate 
 and potassium chloride dissolve readily. Suggest a method of recovering the 
 manganese dioxide and of obtaining a mixture of dry potassium chloride and 
 potassium chlorate practically free from manganese dioxide. 
 
 To make the calculations it will be necessary to know the barometric 
 pressure at the time you perform the experiment, and the temperature of the 
 water in the flask. The temperature of the water may be assumed to be that 
 of the room, and the barometric pressure will be posted on the blackboard. 
 Enter these data in your notebook before leaving the laboratory. Below . is 
 given a table of the vapor-pressure of water at different temperatures. 
 
 Vapor Pressure of Water. 
 
 Temp C. Vapor Pressure Temp. C. Vapor Pressure 
 
 14 2.2 cm. mercury 24 1.2 cm. mercury 
 
 1 6 2.5 " 26 1.3 
 
 18 2.8 " 28 1.5 
 
 20 3-2 30 1-7 
 
 22 3.5 '" 32 2.0 
 
:- , 
 
 Questions Assuming that the levels of the water in the beaker and the 
 flask were the same when the pinch-cock was closed, what was the total pres- 
 sure of the gas in the flask? What was the partial pressure of the water-vapor? 
 Of the oxygen? 
 
 Calculate the volume of the oxygen at i atmosphere pressure and o C, 
 the density of the gas under standard conditions, and the weight of a liter of the 
 gas. Compare your value of the density with the one given by Cady, page 23. 
 
 Problems, i. From your experiment calculate the weight of one liter of 
 oxygen at 2OC and 754 mm. pressure. 
 
 2. Under standard conditions w r hat weight of oxygen would occupy 
 22.4 liters? (Use the density obtained in the experiment.) How is this num- 
 ber related to the atomic weight of the oxygen? 
 
 3. What weight of oxygen could be obtained from i gram of potassium 
 chlorate (KC1O,) by completely decomposing it? 
 
 ASSIGNMENT IV. 
 THE REACTION BETWEEN ZINC AND SULFURIC ACID. 
 
 The purpose of Assignment IV is to determine the quantity of zinc, which, 
 by reacting with an acid, will set free one gram atom of hydrogen. A weighed 
 amount of zinc is allowed to react completely with excess of sulfuric acid and 
 the hydrogen liberated is collected in such a way that its volume can be 
 measured. 
 
 Experiment. Boil about 500 cc. water in your largest beaker to expel the 
 dissolved air. 
 
 Take a piece of zinc weighing about half a gram. If necessary, clean it 
 with sand paper. Weigh to 5 milligrams. Do not try to cut or file the zinc in 
 order to obtain any definite weight. 
 
 Select a beaker of such size that a small funnel when placed in it can be 
 completely covered with water. Place the weighed zinc in the beaker, place the 
 inverted funnel over it, and pour boiled water into the beaker until the funnel 
 is completely covered. 
 
 Pour boiled water into a quarter liter flask until the water completely fills 
 the flask. Without waiting for the water to cool, moisten a piece of filter paper 
 slightly larger than the mouth of the flask, cover the mouth of the flask with 1 the 
 paper, taking care that no bubble of air remains below the paper. Invert the 
 flask (over an empty vessel) and lower it into the beaker in such a manner 
 that the stem of the funnel enters the neck of the flask. If a bubble of air 
 enters the flask repeat this operation. The apparatus now consists of a beaker 
 containing a funnel inverted over the zinc, and a flesk filled with water and 
 inverted over the funnel. Place this apparatus in a basin, or other vessel, to 
 prevent the water from overflowing on the desk during the remainder of the 
 experiment. 
 
 Pour into the water 5 cc. to 8 cc. concentrated sulfuric acid (CAUTION*) 
 If the action of the acid on the zinc is slow add about 5 cc. of the laboratory 
 solution of hydrochloric acid; stir the mixture gently but be careful not to dis- 
 place the zinc from below the funnel. The hydrogen evolved rises through the 
 funnel into the flask, but should not displace all the water in the flask. 
 
 When the zinc has all dissolved (except a few dark-colored flakes of im- 
 purities of negligible weight), place the apparatus in a large basin of water and 
 carefully remove the beaker and funnel without allowing any air to enter the 
 inverted flask. Keep the flask in water until it has cooled to the temperature of 
 the water. Then raise or lower the flask until the level inside and outside the 
 
 * Concentrated sulfuric acid produces dangerous -burns and should not be used care- 
 lessly. The bottle niust not be removed from the lead tray. Carefully pour just enough 
 acid for your experiment into a small beaker. 
 
flask is the same. What is now the pressure of the gas inside the flask?) While 
 the flask is in this position cover the mouth of the flask with the palm of the 
 hand, remove the flask from the water and invert it. While the gas is escaping 
 test to prove that it is hydrogen. 
 
 Fill the flask by means of a graduated cylinder, noting the volume of water 
 needed to take the place of the gas which has escaped. Record in your note- 
 book the barometric pressure (written on the blackboard) and the temperature 
 of the water in which the flask was immersed. 
 
 You now have the weight of zinc taken, and the volume, at a definite temp- 
 erature and pressure, of a corresponding amount of hydrogen saturated with 
 water vapor. What is the partial pressure of the water vapor at the tempera- 
 ture of the experiment? What was the partial pressure of the hydrogen in the 
 flask? 
 
 Calculate from these data : 
 
 1. The volume of the hydrogen under standard conditions of temperature 
 and pressure. 
 
 2. The weight of the hydrogen. .(The density of pure hydrogen under 
 standard conditions is 0.0000899.) 
 
 3. The weight of zinc that would have liberated i gram atom of hydro- 
 gen. Compare this number with atomic weight of zinc. 
 
 From these conclusions write the equation for the reaction. The formula 
 of sulfuric acid is H 2 SO 4 , and both hydrogens are replaced by zinc. 
 
 Problems, i. (a) From the density of hydrogen calculate the weight of 
 22.4 liters of this gas under standard conditions. What is the true molecular 
 weight of hydrogen? What is the formula of this gas? Have you written the 
 above equation correctly? 
 
 (b) If you had not been given the density of hydrogen in the above ex- 
 periment, how could you have calculated a value from the molecular weight? 
 What percentage error would you have made? 
 
 2. (a) What weight of pure zinc sulfate (ZnSO 4 ) could be obtained 
 from the weight of zinc used? 
 
 (b) What is the minimum weight of sulfuric acid (H 2 SO 4 ) necessary in 
 the experiment? 
 
 3. If hydrochloric acid (HC1) had been used instead of sulfuric acid, how 
 much hydrogen would have been liberated? Write the equation. What weight 
 of solid zinc chloride could have been obtained? 
 
 Note on Glass Manipulation. 
 
 To bend a piece of ordinary glass tubing, hold it with both hands in a fan- 
 shaped gas flame and rotate it slowly between the thumb and fingers until a 
 2.y 2 to 3 inch portion is uniformly heated and is soft enough to be bent to the 
 proper angle. Set it aside to cool; glass will remain hot enough to burn the 
 hand for some time after.it no longer appears to be hot. 
 
 To cut glass tubing, scratch it with a file at the proper place, grasp it firmly 
 on each side of this mark (protecting the hands with a cloth), and bend the 
 tube away from the mark. 
 
 Always remove the sharp edges of freshly cut glass at once with a file, or 
 by heating in a gas flame. 
 
 To draw down a piece of tubing to a capillary, heat a portion about i 
 inch long in an ordinary gas flame to a higher temperature than was necessary 
 in bending the tubing. Hold the tube with both hands and rotate it to ensure 
 uniform heating and prevent the hot portion from sagging. Withdraw from the 
 flame and draw apart slowly to obtain a thick-walled capillary. Mount your 
 piece of platinum wire (first used in Assignment VI) by inserting one end of 
 the wire in the capillary near the tube, and heating until the glass closes firmly 
 around the wire. 
 
 10 
 
ASSIGNMENT V. 
 THE ANALYSIS OF COPPER OXIDE 
 (Two students work together.) 
 
 The purpose of Assignment V is to determine the relative weights of 
 copper and oxygen in copper oxide. Hydrogen gas reacts with hot copper 
 oxide to form metallic copper and steam. Two measurements are necessary: 
 the weight of the copper oxide used, and the loss in weight when it is com- 
 pletely reduced to copper. Question : What does this loss of weight repre- 
 sent? 
 
 One of the students signs a white slip for "special apparatus for Assign- 
 ment V", consisting of a thistle tube, a clamp, calcium chloride tube, with two 
 rubber stoppers, 2 glass tubes, 2 rubber tubes, and 2 short glass rods, and 
 returns the apparatus as soon as the experiment is finished. 
 
 Out of a wash bottle make a hydrogen generator by fitting it with a 
 thistle tube (extending almost to the bottorri of the flask) and an outlet tube 
 bent at right angles (see note preceding this Assignment). Place in the wash 
 bottle about 10 grams of zinc, and (to ensure a rapid reaction between Zn 
 and H 2 SO 4 ) cover with a very dilute solution of copper sulfate made by 
 diluting 10 cc. of the laboratory solution to 100 cc. with water. To the out- 
 let tube attach a "drying tube" (containing solid calcium chloride, which has 
 the property of absorbing moisture). Make sure that the apparatus is air- 
 tight, and wrap the flask in a towel. 
 
 Set up the remainder of the apparatus according to the directions of the 
 instructor. Dry the thick-walled glass tube that is to contain the copper 
 oxide by heating it gently. When it is cool weigh it carefully, together with 
 any portion of the apparatus that may come in contact with the 
 copper oxide. Place in the tube about i gram of copper oxide, 
 wipe off any particles that are not in the portion of the tube 
 that is to be heated. Weigh again carefully to obtain the weight of copper 
 oxide used. Attach the apparatus to the hydrogen generator, pour 10 cc. con- 
 centrated sulfuric acid (Caution!) down the thistle tube, and allow the hydro- 
 gen to pass through the apparatus until it has swept out the air. (Caution! 
 Do not place a flame near the outlet nor heat the oxide while the apparatus 
 contains a mixture of oxygen and hydrogen. A dangerous explosion might 
 result.) Suggest a simple test to determine when the hydrogen is no longer 
 mixed with oxygen. 
 
 When pure hydrogen is passing over the copper oxide, begin to heat it 
 very gently with a small flame and continue to heat cautiously until all the 
 oxide changes color. If moisture collects in the farther end -of the tube drive 
 it out by heating the tube carefully. Question : Where does this moisture 
 come from ? 
 
 Allow the tube to cool in the current of hydrogen, and weigh it. If you 
 have time, check this result at once by repeating the heating and weighing; if 
 not, set the tube aside in order that you may do so if the results of the following 
 calculations are unsatisfactory. 
 
 From the results of your experiments calculate the percent of copper 
 and oxygen in copper oxide. 
 
 How many grams of copper are combined with i gram atom of oxygen? 
 Compare this figure with the atomic weight of copper. What, then, is the 
 formula of this oxide of copper? Write the equation for the reaction. What 
 is the true percent of copper and oxygen in this oxide of copper? 
 
 Problems, i. (a) Another oxide of copper is known which contains 
 88.8% of copper. What is its formula? 
 
 (b) Which of the oxides is cuprous oxide and which cupric oxide? 
 
 ii 
 
(c) What are the formulas of cuprous and cupric sulfides? 
 
 (d) Two of the oxides of iron are FeO and Fe 2 CX : which is ferrous 
 oxide arid which is ferric oxide? 
 
 2. How many grams of water can be formed from i gram of cupric 
 oxide? How many grams of hydrogen would be needed? What volume of 
 hydrogen measured at 20 C and i atmosphere would be needed? 
 
 3. How many grams of zinc would be needed to generate enough hydro- 
 gen (by reacting with an acid) to reduce completely one gram molecule of 
 cupric oxide? Compare this weight with the atomic weight of zinc. 
 
 ASSIGNMENT VI. 
 PROPERTIES OF AQUEOUS SOLUTIONS OF ACIDS AND BASES. 
 
 (a) Acids. Prepare a solution of each of the common acids, hydro- 
 chloric, nitric and sulfuric, for use in the following experiments, by adding 
 50 cc. water to 5 cc. of the "dilute" laboratory solution of each of these acids. 
 Label each solution. 
 
 Taste. Taste each solution by dipping a glass rod into the liquid and 
 touching it to the tongue. (Caution: Do not taste any substance in the labor- 
 atory unless directed to do so.) 
 
 Action on Indicators. To about 20 cc. of distilled water in a beaker add 
 enough blue litmus solution to give a distinct but not strong color. Divide 
 this solution into three portions and to each add a few drops of a different acid. 
 
 Repeat, using methyl orange instead of litmus. 
 
 Withdraw a drop of any one of the acids on a glass rod and touch it to 
 blue litmus paper. 
 
 Repeat, using red litmus paper. 
 
 Action on Sodium Carbonate. To 5 cc. of each acid in separate test 
 tubes add a little powered sodium carbonate. 
 
 Action on Metallic Zinc. To 5 cc. of each acid add small pieces of zinc. 
 Is the same gas given off in each case? Test it. Repeat the experiment with 
 zinc and the ordinary laboratory solution of the acids. 
 
 Summarize the properties common to the acids you have studied. Note 
 that any substance is recognized by its properties. Any solution ^ which has 
 the properties thus summarized is an acid solution and contains the substance 
 hydrogen ion. The symbol is H + . 
 
 In addition to the properties of hydrogen ion, the solution of each acid 
 has another group of properties peculiar to itself by means of which the acids 
 may be distinguished from one another. Try the following experiments : 
 
 To 15 cc. water add about 2 cc. silver nitrate solution. Divide the mix- 
 ture into three portions and to each add a few cc. of a different acid. 
 
 In a similar set of experiments use barium chloride solution instead of 
 silver nitrate. 
 
 State how the three acids may be distinguished. 
 
 Thus, in addition to hydrogen ion, hydrochloric acid solution contains 
 chloride ion, C\~ ; nitric acid solution contains nitrate ion, NO 3 ~; and sulfuric 
 acid solution contains sulfate ion, SO 4 ~~. 
 
 Other properties of ions will be discussed in the lectures, and additional 
 evidence of the existence of ions in aqueous solutions will be presented. The 
 study of ionization will be continued in the laboratory in Assignment IX. 
 
 The chemical formulas of these three substances, whose aqueous solutions 
 contain hydrogen ion, are : 
 
 Hydrochloric acid HC1, 
 Nitric acid, HNO 3 , 
 
 Sulfuric acid, H 2 SO 4 . 
 
 Calculate the molecular weight of each acid. Pure HC1 is a gas. and it 
 
 12 
 
is also named hydrogen chloride and hydrochloric acid gas. Pure HNO 3 and 
 H 2 SO 4 are liquids. 
 
 (b) Bases. To 5 cc. portions of the laboratory solutions of sodium 
 hydroxide and potassium hydroxide (taken separately) add 50 cc. water. Use 
 these two solutions and the undiluted laboratory solution of barium hydroxide 
 in the following experiments : 
 
 Taste. Taste each of. these three solutions by withdrawing a drop on a 
 glass rod and touching it to the tongue. 
 
 Touch. Rub a drop of each solution between the thumb and finger. 
 
 Action on Indicators. Treat drops of each of the three solutions with red 
 and with blue litmus paper. 
 
 To 5 cc - of each of the same solutions add a drop of litmus solution. To 
 the sodium hydroxide solution containing the indicator add hydrochloric acid 
 slowly and note the change of color. Then add sodium hydroxide solution 
 until a change is noticed. 
 
 Repeat the last experiment, using fresh 5 cc. portions of the same hydrox- 
 ide solutions and a solution of the indicator methyl orange. 
 
 Repeat, using a solution of the indicator phenolphthalein. 
 
 Action with Ferric Chlcride Solution. To 5 cc. portions of each of the 
 hydroxide solutions add ferric chloride drop by drop. 
 
 Summarize the properties common to the solutions of the bases you have 
 studied. Each solution contains the substance hydroxide ion, OH~. 
 
 In addition to hydroxide ion the solution of each base also has properties 
 characteristic of the base. 
 
 To 5 cc. of each hydroxide solution add a few drops of sulfuric acid. 
 
 Test a drop of each of the hyroxide solutions in a colorless gas flame by 
 means of a looped platinum wire, as demonstrated by the instructor. Can the 
 flame test be used to prove the presence of sodium? 
 
 State how these bases may be distinguished. 
 
 In addition to hydroxide ion, sodium hydroxide solution contains sodium 
 ion, Na + ; potassium hydroxide contains potassium ion, K + ; and barium 
 hydroxide contains barium ion, Ba ++ . 
 
 Each of the hydroxides in the pure state is a solid. The chemical form- 
 ulas for the hydroxides are: 
 
 ^Sodium hydroxide, NaOH, 
 Potassium hydroxide, KOH, 
 Barium hydroxide, Ba(OH) 2 . 
 What is the molecular weight of each hydroxide? 
 
 Problems, i. Make as simple a table as you can to show the colors 
 obtained with different indicators in acid and basic solutions. 
 
 2. Four unlabelled beakers, which are known to contain solutions of 
 sulfuric acid, nitric acid, barium hydroxide, and sodium hydroxide, are to be 
 identified as quickly as possible. State what experiments you would perform, 
 and note clearly the progress towards identification that would be made in 
 each experiment. 
 
 ASSIGNMENT VII. 
 REACTIONS BETWEEN ACIDS AND BASES. 
 
 Experiment. Evaporate , to dryness in a porcelain dish a few drops of 
 (a) hydrochloric acid solution, (b) sodium hydroxide solution. In each case 
 note if there is a residue, add a little water, and test the solution with an 
 indicator. Questions: Taking into account that HC1 and NaOH are both 
 stable substances, state what gases were given off when the solutions were 
 evaporated. What conclusion do you draw in regard to the volatility of 
 hydrogen chloride and sodium hydroxide? How is sodium hydroxide pre- 
 13 
 
pared from metallic sodium and water, hydrogen chloride from hydrogen and 
 chlorine gases? 
 
 Experiment. To 10 cc. sodium hydroxide solution in a beaker add hydro- 
 chloric acid solution slowly until a drop of the mixture, after stirring, reacts 
 acid to litmus paper. Evaporate in a porcelain casserole. While the solution 
 is boiling, test again with litmus paper, and if the solution is no longer acid 
 add HC1. When the solid begins to separate, use a small flame, and, if neces- 
 sary, withdraw the flame and stir the mixture. Finally heat until the mixture 
 is thoroughly dry. Allow the dish to cool, and note the appearance of the 
 residue. Dissolve it in about 10 cc. of water. Test the solution with blue 
 litmus paper and with phenolphthalein test a drop on a looped platinum wire 
 in a colorless gas flame and test a small portion of the solution with silver 
 nitrate in the presence of nitric acid. Questions. Does this solution have 
 either acid or basic properties? What does the solution contain? What was 
 the solid substance obtained on evaporation? What substances were vola- 
 tilized? Why was hydrochloric acid added in excess instead of sodium 
 hydroxide ? 
 
 Write the equation (non-ionic) for the reaction between sodium hydrox- 
 ide and hydrochloric acid, and state what has become of the hydrogen of the 
 acid and the hydroxide radical of the base. t 
 
 This reaction is typical of the action of any acid on any base. 
 
 Write the equations (non-ionic) for the following reactions, consulting 
 a text-book if necessary. 
 
 1. Nitric acid and sodim hydroxide. 
 
 2. Sulfuric acid and sodium hydroxide. 
 
 3. Hydrochloric acid and barium hydroxide. 
 
 Questions and Problems. i. Calculate the weight of sodium chloride 
 that will be obtained on neutralizing one gram of sodium hydroxide with 
 hydrochloric acid. How many molecules of HC1 react with one molecule of 
 Ba(OH) 2 ? What weight of barium hydroxide will be obtained on neutral- 
 izing one gram of barium hydroxide with hydrochloric acid? 
 
 2. What weight of barium hydroxide will react with one mol of hydro- 
 gen chloride? This is called one equivalent of barium hydroxide. 
 
 Suggest a similar definition for one equivalent of sulfuriq acid from its 
 reaction with sodium hydroxide. Note that one equivalent of any acid will 
 react with one equivalent of any base and the resulting solution will contain 
 one equivalent of a corresponding salt. 
 
 How many equivalents of barium hydroxide will react with one mol of 
 sulfuric acid? How many equivalents of barium sulfate will be formed? 
 
 3. Concentration of a Solution. The concentration of a solution is the 
 amount of the dissolved substance contained in a unit volume of the solution. 
 
 The molecular weight in grams (the mol) is frequently adopted as a unit 
 of mass, and the liter is the standard unit of volume. A solution which con- 
 tains in one liter one mol of the dissolved substance is called a mold solution. 
 
 How many grams of sulfuric acid are contained in a liter of a molal 
 solution? 
 
 How many grams of sodium chloride are contained in a liter of a one- 
 tenth molal solution? (o.i M NaCl designates this solution.) 
 
 4. When we are dealing only with neutralization reactions it is conven- 
 ient to use the equivalent (of the acid, or base, or salt) as the unit of mass. 
 A solution which contains one equivalent in a liter is called a normal solution. 
 
 How many grams of HC1 are contained in a liter of a o.io normal solu- 
 tion of hydrochloric acid (o.io N HC1) ? 
 
 How many grams of Ba(OH) 2 are contained in a liter of o.io N barium 
 hydroxide solution? 
 
 How many cc. of o.io N HC1 must be added to 100 cc. of o.io N 
 
Ba(OH) 2 to produce a neutral solution of barium chloride? What color 
 changes would be observed while the acid is being added to the barium 
 hydroxide solution (i) if methyl orange were present, (2) if phenolphthalein 
 weije present? 
 
 ASSIGNMENT VIII. 
 
 TlTRATION OF ACIDS AND BASES. AN ILLUSTRATION OF 
 
 VOLUMETRIC ANALYSTS. 
 
 .Ac-; r^^nt VTII the purpose is to measure the volumes of solutions 
 of r ,acid and base which will exactly neutralize each other. This neutral point, 
 called the end-point, is determined by means of some indicator like litmus or 
 phenolphthalein. The whole experiment is called titration. Question. Does 
 the color of an indicator used in titration change gradually or suddenly as the 
 solution is neutralized? If necessary, perform an experiment to determine 
 this point.) 
 
 Questions. To what volume would you dilute 10 cc. of a 6-normal 
 solution to prepare a 3-normal solution? How would you prepare a o.5-normal 
 solution from a 6-normal solution? 
 
 Experiment. Prepare about 400 cc. approximately o.5-normal sodium 
 hydroxide solution from the laboratory 6-normal solution. Place this 0.5- 
 normal solution in your 500 cc. flask. Prepare also about 250 cc. approxi- 
 mately o.5-normal sulfuric acid solution from the laboratory 6-normal solu- 
 tion, and place it in another flask. Cork and label each flask. Shake the 
 flasks in order that the concentration of the solutions will be uniform through- 
 out. Clean two other flasks, rinse them with distilled water, and set them 
 aside to drain, in order that they may be ready for use later in the experiment ; 
 label one "known H 2 SO 4 solution", and the other "unknown HC1 solution". 
 
 Note. The volumes of the solutions used in a titration are measured by 
 means of burettes. These are glass tubes, marked off in cubic centimeters and 
 tenths (or fifths) of cubic centimeters, and provided with a glass tip out of 
 which the solution may be drawn by opening a pinch-cock. 
 
 Apply at the office for two burettes and clamps, and return them the same 
 day before the office closes. Sign a yellow temporary order slip when you 
 receive these articles, and a blue return slip when you return them. 
 
 Experiment. Fill each burette with distilled water. Note that air may 
 be removed from the small tube below the pinch-cock by tilting the tip upward 
 and allowing the liquid to flow through the pinch-cock. Practice reading a 
 burette : bring your eye to the same level as the liquid and note the reading 
 of the burette corresponding to the bottom of the meniscus ; repeat until con- 
 secutive readings check to better than 0.05 cc. Question. Why is it import- 
 ant to have the eye at the same level as the liquid before making a reading? 
 Never attempt to adjust the volume of the solution in a burette so that the 
 reading will be some exact amount. Allow the water to flow slowly out of 
 the burette. If drops remain on the inner surface of a burette, exchange it at 
 the office for a clean one. 
 
 Rinse one burette with a little of your approximately o.5-normal H 2 SO 4 
 solution, and fill the burette with this solution. Rinse and fill the other burette 
 with the o.5-normal NaOH solution. 
 
 Record the readings of the burettes side by side in your note-book; run 
 about 10 cc. of the acid solution into a clean beaker or flask standing on white 
 paper, and record the final burette reading under the initial reading. Add two 
 drops of phenolphthalein, and about 20 cc. distilled water. Then run in the 
 sodium hydroxide solution from the other burette, a little at a time, and to- 
 wards the end, very carefully, a drop or two at a time, stirring the 
 mixture constantly, until the faintest perceptible .permanent pink color is 
 
 15 
 
obtained. Wash down the inside 'of the beaker by means of a jet of water 
 from the wash bottle. If too much of the basic solution is added, decolorize 
 the solution by adding a little of the acid solution and 4etermine the end-point 
 again. Record the final readings of each burette, and the actual volumes of 
 each solution used in the titration. Calculate the ratio of the two concentra- 
 tions and state which is the more concentrated solution. 
 
 Repeat this experiment, using about 15 cc. and then about 20 cc. of the 
 acid solution, and in each case make the same calculations. Do not fill up the 
 burette before a titration unless there is not enough solution in it for the 
 titration. 
 
 Compare the ratios of the concentrations calculated from these three titra- 
 tions. If any result differs from the average by more than i%, perform 
 additional titrations until you are satisfied that you have determined the ratio 
 of the concentrations with an accuracy better than i%. 
 
 Take your two clean, dry, labelled flasks to the office to obtain a sulfuric 
 acid solution of known concentration, and a hydrochloric acid solution whose 
 concentration you are to determine. 
 
 Empty your acid burette, fill it with the sulfuric acid solution of known 
 concentration, and determine as before from three (or more) titrations the 
 relative concentrations of this solution and your sodium hydroxide solution. 
 Never use less than 10 cc. of any solution in a titration. Calculate the con- 
 centration of your sodium hydroxide solution. 
 
 Finally titrate the unknown hydrochloric solution against the same sodium 
 hydroxide solution, and calculate the concentration of the acetic acid solution. 
 Report this result at once to your instructor. If the result is not sufficiently 
 accurate, first look for an error in your calculations, but, if necessary, repeat 
 the experiment. 
 
 Make a list of the sources of error. (Many of them have been mentioned 
 in the above directions.) 
 
 Save the remainder of the NaOH solution, w 7 hose concentration you have 
 determined, in a corked flask for use in Assignment X. 
 
 Problems. i. If the error in measuring out a volume of solution by 
 means of a burette is o.io cc. what is the percentage error if i cc. of solution 
 is measured? If 20 cc. of solution are measured? Why is it important not 
 to use less than 10 cc. in any titration? 
 
 2. How many cubic centimeters of i molal sulfuric acid will contain o.oi 
 equivalent? How many grams of sodium hydroxide would be needed to neu- 
 tralize this amount of acid? 
 
 3. Chemically pure ("C. P.") sulfuric acid, nitric acid, hydrochloric 
 acid and ammonia, as supplied by the manufacturers, are concentrated 
 aqueous solutions of these substances. The concentration of each solution is 
 guaranteed not to be less than a certain minimum value, and this is tested by 
 measuring the density (specific gravity). The following table contains the 
 density and the percentage composition by weight of the concentrated labora- 
 tory reagents. 
 
 Density % by weight Normal Concentration 
 
 H 2 SO 4 1.84 95-6% H.SO, 
 
 HNO 3 1.42 69.8% HNO 3 
 
 HC1 1.19 37.2% HC1 
 
 NH 4 OH 0.90 28.3% NH 3 
 
 For each reagent calculate the normal concentration and record the 
 results in the fourth column of the table. 
 
 CAUTION! The concentrated acids, especially sulfuric and nitric, pro- 
 duce dangerous burns and should not be used carelessly. These reagents are 
 kept on lead shelves and must not be used unless special directions are given. 
 
 16 
 
ASSIGNMENT IX. 
 STRONG ACIDS. IONIC EQUATIONS. 
 
 In Assignments VI and VII the general properties of acids and bases 
 were studied, and it was recognized that acid solutions contain hydrogen ion, 
 and basic solutions hydroxide ion. The purpose of Assignments IX and X 
 is to compare the concentrations of hydrogen ion in solutions of different 
 acids and to introduce the writing of ionic equations. 
 
 Experiment. Prepare a one-normal solution (about 60 cc.) hydrochloric 
 acid from the 6 TV laboratory reagent (see questions below) and from this 
 solution prepare about 50 cc. each of o.io N, .01 N and .001 N solutions. Be 
 careful to shake each solution after preparing it from a more concentrated 
 solution and water. 
 
 Questions. What volume of 6 N acid is needed to make 100 cc. of i AT 
 acid? 
 
 How much water should be added to 10 cc. of 6 N acid to make i N acid? 
 
 How much water should be added to 10 cc. of i N acid to make o.io 
 N acid? 
 
 What volume of o.oi N acid can be made from 5 cc of o.i acid? 
 
 If hydrochloric acid is ionized completely what is the concentration of 
 hydrogen ion and of chloride ion in each of the solutions prepared? 
 
 Experiment. Pour into marked test tubes 10 cc. of each solution (N, 
 o.i N, o.oi TV o.ooi N}, and pour 10 cc. of w r ater into a fifth test tube. Add 
 to each solution from a glass tube a single drop of methyl violet solution. 
 Hold the tubes in a vertical position, look down at the surface of the solutions, 
 record the color of each solution, and note the smallest concentration of 
 hydrochloric acid that shows with this indicator a different color from water. 
 
 State how the indicator, methyl violet, may be used to determine the ap- 
 proximate concentration of a hydrochloric acid solution. (The color in the more 
 concentrated solutions will fade on standing. It may be restored by adding 
 another drop of the indicator.) Set the o.oi TV solution aside for use in a later 
 experiment (Assignment XIII). 
 
 Repeat the experiment with nitric acid and with sulfuric acid, starting in 
 each case with the 6 TV laboratory acid. Compare the colors obtained with 
 the different acids. Questions. Are the colors characteristic for each acid? 
 If not, what substance determines the color? If hydrochloric acid is com- 
 pletely ionized what conclusion can you draw with respect to the ionization of 
 nitric acid and sulfuric acid? Do these experiments prove that these three 
 acids are highly ionized? 
 
 Summarize briefly the evidence presented in the lectures that these acids, 
 the bases sodium hydroxide and potassium hydroxide, and nearly all salts are 
 highly ionized in dilute solutions. Such substances are called strong electro- 
 lytes, these three acids are strong acids, and sodium hydroxide is a strong base. 
 The term strong salt is seldom used, because there are so few slightly ionized 
 salts. In pure water the concentration of H + and OH~ ions is extremely 
 small and it is an example of a very weak electrolyte. 
 
 We shall next consider what ionic reaction takes place when a strong 
 acid neutralizes a strong base. The equation, 
 
 H,SO 4 + 2NaOH = 2H 2 O + Na,SO 4 , 
 
 which we have been using. hitherto, states that i mol or 2 equivalents of sul- 
 furic acid and 2 mols or 2 equivalents of sodium hydroxide react, and that the 
 resulting solution contains i mol or 2 equivalents of the salt sodium sujfate, and 
 2 more mols of water than were formerly present ; but it does not show the differ- 
 ence between the ions present in the two initial solutions and in the final solution. 
 This difference may be stated in words, or expressed briefly in the following 
 equation : 
 
 17 
 
+ SO 4 -) + (2Na + 4- 2OH-) = 2 H 2 O + 2Na + + SO 4 ~. 
 
 The Na + and SO 4 are the same in the initial and final solutions, and it is not 
 correct to say that they have united to form Na 2 SO 4 . The reaction that has 
 taken place is 
 
 H + + OH- = H 2 O. 
 
 This equation states that i mol or i equivalent of hydrogen ion has combined 
 with i mol or i equivalent of hydroxide ion to form i mol of water. Write 
 the ionic equations for the neutralization of hydrochloric acid and nitric acid 
 solutions by sodium hydroxide solution. State briefly the evidence presented 
 in the lectures to show that the reaction between any strong acid and any 
 strong base is the same in all cases. What reaction will take place whenever 
 a solution with acid properties is mixed with a solution that has basic prop- 
 erties? 
 
 Problems. i. Write ionic equations for the following reactions: 
 
 (a) A sodium chloride solution is evaporated to dryness (solids are 
 not ionized). 
 
 (b) A sodium sulphate solution is evaporated to dryness. 
 
 (c) Hydrogen chloride gas is dissolved in water (gases are not 
 
 ionized). 
 
 (d) A precipitate of silver chloride, AgCl, is formed by mixing 
 
 solutions of silver nitrate and sodium chloride. 
 
 (e) A precipitate of barium sulfate, BaSO 4 , is formed by mixing 
 
 solutions of barium chloride and sodium sulfate. 
 
 2. Does anything happen in the following experiments? 
 
 (a) Dilute solutions of sodium chloride and potassium nitrate are 
 
 mixed. 
 
 (b) Dilute solutions of sodium chloride and nitric acid are mixed. 
 
 3. What is the concentration of Na + , of Cl~, in a solution prepared by 
 dissolving 1.17 grams sodium chloride in i liter of water? 
 
 ASSIGNMENT X. 
 WEAK ACIDS. EQUILIBRIUM. 
 
 Experiment. Prepare normal and o.i normal acetic acid solutions* from 
 the laboratory 6-normal solution. 
 
 Determine if the o. i normal aqetic acid soution has acid properties by 
 testing it with litmus, methyl orange, sodium carbonate, and by tasting it (Cf. 
 Assignment VI). 
 
 Estimate approximately the concentration of hydrogen ion in the normal 
 and o.i normal acetic acid solutions by testing 10 cc. of each solution with a 
 a drop of methyl violet solution, and comparing the colors with those obtained 
 with hydrochloric acid solution and water in the preceding Assignment. 
 Notes and Questions. The reasoning depends upon the comparison of the 
 lowest concentration of solutions of the two acids at which methyl violet gives a 
 color distinctly different from that with water. The acetic acid in the solu- 
 tion must be present either in the form of ions, H + and Ac', or in the un-ion- 
 ized form, HAc. From your estimate of the concentration of H + in the o.i 
 normal solution, calculate the fraction which is ionized, and the fraction which 
 is un-ionized. The fraction of the acid which is in the form of ions is called 
 the degree of ionization. State also the concentrations of acetate ion, Ac", and 
 of the un-ionized acid, HAc, in the o.i normal acetic acid solution. Is acetic 
 acid a weak or a strong acid? 
 
 Questions. The con^ntration of the ions in acetic acid solutions have 
 been determined more accurately than is possible by these color experiments. 
 
 * Save some 0.1 normal acetic acid solution for use in Assignment XIII. 
 
 18 
 
The concentrations of hydrogen ion in these solutions at room temperature are 
 given in the following table. 
 
 Concentration Concentration Concentration Concentration of Degree of 
 
 of acid. of H + . of Ac.~ un-ionized HAc. lonization. 
 
 i N .004 N 
 
 0.5 N .003 N 
 
 o.i N .0013 N 
 
 Fill in the remaining columns of the table. Does the degree of ionization 
 decrease or increase as the concentration of the acid decreases? 
 
 In any given solution containing acetic acid each of the substan- 
 ces, H + , Ac", and HAc, has a definite concentration. The three substances are 
 in equilibrium, and their concentration may be changed by altering the experi- 
 mental conditions, e. g., by raising or lowering the temperature, or by increas- 
 ing or decreasing the concentration of one or more of the substances involved. 
 A careful study of this equilibrium will enable the student to understand more 
 readily the many other examples of equilibrium that will be met with in this 
 course. 
 
 First consider the restriction that these three substances are not indepen- 
 dent of each other, since, namely, H + and Ac~ form HAc when they unite, 
 and HAc forms H + -f- ACT when it ionizes. It is important to realize that 
 whenever there is a disturbance of this equilibrium, either the ions unite to 
 form un-ionized HAc, or HAc breaks up into the ions. These ideas are all 
 involved in the statement that 
 
 H+ + Ac- = HAc 
 
 is a reversible reaction; and the problem is to determine how the reaction may 
 be made to proceed to the right or to the left. 
 
 Questions. What reaction takes place (a) when acetic acid gas is dis- 
 solved in water, (b) when water is added to a normal solution of acetic acid? 
 In the latter case consider the degrees of ionization which you calculated for 
 acetic acid at different concentrations. 
 
 Experiment. The reaction between sodium acetate and hydrochloric acid. 
 Prepare some approximately half normal hydrochloric acid. Measure out 
 three 15 cc. portions in test tubes, and add two or three* drops of methyl violet 
 to each. Measure out $ cc. 2 N sodium acetate solution and add the solution, 
 a few drops at a time, to one of the 0.5 N hydrochloric acid solutions; shake 
 the mixture after each addition, recording the color and estimating the volume 
 of the sodium acetate solution added. Repeat the experiment to check your 
 results. What conclusions can you draw with regard to changes in the hydro- 
 gen ion concentration? (Note that sodium acetate, like other salts, is present 
 in solution as ions, in this case Na + and Ac".) State what reaction has taken 
 place and write the equation. Could the result of this experiment have been 
 predicted from your previous experiments on acetic acid in this Assignment? 
 
 Experiment. The reaction between sodium acetate and acetic acid. Pre- 
 dict what will happen when sodium acetate solution is added to acetic acid 
 solution. Give your reasoning. Test your answer experimentally by adding 
 2 N sodium acetate solution to normal acetic acid containing a drop of methyl 
 violet. 
 
 Review the lectures on equilibrium while answering the following prob- 
 lems. 
 
 Problems.' i. A vessel is partially filled with water; all the air in it is 
 removed, and the vessel is closed. What is the total pressure in the vessel at 
 22 (see Vapor Pressure of Water, Table, Assignment III)? What will the 
 total pressure become if the vessel is heated to 30 ? How has the weight of water 
 
 19 
 
vapor in the vessel changed (qualitatively) ? What reaction has taken place 
 during the change from 22 to 30? If the vessel is again brought to 22 
 what will be the pressure? Write the reaction that has taken place during 
 cooling. Suggest a method for making this reaction go in either direction 
 without changing the temperature. 
 
 2. State in words what is represented by each of the following equations. 
 Suggest experimental conditions under which each reaction can be made to 
 proceed (i) towards the right and (2) towards the left. Under what con- 
 ditions will equilibrium be established in each case? 
 
 (a) HC1 (gas) = H + + Cl~ (in solution) 
 
 (b) Zn (solid) + H 2 O (gas) = ZnO (solid) + H 2 
 
 3. Are the gases, hydrogen, oxygen, and nitrogen slightly soluble or very 
 soluble in water? Give reasons for your answers. What happens when water 
 saturated with one of these solutions is boiled? How is the equilibrium 
 between the gas phase and the solution altered by an increase of temperature? 
 What substances are present in solutions of each of these gases? 
 
 ASSIGNMENT XT. 
 NEUTRALIZATION OF ACETIC ACID SOLUTION BY A STRONG BASE. 
 
 The purpose of Assignment XI is to study the reaction between solutions 
 of acetic acid and of the strong base, sodium hydroxide, and continue the 
 writing of ionic equations. 
 
 Question. How many grams of acetic acid are contained in a liter of 
 molal solution? The formula is HC 2 H 3 O 2 . 
 
 Experiment. Procure two burettes at the office, fill one with your sodium 
 hydroxide solution of known concentration (Assignment VIII), and the other 
 with the special acetic acid solution marked for use in this Assignment. The 
 concentration of the latter solution is in the neighborhood of molal ; the exact 
 value will be announced. Titrate the two solutions, using phenolphthalein as 
 indicator, and save the solutions obtained in the titrations. Note if the end- 
 point is sharp, and if the results of two or more titrations are concordant. 
 Calculate the normal concentration of the acetic acid solution. Questions. 
 What conclusions do you draw from your results? How many mols of acetic 
 acid are neutralized by i equivalent of 'sodium hydroxide? 
 
 Experiment. Acidify the solution obtained in one of the above titrations 
 with a little acetic acid, and evaporate it to dryness, taking care not to heat 
 the residue strongly while driving off the last portions of water and acetic 
 acid. Questions. What is the solid residue? From the results of the titra- 
 tions write the equation (non-ionic) for the neutralization of acetic acid by 
 sodium hydroxide. Point out how it is possible for this solution of a weak acid, 
 in which the concentration of H + is so small, to neutralize so much sodium 
 hydroxide. Which is the weaker electrolyte, acetic acid or water? 
 
 If you have answered the questions in the preceding paragraph correctly 
 you will realize that it is impossible to represent all that has happened in this 
 neutralization reaction in a single equation. However, by following the plan 
 adopted in Assignment IX we can find the one equation that shows most 
 satisfactorily the difference between the initial solutions and the final solution. 
 
 HC 2 H,O 2 -f (Na + + OH-) == H 2 O + Na + + C 2 H 3 O 2 -. 
 
 This equation shows that NaOH and NaC 2 H 2 O 2 are strong electrolytes, that 
 the weak acid is mainly present in the un-ionized form, and that the sodium 
 ion takes no part in the reaction. Accordingly the equation 
 
 HC 2 H 3 2 + OH- = H 2 + C,H 3 2 - 
 20 
 
represents the main reaction. It tells nothing, however, about the "mechanism" 
 of the reaction. 
 
 Problems, i. How many grams of sodium acetate can be prepared (a) 
 from 50 cc. normal acetic acid solution, (b) from 100 cc. 0.5 normal acetic acid 
 solution? 
 
 2. Outline experiments to distinguish between 
 
 (a) i.o TV HNO, and o.i N HNO 3 , 
 
 (b) o.oi N HNO 3 and i.o N acetic acid. 
 
 Optional Experiment. This is to be ommitted by the majority of the 
 class, and should not be porformed without consulting the instructor.' Repeat 
 the titration of the acetic acid and sodium hydroxide solutions, using methyl 
 orange as indicator. 
 
 ASSIGNMENT XII. 
 STRONG AND WEAK BASES. EQUILIBRIUM. 
 
 The purpose of Assignment XII is to compare the concentration of 
 hydroxide ion in solutions of different bases, and to continue the study of 
 equilibrium. 
 
 \ 
 
 Experiment. Prepare solutions of sodium hydroxide which are approxi- 
 mately normal, o.i normal, o.oi normal, and o.ooi normal. (State in your 
 note-book how you prepared these solutions.) To 10 cc. of each solution in 
 a test tube, add i drop of a solution of the indicator, trinitrobenzol. Record 
 the color obtained in each case, and note carefully the most dilute solution that 
 gives a color with the indicator. 
 
 Repeat the experiment with potassium hydroxide solution, and compare 
 the colors obtained at each concentration. If sodium hydroxide in solution is 
 completely ionized, what conclusion can you draw with respect to potassium 
 hydroxide? What concentrations of hydroxide ion can be measured by means 
 of this indicator, trinitrobenzol? 
 
 Try similar experiments with ammonium hydroxide solutions of various 
 concentrations, say normal and o.i normal*. Point out the sodium hydroxide 
 solution and the ammonium hydroxide solution which give approximately the 
 same faint color with the indicator, and estimate the concentration of hydrox- 
 ide ion in this ammonium hydroxide solution. What is the concentration of 
 ammonium ion NH 4 + in this solution? What is the concentration of the un- 
 ionized ammonium hydroxide, NH 4 OH? What is the "degree of ionization"? 
 Is ammonium hydroxide a strong or a weak base? 
 
 In every solution of ammonium hydroxide, just as in the case of acetic 
 acid, there is a definite equilibrium between the un-ionized substance and the 
 two ions. This equilibrium is represented by the equation 
 
 NH 4 OH = NH 4 + + OH-. 
 
 Taking into account the results obtained in these experiments, predict 
 what will happen when solutions of sodium hydroxide and ammonium chloride 
 are mixed. Test your answer by means of an experiment with these solutions 
 and the indicator, trinitrobenzol. Write an equation for the reaction that 
 takes place. 
 
 How will the equilibrium in an ammonium hydroxide solution be disturbed 
 when ammonium chloride is added? What reaction will take place? Test 
 your answer by an experiment with i normal NH 4 OH, ammonium chloride 
 solution and trinitrobenzol. 
 
 Problems. i. An ammonium hydroxide solution has a characteristic 
 odor; this is due to the NH 3 gas given off by the solution. Which solutions 
 
 *Save some 0.01 normal NaOH and 0.1 normal NH OH solutions for use in Assignment XIII. 
 
 4 
 
 21 
 
smell more strongly of ammonia, (a) a dilute solution or a concentrated solu- 
 tion? (b) a cold solution or a hot solution? Write an equation to show the 
 formation of NH 4 OH from NH 3 and water. Is this a reversible reaction? Is 
 there an equilibrium between gaseous NH 3 and the solution? Give reasons for 
 your answers. In this discussion point out what happens (a) when NH 3 gas 
 is passed into pure water, (b) when the resulting solution is boiled in an open 
 vessel, (c) when the gas space above a solution is evacuated. 
 
 2. How will the equilibrium NH 4 OH : : NH 4 + -f OH- be disturbed 
 when sulfuric acid is added to an ammonium hydroxide solution? State 
 how you would prepare ammonium sulfate from ammonium hydroxide and 
 sulfuric acid. Write an equation to represent the main reaction in aqueous 
 solution. 
 
 ASSIGNMENT XIII. 
 HYDROLYSIS. 
 
 In earlier Assignments neutralization reactions have been shown to depend 
 upon the fact that the reaction H + -(- OH~ = H 2 O takes place practically 
 completely. The concentration of H + and OH" in pure water is, however, not 
 zero; each has a definite, although very small, value, namely icr 7 N (i/io 7 ). 
 Water is accordingly a very weak electrolyte ; in other words, the reaction 
 between H + and OH~ does not go quite to completion, and there is in all 
 aqueous solutions an equilibrium 
 
 H,O = H + -f- OH-. 
 
 In Assignment XIII we shall study some experiments closely connected with 
 this fact. 
 
 A solution of sodium acetate may be prepared in two ways : ( i ) by mixing 
 exactly equivalent amounts of acetic acid and sodium hydroxide (reaction, 
 HAc + Na + + OH- = Na + -f Ac- + H 2 O, or more simply, HAc + OH- 
 = Ac" -j- H 2 O), and (2) by dissolving solid sodium acetate in water. Both 
 solutions must have exactly the same properties. 
 
 Prepare some O.OT N NaOH, o.oi N HC1, o.i N acetic acid, and o.i N 
 ammonium hydroxide for use in the following experiments. 
 
 Experiment. To 10 cc. 4 N sodium acetate solution add one drop litmus 
 solution. To another 10 cc. portion of the solution add one drop phenol- 
 phthalein solution. Repeat with 10 cc. portions of distilled water. What 
 results did you obtain in similar experiments with NaCl solution in Assign- 
 ment VII? (Repeat these experiments if necessary.) 
 
 Questions. In what respect does the sodium acetate solution differ in 
 properties from water and from a sodium chloride solution? Is the reaction 
 between equivalent amounts of acetic acid and sodium hydroxide complete? 
 What substances are present in the resulting solution? State what happens 
 when solid sodium acetate is dissolved in water. The reaction that takes place 
 between the acetate ion and water is an example of hydrolysis. Write this 
 reaction as well as you can in a single equation, and compare the equation 
 with that given above for the neutralization of acetic acid by sodium hydrox- 
 ide. Which of the following substances in the 4 N sodium acetate solution 
 have (i) large concentrations, (2) small concentrations, and (3) extremely 
 small concentrations: H 2 O, Na + , Ac", HAc, OH~, H + . Test your answer by 
 the following experiments : 
 
 Experiment. To 10 cc. 4 N sodium acetate solution containing i drop 
 phenolphthalein add o.i N acetic acid, drop by drop. State in words what 
 happens in this reaction. Write the equation for the main reaction. 
 
 Experiment. To 10 cc. distilled water containing i drop phenolphthalein 
 add o.oi N NaOH solution drop by drop until (after shaking) the color is the 
 
 22 
 
same as in the 4 N NaAc solution. Estimate the volume of one drop by 
 counting the number of drops in, say 5cc., and calculate the concentration of 
 hydroxide ion in 4 N sodium acetate. What is the concentration of un-ionized 
 acetic acid in the same solution? What is the concentration of Na + and of 
 Ac"? (Assume complete ionization of sodium acetate and sodium hydroxide.) 
 
 Questions. Do these results furnish any evidence that the reaction H + -)- 
 OH" = H 2 O is reversible? How is this equilibrium disturbed by the addition 
 of sodium hydroxide? Is the concentration of hydrogen ion in the 4 N 
 sodium acetate solution greater or less than in pure water? 
 
 Experiment. Repeat the above experiments with 10 cc. 4 TV ammonium 
 chloride. State whether the solution has an acid or basic reaction. 
 
 To 10 cc. of 4 N ammonium chloride containing a drop of litmus solution 
 add gradually o.i N NH 4 OH until the solution is neutral. What reaction has 
 taken place? Discuss the hydrolysis of ammonium chloride and the neutraliz- 
 ation of the weak base ammonium hydroxide by an equivalent amount of 
 hydrochloric acid. Write a single equation to represent the main reaction in 
 each case. Devise and try an experiment to determine the hydrogen ion 
 concentration in 2 N ammonium chloride. 
 
 Questions. How do the concentrations of H + and OH~ in pure water and 
 in solutions of NaCl, NaNO 3 , K 2 SO 4 , compare with each other? What con- 
 clusions can you draw with regard to (a) the hydrolysis of the salt of a strong 
 acid and a strong base, (b) the hydrolysis of the salt of a weak acid and a 
 strong base, as potassium acetate, (c) the hydrolysis of the salt of a strong 
 acid and a weak base, as ammonium nitrate? 
 
 State whether or not you would expect a salt of a weak acid and a weak 
 base (as ammonium nitrate) to be hydrolyzed. Ammonium acetate, like other 
 salts, is completely ionized in dilute solution. When solid ammonium acetate 
 is dissolved, what reaction will take place between ( i ) ammonium ion and 
 water, (2) between acetate ion and water? Write a single equation to repre- 
 sent the hydrolysis. Will the neutralization reaction between equivalent 
 amounts of acetic acid and ammonium hydroxide take place completely? Will 
 it take place more or less completely than the neutralization of acetic acid by 
 sodium hydroxide? 
 
 Experiment. Test a normal solution of ammonium acetate with indicat- 
 ors. Explain your results, taking into account the fact that acetic acid and 
 ammonium hydroxide are ionized to about the same extent. 
 
 Problems, i. In the table given below fill in the formulas of the com- 
 pounds which the given positive and negative ions form with each other. 
 
 Now mark with a star (*) those compounds which are only slightly ion- 
 ized in solution, and mark with an "h" those compounds which are hydrolyzed 
 in solution. 
 
 Study this table carefully and point out parallelisms between the strengths 
 of acids and bases and the hydrolysis of the corresponding ions. 
 
 H + Na + K + NH 4 + 
 
 OH- 
 
 ci- 
 
 NO.,- 
 
 so 4 - 
 
 Ac- 
 
 2. Summarize the behavior of solutions towards the five indicators you 
 have used in Assignments VI to XIII by completing the following table. 
 
 23 
 
Concentration. Methyl violet. Methyl orange. 
 
 H + i.o normal yellow red 
 
 H + o.i 
 H+ o.oi 
 
 H + o.ooi violet 
 
 H + slightly greater than in red 
 
 water yellow 
 
 H + and OH~ equal in 
 
 water 
 OH" slightly greater than 
 
 in water 
 
 OH~ o.oi normal 
 OH- o.i 
 OH- i.o 
 
 State how you would estimate roughly the H + cr OH" concentration in any 
 colorless unknown solution. Suggest a definition for an indicator. 
 
 ASSIGNMENT XIV. 
 CARBON DIOXIDE, CARBONATES, BICARBONATES, CARBONIC ACID. 
 
 In Assignments VI and X solutions of acids were found to react with 
 sodium carbonate to give a colorless, odorless gas. This gas is carbon dioxide, 
 CO 2 . In Assignment XIV we shall study this reaction more carefully, .and 
 shall investigate the behavior of an aqueous solution of carbon dioxide. 
 
 Experiment. Add a small quantity of powdered sodium carbonate, 
 Na 2 CO 3 , to 10 cc. water, shake the mixture several times, and filter if there is 
 a residue. Add some of the clear solution, a little at a time, to a solution of 
 (a) hydrochloric acid, (b) acetic acid. Is Na 2 CO 3 a readily soluble or a 
 difficultly soluble substance? Like other salts, it is highly ionized in aqueous 
 solution. Complete the equations 
 
 H + + CO,- = CO, -f 
 HAc + CO 3 ~ = CO 2 + 
 
 A Test for Carbon Dioxide. Moisten the end of a glass tube with Ba( OH) 2 
 solution and hold it over a solution from which CO 2 bubbles are issuing 
 (or pass CO 2 gas into i or 2 cc. Ba(OH) 2 solution). The white solid is 
 barium carbonate, BaCO 3 , and its formation from Ba(OH) 2 solution is a test 
 for carbon dioxide. Prove that the breath contains CO 2 . Is BaCO, readily 
 soluble or difficultly soluble in water? Complete the equation 
 
 CO 2 + Ba ++ + OH- = BaCO 3 + 
 
 Note that Ba(OH) 2 is a strong base. Outline an experiment to prove this. 
 
 Experiment. Repeat the first experiment, using sodium bicarbonate, 
 
 NaHCO,, instead of Na 2 CO 3 . Report whether it is readily soluble or diffi-. 
 
 cultly soluble in water. Prove that the gas evolved is CCX. Complete the 
 
 equati ns H+ + HC0 3 - = CO. + 
 
 HAc + HCO 3 - = CO 2 + 
 
 and compare them with the equations for the corresponding reactions with 
 ca'rbonate ion. 
 
 Experiment. Limestone is calcium carbonate, CaCO 3 . Is it readily or 
 difficultly soluble in water? Try an experiment, if necessary. Test the action 
 of hydrochloric and acetic acid on small amounts of solid calcium carbonate, 
 and complete the equations 
 
 H + -f- CaCO 3 (solid) = 
 HAc + CaCO 3 (solid) = 
 
 24 
 
State what you think will happen when solid BaCO 3 is treated with hydro- 
 chloric acid. Try the experiment and write the equation for the reaction. 
 
 Set up a carbon dioxide generator similar to that used in preparing hydro- 
 gen (Assignment V) ; sign a white slip for the thistle tube. Place in the flask 
 a few small lumps (about 5 grams) of limestone, 'cover with water, and add 
 a little hydrochloric acid through the thistle tube. 
 Show by experiments : 
 
 Whether CO 2 is denser or lighter than air. 
 Whether it is inflammable or supports combustion. 
 Whether or not it dissolves in water. Give details of the last experi- 
 ment. 
 
 The solution of carbon dioxide in water contains carbonic acid, H 2 CO 3 . 
 Prove by experiment that 
 
 CO 2 (gas) + H 2 O = H 2 CO 3 (in solution) 
 
 is a reversible reaction, and that therefore is an equilibrium. At room tempera- 
 ture a solution in equilibrium with CO 2 gas at i atmosphere pressure con- 
 tains about 0.04 mol H 2 CO 3 in i liter. How is this equilibrium altered by an ' 
 increase of temperature? 
 
 Properties of a Solution of Carbonic Acid. Prepare a small quantity of 
 a nearly saturated solution of CO 2 and test it with indicators (see Problem 
 2, Assignment XIII). Is H 2 CO 3 a weak or a strong acid? Decide if possible 
 whether it is weaker or stronger than acetic acid. The equilibrium in the 
 solution depends on the reversible reaction 
 
 H 2 CO 3 = H + -f HCO 3 -. 
 
 Explain what happens in the following experiments : 
 
 To 10 cc. water add i drop litmus solution,] ust enough to give a faint 
 bluish color. Pass in CO 2 gas until the color changes. Heat the solution to 
 boiling and cool it again. 
 
 To 10 cc. water containing i drop phenolphthalein add a drop of dilute 
 NaOH solution, just enough to give a faint pink color; pass in CO, gas. 
 
 Review the above experiments on the action of acids on NaHCO 3 solution, 
 Na 2 COo solution, and solid CaCO 3 , and state what happened in the solution 
 before CO 2 gas was evolved. 
 
 The Neutralization of Carbonic Acid in Steps. Which of the following 
 equations represents the main reaction when solutions of equimolal quantities 
 of H 2 CO, and NaOH are mixed? 
 
 H 2 CO 3 -f OH- = HCO 3 - + H,O 
 H+ -f OH- = H 2 O 
 
 Experiment. The reaction between bicarbonate ion and hydroxide ion. 
 To a solution of sodium bicarbonate containing trinitrobenzol, add slowly from 
 a graduated cylinder 0.5 normal NaOH solution. Perform a blank experiment 
 (for comparison) by adding the NaOH solution to an equal volume of water 
 containing the same indicator. Have you any evidence that the reaction 
 
 HCO 3 - + OH- = H 2 O -f- CO 3 ~ 
 
 has taken place? Is it correct to say that bicarbonate ion is an acid? Try to 
 devise an experiment to test if the reaction just written is reversible. 
 
 Write the equation for the main reaction between H 2 CO 3 and excess 
 sodium hydroxide solution. 
 
 Experiment. The hydrolysis of carbonate ion. Test 10 cc. portions of 
 normal Na 2 CO 3 solution with indicators. To the portion containing phenol- 
 phthalein add o.i normal hydrochloric acid very slowly from a graduate. 
 Compare your results with those obtained with sodium acetate solutions in 
 Assignment XIII, and state which solution has the greater concentration of 
 OH~. The reaction between carbonate ion and water is 
 
 25 
 
CO 3 ~ + H 2 O = HCO 3 - -f OH-, 
 
 and this experiment furnishes a proof that HCO 3 ~ is a very weak acid. Which 
 is the weaker acid, acetic acid or bicarbonate ion ? Give your reasoning. 
 
 It is evident from the preceding experiments that the weak dibasic acid 
 H,CO 3 ionizes in two stages, namely, 
 
 H,CO, = H + + HCO,-, and 
 HCCV = H + + CO,-. 
 
 The second acid, HCO 3 ~, is a much weaker acid than the first, H 2 CO :; . Sug- 
 gest experiments or give reasons to prove this fact. There are two series of 
 salts : the "normal salts" as Na 2 CO 3 and CaCO 3 and the "acid salts" as 
 NaHCO :i . The chief negative ions present in solutions of these two types of 
 salts are CO 3 ~~ and HCO 3 ~ respectively. Sodium bicarbonate is also named 
 sodium acid carbonate or sodium hydrogen carbonate ; and the bicarbonate ion 
 is named similarly the acid carbonate ion or the hydrogen carbonate ion. 
 
 The behavior of other weak polybasic acids is similar to that of carbonic 
 acid. Thus the weak dibasic acid, hydrogen sulfide in solution, H 2 S, ionizes 
 in two stages, and there are two series of salts. The second acid. HS~, is 
 weaker than the first, H,S. 
 
 Sulfuric acid, H 2 SO 4 , is an example of a strong polybasic acid. Review 
 your experiments with H 2 SO 4 in Assignments VIII, and IX, and decide 
 whether hydrogen sulfate ion, HSO 4 ~, is a strong or a weak acid. What ions 
 are present in large quantity in a dilute solution of potassium acid sulfate, 
 KHSO 4 ? Try experiments if necessary. 
 
 Problems, i What general statement can you make about the action of 
 an acid on a solution of a salt of a weaker acid? Give examples and write 
 equations. 
 
 2. A solution of sodium bicarbonate is prepared by dissolving 8.4 grams 
 of this salt in enough water to give i liter of solution, (a) Calculate approxi- 
 mately the concentration of HCO 3 ~ in this solution. (b) What weight of 
 sodium carbonate can be prepared from this solution? (c) How many cc. 
 of normal NaOH solution are necessary in (b) ? What volume of CO 2 under 
 standard conditions can be prepared from this quantity of NaHCO 3 ? 
 
 Optional Experiment. Test a solution of NaHCO 3 with indicators. 
 
 Considering the facts that H 2 CO 3 is a stronger acid than HCO ; r, predict 
 what will happen when CO 2 is passed into a solution of NaXCX. Test your 
 answer experimentally. Write the equation for the reaction. 
 
 Predict what will happen when a solutoin of NaHCO 3 is boiled. How 
 would the concentration of OH~ change during the boiling? Test your answer 
 by experiments. Write the equation or equations. Is the reaction considered 
 in the preceding paragraph reversible? 
 
 Note. When you have completed Assignment XIV, clean your two sample 
 bottles, label them with your desk number and name, mark them No. i and 
 No. 2 respectively, and deposit them at the office. Your two unknowns (for 
 analysis) will be ready for you at the next laboratory period. See Assign- 
 ment XV. 
 
 ASSIGNMENT XV. 
 REVIEW. ANALYSIS OF SOLUTIONS FOR IONS ALREADY STUDIED. 
 
 Give formulas and names of four or five of the common sodium, potas- 
 sium and ammonium salts. Are these salts readily soluble or difficultly 
 soluble in water? Are all of them highly ionized in dilute solution? 
 
 What is the color of each of the ions, Na + , K + , NH 4 + ? Give other prop- 
 erties of these ions, and specify a characteristic property of each that may be 
 used as a test. Devise methods of testing for each of these ions in the pres- 
 ence of the other two. Try your methods with known solutions : for example, 
 divide a solution contining Na + and NH 4 + into two parts, to one part add a 
 
 26 
 
small quantity of a potassium salt solution, and be sure that you can distin- 
 guish between these two solutions. Is it Na + or NH 4 + that causes difficulty 
 in testing for K + ? 
 
 Suggest a method of testing a solution for the presence of a carbonate 
 based on the evolution and detection of CO 2 . Try your method. 
 
 State how you would test a solution for the presence of Cl~, for the 
 presence of SO 4 ". Write the ionic equations. Apply these tests to a solu- 
 tion of a carbonate free from Cl~ or SO 4 ~ in order to determine if carbonate 
 interferes with your tests. If it does, suggest and try a method of preventing 
 this interference. 
 
 Experiment. Heat to boiling a normal solution of (NH 4 ) 2 SO 4 in a small 
 beaker, and note if there is an odor. Now add i cc. 6 normal NaOH, and 
 again heat to boiling. While you are heating the mixture note the odor and 
 test the vapor with moist litmus. Questions. What reaction takes place (a) 
 when cold solutions of (NH 4 ) 2 SO 4 and NaOH are mixed? (b) when the 
 solution is boiled? What volume of 6 normal NaOH would be needed to 
 react completely with 5 cc. normal (NH 4 ) 2 SO 4 ? If a mixture consisting of 
 5 cc. normal (NH 4 ) 2 SO 4 and 3 cc. 6 normal NaOH were evaporated to dry- 
 ness, what substances would be present in the solid residue? 
 
 Experiment. Mix 5 cc. normal NaCl and 3 cc. 6 normal H 2 SO 4 in your 
 casserole, evaporate the mixture on the porch* until heavy white fumes of 
 sulfuric acid are given off. When the residue is cold, pour it into cold water, 
 and test the solution for chloride. Questions. Did any reaction take place 
 (a) when the two solutions were mixed, (b) when the mixture was evapor- 
 ated? Which is the more volatile, HC1 or H 2 SO 4 ? 
 
 Test the volatility of nitric acid and acetic acid by evaporating solutions of 
 these acids in a porcelain dish on the porch. 
 
 Salts of sodium and potassium, and of metals in general, are not easily 
 volatilized. Test the volatility of ammonium salts, as NH 4 C1 and (NH 4 ) 2 SO 4 
 by heating small quantities of the solids in a porcelain dish on the porch. 
 
 Analysis Nos. i and 2. Analyze the two "unknowns" for 
 H + , Na% K + , NH 4 + , 
 OH-, C1-, S0 4 - C0 3 - 
 
 In these and all later analyses estimate roughly the concentration of H + or OH~ 
 in the original solution ^see Problem 2, Assignment XIII). 
 
 (Optional work. Distinguish carbonates and bicarbonates.) 
 While making these analyses record in your note-book all your experi- 
 ments and observations, and state clearly the conclusions that you draw. 
 Submit this written report of your analyses to your instructor before pro- 
 ceeding with the next Assignment. 
 
 Problems, i. State how you would convert: 
 
 (a) Solid sodium carbonate into solid sodium nitrate. 
 
 (b) Solid sodium nitrate into solid sodium sulfate. 
 
 (c) Solid ammonium chloride into solid sodium chloride. 
 
 2. How would you recover the potassium chloride from a mixture of 
 potassium chloride and ammonium chloride. 
 
 ASSIGNMENT XVI. 
 CHEMISTRY OF CALCIUM. 
 
 In Assignment XVI we shall study the chemistry of metallic calcium and 
 of calcium ion. We shall find that certain compounds of this element differ 
 from the corresponding compounds of sodium, potassium, and ammonium in 
 
 * All operations which give rise to noxious fumes must be performed on the 
 porch, not in the laboratory. 
 
 2 7 
 
that they dissolve to a much smaller extent in water; and we shall study the 
 equilibrium between these solids and their "saturated" solutions. Review the 
 lectures on calcium.) 
 
 The solubility of a substance in water is the concentration of its satur- 
 ated solution. Question. If the solubiltiy of NaCl in water at room tempera- 
 ture is 5.4 mols per liter, how many grams of NaCl are contained in 100 cc. 
 of saturated solution? How could you prove that NaCl (solid) = NaCl (in 
 solution) is a reversible reaction? 
 
 Note. In future write equations for all reactions. 
 
 Obtain from the office a piece of metallic calcium. Describe its proper- 
 ties as far as you can observe them by physical examination. In what 
 respects does calcium show the physical properties of a metal? 
 
 Experiment. Drop the calcium into 20 cc. of water. Stir the mixture or 
 warm gently until the metal has dissolved. Test the solution for OH". What 
 reaction occurs? The white solid formed, calcium hydroxide, is a strong base 
 but only slightly soluble. Question. Name other metals that react readily 
 with water? What reactions occur? 
 
 Write the ionic reaction for the dissolving of solid calcium hydroxide to 
 form a saturated solution, and indicate the equilibrium involved. How can 
 the equilibrium be shifted so that more Ca(OH) 2 will dissolve? So that less 
 will dissolve? 
 
 Experiment. Divide the solution of Ca(OH) 2 and suspended solid into 
 three parts. Set aside one of these (corked) for future use. To one part 
 add 10 cc. of i N NH 4 C1. Heat the solution and test the odor. To another 
 part add 3 cc. of 6 N HC1. What reactions have taken place in these two 
 cases? Note the relation of these reactions to the equilibrium discussed in 
 the preceding paragraph. 
 
 To 5 cc. i N CaCU add Na 2 CO 3 solution in excess, heat the mixture to 
 boiling, filter. (The instructor will explain the details of the process of filtra- 
 tion.) The white precipitate is calcium carbonate. Write the simplest ionic 
 equation for its formation. Prove that the fiiltrate contains chloride ion. 
 
 The fact that the amount of calcium ion in the filtrate must be very small 
 can be deduced from the small solubility of CaCO 3 in water, 0.00013 mols per 
 liter. 
 
 In any solution saturated with solid calcium carbonate there is an equili- 
 brium between the solid substance and its ions, which may be expressed by 
 the equation 
 
 CaCO 3 (solid) = Ca ++ -f CO 3 ". 
 
 When will a precipitate of calcium carbonate be formed? How is this equi- 
 librium disturbed by the addition of a strong acid? Review or repeat your 
 experiments (Assignment XIV) on the reaction of hydrochloric and acetic 
 acids on CaCO 3 . Which is the stronger acid, HAc or H 2 CO 3 ? 
 
 Experiment. Pass CO 2 gas into a solution of CaCL. What conclusion 
 do you draw with regard to the concentration of CO 3 ~ in a solution of 
 carbonic acid? 
 
 To a solution of CaCl 2 add NaOH solution slowly a few drops at a time. 
 Repeat the experiment with NH 4 OH solution. What conclusions do you 
 draw with regard to the relative concentrations of OH" in NaOH and NH 4 OH 
 solution? 
 
 Now pass CO 2 gas into the CaCl 2 solution containing NH 4 OH, and warm 
 the mixture gently. 
 
 Filter the third portion of the mixture obtained by the action of calcium 
 on water, and pass a few bubbles of CO 2 gas into the filtrate. Under what 
 conditions will CO 2 gas form a precipitate of CaCO 3 in a solution containing 
 Ca ++ ? Which is the less soluble substance, CaCO 3 or Ca(OH) 2 ? 
 
 28 
 
Test the action of excess CO 2 on the same mixture by passing CO 2 gas 
 into a small quantity of the mixture for some time. The reaction involves 
 the formation of bicarbonate ion, HCO 3 ", H 2 CO 3 + CaCO 3 (solid) = Ca ++ -f- 
 2HCO 3 ~. There is an equilibrium. Point out how the reverse reaction can 
 be made to take place. (Optional Question. Point out how these results are 
 related to the results of the optional experiments at the end of Assignment 
 XIV.) 
 
 Experiment. Determine whether calcium salts give a characteristic flame 
 test. Can you distinguish the flame with a calcium salt from that with sodium 
 or potassium salts? 
 
 Experiment. To 10 cc. normal CaCl 2 solution add 2 cc. 6 normal 
 HoSO 4 . If no precipitate appears at once, heat the solution gently, and let it 
 stand. Filter. Test a portion of the filtrate for Ca ++ by adding NH 4 OH 
 until the solution is no longer acid, and then (NH 4 ) 2 CO 3 solution, and warm 
 the mixture. Test another portion for SO 4 by adding BaCL solution. What 
 conclusion do you draw with regard to the solubility of CaSO 4 ? Which is 
 the less soluble, (i) CaSO 4 or CaCO 3 , (2) CaSO 4 or BaSO 4 ? 
 
 Predict what will take place when solid CaSO 4 is heated with excess 
 NaXO 3 solution? Experiment. Test your prediction by boiling some solid 
 CaSO/with normal Na,CO 3 solution/ Test the filtrate for SO 4 ". Wash 
 the precipitate with water, and test it for carbonate. 
 
 Test whether the reverse reaction will take place by heating solid CaCO 3 
 with sodium sulfate solution, filtering and testing the precipitate and filtrate. 
 
 Problems. i. Arrange the compounds of calcium according to their 
 solubilities in water, distinguishing readily soluble, moderately soluble, and 
 difficultly soluble substances. Point out the compounds which are much more 
 soluble in dilute hydrochloric or nitric acids than in water. 
 
 2. Can the following substances be present at high concentrations in the 
 same solution? If not, what is formed? 
 
 (a) H + and NO,- (e) H + and HCO 3 - 
 
 (b) H + and OH- (f) Ca ++ and CO 3 ~ 
 
 (c) H+ and S0 4 - (g) Ag + and Cl~ 
 
 (d) H + and CO 3 ~ (h) H 2 CO 3 and CO 3 " 
 
 ASSIGNMENT XVII. 
 
 ACTION OF ACTDS ON METALS. SOLUBLE SALTS OF ZINC, 
 COPPER AND SILVER. 
 
 Describe the physical properties of zinc, copper and silver, noting any 
 characteristic property of each metal. 
 
 Zinc and Acids. Review your experiments on the action of acids on 
 zinc. Write the ionic equations for the reaction between zinc and solutions 
 of hydrochloric and sulfuric acids. What is the color of Zn ++ ? Is it correct 
 to say that an atom of the metal has combined with 2d~ or with SO 4 "? How 
 can solid zinc chloride and solid zinc sulfate be prepared from these solu- 
 tions ? 
 
 From an examination of the ionic equation which you have just written 
 state how you would expect the speed of the reaction to alter with the concen- 
 tration of H + . Experiment. Test your answer by treating small pieces of 
 zinc with (a) a small volume of 6 normal HC1, (b) the same amount of HC1 
 in a large volume of water, and. (c) 6 normal acetic acid solution. In the 
 same experiments determine how the speed of the reaction is altered by an 
 increase of temperature. 
 
 When Zn reacts with 6 normal HNO 3 the principal reaction, written in 
 its simplest form, is 
 
 29 
 
3Zn( solid) + 2NO 3 - -f 8H + == 3Zn ++ -f 2NO(gas) -f 4H 2 O 
 
 NO gas is colorless. What is the reaction between NO and oxygen of the air? 
 Repeat the experiment with Zn and HNO 3 solution, if necessary. Place the 
 metal and acid in a small beaker and cover it with a watch glass. 
 
 Copper and Acids. Experiment. Try the action of 6 normal HC1 and 
 of 6 normal H 2 SO 4 on metallic copper. After a rew minutes remove the 
 pieces of metal, wash them with water, and note if the appearance has changed. 
 Then test the action of fresh solutions of the acids on the clean metal. Is 
 there any evidence of chemical action? 
 
 Try the action of 6 normal HNO 3 on metallic copper. Save the solution. 
 What is the color of the copper ion (cupric ion), Cu ++ ? The reaction is 
 similar to that between Zn and nitric acid solution. Write the equation. 
 
 Silver and Acids. Experiment. Try the action of 6 normal HC1, 6 
 normal PLSO.,, and 6 normal HNO 3 on silver. Compare with the correspond- 
 ing actions in the case of zinc and copper. Save any solution you obtain. 
 What is the color of silver ion, Ag + ? Write the equation for the reaction 
 with nitric acid solution, noting that the change from i atom Zn to Zn ++ 
 corresponds to the change from 2 atoms Ag to 2Ag + . 
 
 A summary of the results of these experiments is asked for in Problem 
 i below. 
 
 Solubility of Nitrates. Chlorides and Sulfates of Metals. What conclu- 
 sion can you draw from the above experiments with regard to the solubility 
 of the nitrates of zinc, copper and silver. State how you could prepare the 
 solid nitrates from your solutions. Give the formulas. Note that the nitrates 
 of all metals are soluble in water. 
 
 Experiment. To a small portion of your solution containing Ag + add 
 a few drops HC1 solution; to another portion add a few drops H 2 SO 4 solu- 
 tion. Repeat these experiments with your solution containing Cu ++ . What 
 conclusions can you draw with regard to the solubility of AgCl, Ag 2 SO 4 , 
 CuCL and CuSO 4 in water? Note that nearly all the chlorides and sulfates 
 of metals are soluble in water. There are only a few exceptions in each 
 case. Give an example of a difficultly soluble sulfate. 
 
 Preparation of Sulfates from, Nitrates and Chlorides. Suggest a method 
 of preparing solid copper sulfate from copper nitrate based on the difference 
 in volatility of HNO, and H 2 SO 4 . Try your method with your solution 
 which contains Cu ++ , NO 3 ~ and H + (which is usually spoken of as a solution 
 of copper nitrate and nitric acid). Collect some crystals of copper sulfate by 
 evaporating the final solution to a very small volume, and letting it cool 
 slowly. Wash the crystals with a very little water and save them. 
 
 Suggest a similar method of preparing solid zinc sulfate from a solution 
 of zinc chloride. How would you test whether the final product is free from 
 chloride ? 
 
 The problem of preparing a soluble chloride or nitrate from a sulfate 
 will be considered later. Suggest a method now if you can. 
 
 Conversion of Soluble Chlorides into Nitrates, and Nitrates into Chlor- 
 ides. Experiment. Mix 2 cc. concentrated HNO 3 solution and 5 cc. concen- 
 trated HC1 solution, and let the mixture stand. Is there any evidence of 
 chemical action? The same reaction takes place when a dilute solution of 
 the two acids is evaporated to a small volume. Do not attempt to write the 
 equation. Both acids are destroyed, and this reaction may be used to remove 
 completely chloride or nitrate from a solution by adding a concentrated solu- 
 tion of the other acid in large excess and evaporating almost to dryness. 
 
 Experiment. Dissolve about 0.5 g. zinc in hydrochloric acid, evaporate 
 the solution to a small volume, add some concentrated HNO 3 solution, and 
 evaporate the mixture almost to dryness on the porch. Test the residue for 
 
 30 
 
chloride, and repeat the treatment with concentrated nitric acid, if necessary. 
 
 How would you reconvert the zinc nitrate into zinc chloride? 
 
 Water of Crystallisation. Note the appearance of the copper sulfate 
 crystals you have prepared. Examine also the large crystals in the bottle of 
 copper sulfate in the laboratory. The formula is CuSO 4 5H 2 O, and the water 
 present in the solid is called "water of crystallization". Many other salts, 
 such as crystalline zinc sulfate, also contain water of crystallization. 
 
 Experiment. Heat a small quantity of CuSO 4 5H 2 O in a porcelain evap- 
 orating dish, and note the change in its appearance. The residue is anhy- 
 drous copper sulfate, CuSO 4 . Write the equation for the reaction. When 
 the dish is cool add a few drops of water. What reaction will take place 
 when water vapor is passed over anhydrous CuSO 4 at a low temperature? 
 Under what conditions will there be an equilibrium? 
 
 Give another example of an equilibrium involving two solids and a gas. 
 How does the pressure of the gas at equilibrium vary with the temperature? 
 
 Problems, i. Summarize what you know about the action of acids on Na, 
 K, Ca, Zn, Cu, Ag and any other metals discussed in the lectures by dividing 
 the metals into the three following classes : 
 
 1. Those which react readily with water. Hydrogen is evolved. 
 
 II. Those which react readily with HC1 or H 2 SO 4 solution, but not 
 with water. Hydrogen is evolved. 
 
 III. Those which dissolve readily in HNO 3 , but not in HC1, H 2 SO 4 or 
 
 water. H 2 is not evolved. 
 Write the equation for the action of HC1 or H 2 SO 4 on metallic sodium. 
 
 2. How would you prepare solid silver sulfate from ' metallic silver? 
 What weight of the sulfate could be obtained from I gram of silver? 
 
 ASSIGNMENTS XVIII TO XX. 
 CHEMISTRY OF CUPRIC ION, SILVER ION, AND ZINC ION. 
 
 In the next three Assignments we shall study the chemistry of the ions 
 Cu ++ , Ag + , and Zn ++ , especially the preparation of difficultly soluble com- 
 pounds, methods of dissolving them, and the equilibria involved. 
 
 While doing this work collect the information necessary to complete the 
 following table, performing additional experiments when necessary. This 
 table is intended to summarize the action of various reagents on the different 
 positive ions. Leave a blank when there is no action. Mark the colors of 
 the precipitates, and add notes in the table, or below the table, to show 
 methods of dissolving them. Be sure that you can write the equation for 
 each reaction. The information thus summarized will be used in Assignment 
 XXI in planning methods of qualitative analysis. 
 Reagents Ca ++ Cu ++ Ag + Zn ++ 
 
 Cl- AgCl, white 
 
 soluble in ? 
 
 OH- Ca(OH) 2 , 
 
 OH~ in excess moderately 
 
 soluble. 
 NH 4 OH 
 
 NH 4 OH in excess 
 CO 3 - CaCO 3 
 
 soluble in 
 
 acids. 
 H 2 S in acid 
 
 solution ? 
 
 S in alkaline 
 
 solution ? 
 
 Ferrocyanide ion 
 Fe(CN). ? 
 
ASSIGNMENT XVIII. 
 HYDROXIDES AND OXIDES. 
 
 Experiment. To solutions containing Cu ++ , Ag + and Zn ++ , in separate 
 test tubes, add a few drops NaOH solution. In the experiments with copper 
 and silver continue to add NaOH solution until the mixture, after shaking, 
 reacts strongly alkaline to litmus. Collect the precipitates on filters. The 
 precipitates are the hydroxides of copper and zinc and the oxide of silver, 
 Ag 2 O. Write equations to show the formation of the solids from the ions. 
 There is an equilibrium in each case. How will each equilibrium be affected 
 by the addition of a strong acid? 
 
 Test the action of nitric acid on the precipitates just obtained. Com- 
 pare your equations with that for the action of a strong acid on solid calcium 
 hydroxide. 
 
 To solutions containing Cu + % Ag + , and Zn ++ , in separate tubes, add a few 
 drops NH 4 OH solution. Do not add excess. 
 
 Experiment. Collect some zinc hydroxide on a filter. Treat a portion 
 with hydrochloric acid, and another portion with sodium hydroxide solution. 
 All hydroxides and oxides of metals react with acids to from water (cf. neu- 
 tralization), and the reaction usually takes place rapidly. Some of them, as 
 in the case of zinc, also react with strong bases to form water, and the 
 reaction again corresponds to neutralization. Compare 
 
 H 2 ZnO 2 (solid) + 2OH- = 2H 2 O + ZnO,~ 
 and H 2 CO, + 2OH~ = 2H 2 O + CO 3 ~. 
 
 ZnO,r~ is zincate ion, and Na 2 ZnO 2 is sodium zincate. What reaction takes 
 place when Zn ++ is treated with excess OH~? 
 
 State if any other hydroxide of a metal which reacts with a strong base 
 in a similar manner has been discussed in the lectures, and write the equation. 
 
 Experiment. Collect some Cu(OH) 2 on a filter, wash it once with water, 
 and heat some of it in a porcelain dish. Cupric oxide has been formed. In 
 order to determine if this is a reversible reaction, allow the dish to cool and 
 add water. Also heat a mixture of cupric hydroxide and water to boiling. 
 What reaction would take place if NaOH solution were added to a solution 
 of copper nitrate at ioo? 
 
 The hydroxides of all metals except the alkali metals are decomposed into 
 the oxide and water vapor when the dry solids are heated strongly. In most 
 cases the oxides will remain unchanged in contact with water, ZnO is an 
 example. In some cases, however, the reaction is reversible. What happens 
 when water is added to calcium oxide? Under what conditions will there 
 be an equilibrium? 
 
 The oxides of the alkali metals, as Na 2 O, react violently with water 
 Write the equation. Solid NaOH melts without decomposition. 
 
 Problems, i. Write equations for the action of HC1 solution on calcium 
 oxide, on ferrous and ferric oxides (FeO and Fe 2 O 3 ), and on ferrous and 
 ferric hydroxides. 
 
 2. How would you prepare solid copper nitrate (a) from cupric hydrox- 
 ide, (b) from cupric sulfate? 
 
 3. Give three examples, as widely different as possible, in which there 
 is an equilibrium involving two solids and a gas. In each case state how the 
 pressure of the gas at equilibrium alters with the temperature. 
 
 32 
 
ASSIGNMENT XIX. 
 COMPLEX IONS CONTAINING AMMONIA. 
 
 The ions of certain metals have the power of forming compounds with 
 ammonia, NH,. The ammonia is supplied by adding NH 4 OH solution. 
 There is an equilibrium between the ion of the metal, NH 3 or NH 4 OH and 
 the complex ion ; and when NH 4 OH is present in excess the concentration of 
 the ion of the metal is often very small. Some examples will be studied in 
 Assignment XIX. 
 
 Experiment. Collect some copper hydroxide on a filter, wash it with a 
 little water, and treat it with 6 normal NH 4 OH solution. Give reasons why 
 the reaction cannot be similar to that of OH" on zinc hydroxide. The deep 
 blue substance is Cu(NH 3 ) 4 ++ . How is the equilibrium between solid 
 Cu(OH). : and its ions affected by the addition of NH 4 OH? Which solution 
 has the smaller concentration of Cu ++ , a saturated solution of Cu(OH) 2 or 
 the solution containing the complex ion? 
 
 From a consideration of the equilibrium 
 
 Cu ++ + 4 NH 4 OH = Cu(NH 3 ) 4 ++ -f 4H 2 O 
 
 predict what will happen when the solution is acidified with nitric acid. Test 
 your conclusions. 
 
 Silver and zinc also form complex ions with amonia, Ag(NH 3 ), + and 
 Zn(NH,) 4 ". 
 
 Experiment. Prepare some silver oxide and zinc hydroxide, and test the 
 action of excess NH.OH solution. 
 
 Predict what will be the effect of treating a silver chloride precipitate 
 with NH 4 OH solution. Test your conclusions by an experiment. Acidify the 
 final solution with nitric acid. 
 
 In addition to the alkali metals and alkaline earth metals there are many 
 metals which do not form complex ions with amonia. Give examples if 
 you can. 
 
 Problems, i. Write equations for the reaction between Cu ++ and 
 NH 4 OH (i) when a few drops NH 4 OH solution are added, and (2) when 
 excess NH 4 OH is added. 
 
 2. Give at least one example of the formation of a complex ion of a 
 metal, other than an ammonia complex. 
 
 ASSIGNMENT XX. 
 CARBONATES, SULFIDES, FERROCYANIDES. 
 
 Experiment. Try the action of Na 2 CO 3 solution on Cu ++ , on Ag+, and on 
 Zn ++ . In each case collect the precipitate on a filter, wash it with water, and 
 test a portion for carbonate. Predict the action of nitric acid solution and of 
 ammonium hydroxide solution on these precipitates, and test your prediction 
 by experiments with HNO 3 and NH 4 OH solutions. 
 
 Predict the action of NH 4 OH solution on calcium carbonate, and test your 
 answer by an experiment. 
 
 Experiment. To solutions containing small amounts of Cu ++ , Ag + , and 
 Zn ++ , in separate experiments, add i cc. 6 normal H 2 SO 4 ; dilute each solution 
 to about 20 cc., and pass H 2 S gas* into it until the liquid is saturated with the 
 gas. To determine this, close the end of the test tube or flask, shake 
 thoroughly, and test the odor cautiously. In the experiment with Zn ++ and 
 H 2 S add 5 cc. 2 normal sodium acetate solution, and again saturate with H 2 S 
 gas. Collect the precipitates on separate filters. 
 
 * Caution. H^S is poisonous. Work with it on porch, and do not breathe it. 
 
 33 
 
Predict the action of 0.3 normal H 2 SO 4 on each precipitate. Treat por- 
 tions of each precipitate with 0.3 normal H 2 SO 4 and with 2 normal H 2 SO 4 . Is 
 zinc sulfide more or less soluble than copper sulfide? Give your reasoning. 
 
 To solutions containing Cu(NH,) 4 + + , Ag(NH 3 ) 2 + , and Zn(NH.,)/ + , in 
 separate experiments, add (NH 4 ) 2 S solution. Repeat the experiments, using 
 H 2 S gas instead of (NH 4 ) 2 S solution. State in each case which has the 
 smaller concentration of the ion of the metal, a solution containing the complex 
 ion, or a solution saturated with the sulfide. State also which is the less 
 soluble, Cu(OH), or CuS, Ag,O or Ag 2 S, Zn(OH) 2 or ZnS. 
 
 Predict the action of H 2 S on mixtures of copper carbonate and water, 
 silver carbonate and water, and zinc carbonate and water. Test your answers 
 by experiments. State. in each case which is the less soluble, the carbonate or 
 the sulfide. 
 
 Test the action of H 2 S on solutions of nitric acid of different concentra- 
 tions, O.i normal, 2.0 normal, and 6 normal. Saturate each solution with H 2 S 
 gas, and heat the mixture almost to boiling. The principal reaction is 
 
 3H 2 S+2NO 3 -+2H + =3S( solid )+2NO+4H 2 O. 
 
 Compare this equation with that for the action of nitric acid on metallic Zn 
 or Cu. 
 
 Collect some CuS on a filter, transfer it to a casserole, add 2 normal HNO-, 
 and boil the mixture. Filter, and test the filtrate with excess NH 4 OH. The 
 residue collected on the filter is sulfur; the dark color is due to a little CuS 
 enclosed in the sulphur. Complete the equation 
 
 3CuS (solid) +2NO 3 -=3S (solid) +2NO. 
 Repeat the experiment, using silver sulfide instead of copper sulfide. 
 
 Experiment. To very dilute solutions containing Cu ++ , Ag + , and Zn ++ , 
 respectively, add a few drops potassium ferrocyanide solution, K 4 Fe(CN) . 
 Repeat these experiments in the presence of (i) a little acetic acid, (2) a little 
 hydrochloric acid. A ferrocyanide frequently furnishes a characteristic final 
 test for an ion, on account of the color ; but ferrocyanides are seldom used in 
 making separations. 
 
 Experiment. Determine which is the most delicate test for copper by 
 preparing very dilute solutions of Cu ++ , e.g. N/iooo, N/io,ooo, etc., and testing 
 separate 10 cc. portions with NH 4 OH, with H 2 S and with K 4 Fe (CN) G . 
 
 Problems, i. Hydrogen sulfide in solution is a weak dibasic acid which, 
 like H 2 CO 3 , ionizes in two stages. Review the neutralization reactions of car- 
 bonic acid (Assignment XIV), and write equations for the neutralization of 
 (i) H 2 S with one mol OH~, (2) H 2 S with excess OH-, and (3) HS~ with 
 OH". Write the equation for the hyd'rolysis of sulfide ion, S~~. 
 
 2. Is it correct to state that all difficulty soluble salts of a weak acid dis- 
 solve when a solution of any strong acid is added? Explain, giving examples. 
 
 3. What is a "basic salt?" Give an example of a basic carbonate. 
 
 Note. Leave your two clean sample bottles at the office, labelled with your name 
 and desk number, and marked No. 3 and No. 4, in order that two unknowns may be ready 
 for you at the next laboratory period. 
 
 ASSIGNMENT XXI. 
 REVIEW. QUALITATIVE ANALYSIS. 
 
 Test what you know of the chemistry of cupric copper, silver and zinc, as 
 studied in the preceding Assignments, by preparing from memory a summary 
 for each metal showing 
 
 I the formulas of the ions (including the complex ions) ; 
 
 II the readily soluble and moderately soluble compounds; 
 
 III the difficulty soluble compounds noting which is the least soluble. 
 
 34 
 
Check each item by reference to your laboratory notes, and correct your 
 mistakes. 
 
 Complete the table given just before Assignment XVIII, and use this table 
 in planning how to analyze the various solutions of salts referred to below. 
 Perform additional experiments whenever you are not certain that your methods 
 are satisfactory. 
 
 Give a method of detecting silver in a solution which may contain Cu ++ , 
 Ca ++ and Na + . How could you separate silver from a solution containing any 
 of these ions? 
 
 Give two methods for each of the following separations : (a) Cu ++ f rom Zn ++ , 
 (b) Cu ++ from Ca ++ . (c) Cu ++ from Na + , (d) Zn + + from Ca ++ , (e) Zn + + f rom Na + . 
 
 A solution contains Ag + , Cu ++ , Zn ++ , Ca ++ , and Na+. Devise a series of 
 operations which would enable you to prepare a compound of each element 
 from a single portion of the solution. Experiment. Test your method with a 
 solution containing all of these ions, and apply additional tests, when necessary, 
 to prove that each of these ions was present in the original solution. Such a 
 series of operations is a scheme of qualitative analysis for these ions. 
 
 A knowledge of the positive ions present in a solution often enables the 
 conclusion to be drawn that certain negative ions cannot be present in appreci- 
 able amounts. What negative ions need not be tested for when the following 
 are present at moderate concentrations? (a) H + , (b) Ca ++ , (c) Zn ++ , (d) Cu ++ , 
 (e) Ag+. 
 
 Experiment. Test for nitrate ion. NO 3 ~ To about 2 cc. of the solution 
 to be tested add excess ferrous ammonium sulfate (FeSO 4 :(NH 4 ), J SO. 1 ) solu- 
 tion, or add the solid salt, filter if there is a precipitate, hold the test tube in a 
 slanting position and pour carefully down the side (from a small beaker) 3 to 
 5 cc. concentrated sulfuric acid. The concentrated acid sinks to the bottom of 
 the test tube and a dark brown ring forms on its surface when nitrate is 
 present. This brown substance decomposes when the solution is heated. 
 
 Analysis 3 and 4. Test for 
 
 H + , Ag + , Cu ++ , Zn+ + , Ca+ + , Na+, K + , NH 4 + , 
 OH-, C1-, N0 3 -, SO 4 - CO 3 ~, S-. 
 
 When an ion is found to be present, try to decide whether it is present in a 
 large amount, in a small amount, or as a mere trace. 
 
 In the second term A. A. Noyes' "Qualitative Analysis," new edition, 
 1914, will be used; it is for sale at the Co-operative Store, price $1.50. The 
 new edition is the same as that used in 1914-15, but differs greatly from that 
 used in 1913-14. 
 
 35 
 
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