D UC-NRLF 277 LABORATORY EXERCISE S TO ACCOMPANY STORER AND LINGSM'S CHEMISTRY LIBRARY OF THE UNIVERSITY OF CALIFORNIA GIETT i LABORATORY EXERCISES IN GENERAL CHEMISTRY COMPILED FROM VARIOUS SOURCES BY G. W. SHAW, A.M. FORMERLY PROFESSOR OF CHEMISTRY AT OREGON STATE AGRICULTURAL COLLEGE FOR USE IN CONNECTION WITH STOKER AND LINDSAY'S MANUAL OF CHEMISTRY \ S rf A f? OK THE UNIVERSITY OF lUFOR]^ NEW YORK : CINCINNATI : CHICAGO AMERICAN BOOK COMPANY COPYRIGHT, 1901, BY CK W. SHAW. ENTERED AT STATIONERS' HALL, LONDON. SHAW. LAB. EX. W. P. i PREFACE. THIS pamphlet claims to present nothing that is new and untried. Beginners in chemistry need such experiments as give certain and well-defined results. The pamphlet has been prepared to use in the laboratory in connection with Storer & Lindsay's " Manual of Chemistry " in the class room, hence the greater number of the experiments have been drawn from that work, yet many from other sources have been included. A large number of manuals have been examined in the prepara- tion of the pamphlet, but the compiler is particularly indebted to those of Williams and Remsen. The raison d'&tre for the book is simply that of facilitating work in the laboratory by having before the student only such matter as is necessary for the work in hand. It has been the endeavor of the compiler to include only such experiments as are truly instructive and well illustrate the subject, and which lead up easily and naturally to a knowledge of the science. The book is to be used in a laboratory in which an instructor is always present with the class to furnish such individual direction as may be necessary. Most of the experiments should be performed on the lecture table before they are attempted by the student. Any text-book may be consulted outside the laboratory ; in the laboratory itself no other book than this is allowed. In 3 J 07459 4 PREFACE. some cases a sketch of the necessary apparatus, or an example of the apparatus itself, is placed in the laboratory, where it can be seen by the students. The last twenty minutes of the laboratory period should be taken to question the class rapidly on the work of the day. The books should be inspected after each experiment has been written out. OF THE UNIVERSITY TO THE STUDENT. 1. Unless special directions to the contrary are given no text-book will be allowed in the laboratory except this pamphlet. 2. Provide yourself with a working apron to protect your clothing ; also with towel or cloth to clean the desk. 3. Neatness is essential to success in chemical work. Both the desk and apparatus must be kept deem at all times. 4. Pupils are held responsible for apparatus, and must replace anything lost or broken. 5. Never mix chemicals or reagents except as directed. 6. In experimenting have all apparatus neatly arranged. Have every stopper and connection fit tightly. 7. Have flasks and tubes perfectly dry on the outside before applying heat. 8. Keagents for general use must not be taken to the stu- dent's desk, but used at the side table. Any excess of a reagent must not be poured back from a test tube or beaker into the reagent bottle, nor should a .stirring rod be dipped into a reagent bottle. 9. In experimenting follow directions as closely as pos- sible. Kead an experiment through before performing any part of it. Ask an explanation of anything you do not understand. 10. Observe very carefully everything that takes place ; and endeavor to distinguish essential from non-essential phe- nomena; express in writing the results of your observation and the conclusions as to facts taught by the experiment. 6 6 TO THE STUDENT. 11. A student's standing is largely determined by the quality of his laboratory work and notebooks. The notes must be written clearly and distinctly. Neatness will be in- sisted upon. Do not crowd the notes, but leave room for remarks, or corrections. In general answer the following questions : (a) What is the object of the experiment ? (b) What materials did you use ? (c) What apparatus did you use ? (d) What did you do ? (e) What did you observe ? (/) What are your conclusions ? 12. At the end of the term no allowance is made for appa- ratus that is not clean and in proper condition to be served out to other students. I have read the above directions carefully, and endeavored to understand them. Signed LABORATORY EXERCISES EXERCISE I. PRELIMINARY. 1 to 5.1 1. Ascertain the number of your desk. 2. Obtain key to desk at the supply room. 3. Examine and take inventory of contents of desk. 4. Take receipt to the supply room. Experiment 1. Measure in a graduate 10 cc. of water, then pour it into a test tube. Note what proportion of the tube is filled. Pour out the water, then pour into the tube as near the same quantity of water as possible, estimating by the eye alone. Verify your estimate by measuring in a graduate. Repeat till you can estimate closely. Hereafter, unless great accuracy is required, estimate volumes without measuring. Supplementary. Learn the metric table for length, weight, and vol- ume. How many cubic centimeters in a liter ? What is the weight in grams of a liter of water? Mercury weighs 13,6 times as much as water. What will 5| cc. of mercury weigh ? A rectangular block of marble is 3x4x2 cm., and has a sp. gr. 2.5. What does it weigh in grams? 2. (a) Try to dissolve a small quantity of sulphur in carbon bisul- phide. Try to pick up small pieces of iron with a magnet. Try the effect of dilute hydrochloric acid on sulphur. Is any gas given off ? Try the same acid on iron particles. (&) Mix thoroughly small quantities of sulphur and iron. Is a new substance formed? Is the sulphur still sulphur? Is the iron still iron? Weigh out 3 g. of sulphur and 6 g. of iron ; mix thoroughly, and put it into a tube of hard glass closed at one end. Hold the tube by 1 Lesson in the " Manual of Chemistry " to precede the exercise. 7 8 LABORATORY EXERCISES. the open end by means of a strip of paper folded several times, and heat the mixture over a burner. When no further change takes place allow the tube to cool ; break it, and examine the contents for sulphur and for iron. 3. Heat a small piece of lead foil on the lid of a porcelain crucible as long as any change occurs. Warm slowly, as sudden heat will cause the porcelain to break. While this is going on, perform the following : 4. Fit to any small flask or bottle a perforated cork to which has been adapted a short piece of glass tubing. Over the end of this glass tube slip a short piece of caoutchouc tubing. Suck part of the air out of the flask, and then nip the caoutchouc tubing with the thumb and finger so that no air shall reenter. Immerse the neck of the flask in a basin of water and release the tubing. Explain all results. 5. Adapt the throat of a funnel to a perforated cork which fits tightly into the neck of a bottle or flask, and then fill the funnel with water. Why does the water not enter the bottle ? Loosen the cork so that the mouth of the bottle is not tightly closed. Explain the action. 6. Burn some magnesium wire, observe carefully the result of the combustion, and describe the product. EXERCISE H. CHEMICAL AND PHYSICAL CHANGES. 5 to .12. 7. Upon a small piece of quicklime put a few drops of water. Carefully describe all changes which take place. 8. Examine a piece of marble. Notice whether it is hard or soft. From a piece of glass tubing of about | inches internal diameter cut off a piece about 4 inches long by making a mark across it with a triangular file, and then seizing it with both hands, one on each side of the mark, pulling, and at the same time pressing slightly as if to break it. Clean and dry it, and hold one end in the flame of a labora- tory burner until it melts together. During the melting twirl the tube constantly between the finger and thumb so that the heat may act uniformly upon it. Heat a piece of the marble in this tube. Does it change in any way ? Will it dissolve in water ? Treat a small piece with dilute hydrochloric acid. What takes place? LABORATORY EXERCISES. 9 After the action has continued for about half a minute insert a lighted match into the upper part of the tube. Does the match con- tinue to burn ? Does the substance in the tube burn ? Is the invisible substance in the upper part of the tube ordinary air ? How do you know ? Does the solid substance disappear ? In order to tell whether it has been changed chemically the hydro- chloric acid must be driven off. This can be done by boiling, when it passes off in the form of vapor, just as water does, and then whatever is in solution will remain behind. For this purpose put the solution in a small, clean porcelain evaporating dish, and heat slowly till the liquid has disappeared. After the liquid has evaporated and the substance in the evaporat- ing dish is dry, examine it and carefully compare its properties with those of the substance which was put into the test tube. Is it the same substance ? Is it hard or soft ? Does it change when heated in a tube? Is there an appearance of bubbling when hydrochloric acid is poured on it ? Does it dissolve in water ? Does it change when allowed to lie in contact with the air ? In order to learn whether a substance is soluble in water proceed as follows : Put a piece about the size of a pea in a test tube with pure water. Thoroughly shake, and .then, as heating usually aids solution, boil. Now pour off a few drops of the liquid on a watch glass, and, by gently heating, cause the water to pass off as steam. If there is any- thing solid in solution, it will be left on the platinum foil or watch 9. Examine crystals of salt. Note whether they are hard or soft. Put into a test tube about 5 g. of the salt and just cover with water. Heat till the salt dissolves. Taste a drop of the liquid. Have the salt particles been divided by dissolving ? How do you know ? Arrange a filter paper (ask the instructor how) and filter the solu- tion. Taste a drop of the filtrate. Has the salt passed through the paper? Dilute a portion of the liquid with an equal volume of \\ T ater, and shake vigorously. Have the salt particles been further divided ? Evaporate the remainder of the liquid as directed in Exp. 8. Com- pare the residue with the original salt. 10. Add a little hydrochloric acid, HC1, to about 5 cc. of a solution of lead nitrate, Pb(N0 3 ) 2 . The hydrochloric acid and lead nitrate 10 LABORATORY EXERCISES. have been changed to lead chloride and nitric acid. The lead chloride is insoluble and is therefore thrown down (precipitated). Indicate which of the above experiments represent physical and which chemical changes. EXERCISE III. AIR AND OXYGEN. 12 to 19. 11. In porcelain crucibles carefully ignite weighed amounts of (a) zinc dust, stirring occasionally; (b) copper filings. 1 Weigh again. Was anything given off or absorbed ? Could it have come from the crucible ? From the gas ? From whence then ? Red oxide of mer- cury was thus obtained by early experimenters. It is mercury rust (just as the above are zinc rust and copper rust, respectively). 12. Heat a gram of red oxide of mercury in an ignition tube made as in Exp. 8, holding the tube nearly horizontally. Weigh the tube before and after the heating. What evidence is there of any change ? What is deposited on the tube ? Whence did it come ? During the heating insert into the tube a splinter of wood with a spark on the end. What follows ? Take it out and put it back a few times. Is there any difference between the character of the burning in the tube and out of it ? What difference? What causes the difference? How do you know that the red substance which you put into the tube has been changed ? How could you collect the gas given off ? Is the change above an example of analysis or of synthesis ? How do you explain the loss in weight? 13. How to collect a gas over water. Gases insoluble in water can be collected over that liquid. To collect gas in this way, fill with water the vessel to be filled with gas. Place over its mouth a glass plate, and, holding the plate firmly over the mouth of the vessel invert it with the mouth under water, in the pneumatic trough. Does the vessel remain filled with water ? Why ? Place the end of a glass tube under the inverted cylinder, and blow gently through the tube. Explain what happens. i These should have been freed from oil by means of ether, and carefully dried. LABORATORY EXERCISES. 11 14. Arrange an apparatus as shown in the model. In the flask (or in a large test tube) put 4 to 5 g. potassium chlorate, and gently heat by means of a lamp. When the gas comes off freely bring an inverted cylinder filled with water over the end of the tube, and let the bubbles of gas rise in the cylinder. Confine the gas by placing a glass plate over the mouth of the vessel and inverting it. Insert into it a stick with a spark on its end. AVhat takes place? Is the gas contained in the vessel ordinary air? What caused the chemical change in this case ? In what respects is this chemical change like that in the last experiment? Which of the above illustrates the combustion of fuel ? 15. Make a deflagrating spoon by hollowing out the end of a piece of crayon and attaching it to a wire. In the spoon thus prepared place a small piece of roll brimstone (or a little sulphur) and allow it to burn in the air. Notice the odor of the fumes. Now set fire to another small portion and introduce it in the spoon into one of the vessels containing oxygen. Notice the odor of the fumes given off. Do they appear to be the same as those given off when the burning takes place in the air? Try a bit of phosphorus in a similar manner. Compare the pro- ducts formed when carbon, sulphur, phosphorus, etc., are burned in air and in oxygen. EXERCISE IV. NITROGEN. 16. Float a small evaporating dish on water contained in a pneu- matic trough. Put into the dish a small piece of phosphorus, 1 and set fire to it by means of a hot wire. Quickly place a jar over it on a support which will prevent the jar from sinking more than an inch or two in water. Why is the air at first forced out of the vessel ? Why does the air afterward rise in the vessel ? After the burning has stopped, and the vessel has cooled down, about what proportion of the air is left in the vessel? Cover the mouth of the jar with a glass plate and turn it mouth upward. Try the effect of introducing one after the other several burning bodies into the gas, as, for example, a piece of sulphur, etc. Explain all that you have seen. 17. 2 Put in a side-neck test tube about 2 g. of ammonium chlo- ride, NH 4 C1., 3 g. of sodium nitrite, NaN0 2 , and moisten with a few 1 Phosphorus should always be cut under water, and he handled with for- ceps, never with the fingers. 2 To the Instructor. The following may be substituted for 16. 12 LABORATORY EXERCISES. drops of water. Apply gentle heat and collect the gas over water. Try the same combustion experiments as indicated in Exp. 16. Nitrogen may also be generated by passing air over red hot copper, Cu. What compound would be formed f Compare the properties of oxygen and nitrogen. From the percentage of nitrogen in the air about 1 per cent, must be subtracted for the recently discovered element argon, A. 17 (a). Into a flask fitted with delivery tube and thistle funnel put 50 cc. of reagent ammonia. Rub about 20 g. of bleaching powder into a thin paste with water, and add it gradually through the funnel while the flask is being gently heated. Collect the gas over water. EXERCISE V. EFFECT OF TEMPERATURE ON VOLUME OF GASES. 18. Fit a small flask with a glass tube about 20 cm. long by means of a perforated rubber stopper. Support the tube or flask in a vertical position, flask uppermost, in a vessel of colored water. Apply heat to the flask by means of the burner. Note the result. What escapes through the tube? Why? CAUTION. Do not continue the heating too long or the flask may be broken by the inrush of water on the removal of the heat. Why does the water rush in after the removal of the heat ? State the effect of heat on the volume of a gas. Suppose the flask to have been tightly stoppered, and heat to have been applied in a like manner, how would the pressure under which the gas existed have been affected? COEFFICIENT OF EXPANSION OF GASES. 19. For the Instructor. Into a glass tube about 20 cm. long and of 1 mm. bore, introduce for an index about 5 mm. of mercury by placing one end of the tube in a bottle of mercury and holding the finger on the stopper end of the tube as it is removed from the bottle. In- cline the tube so as to work the mercury column to about the center of the tube. Now, holding the tube in a horizontal position, close one end of it by holding in a Bunsen flame. Fasten a chemical thermom- eter to the tube by means of wire or string so it may be used as one of the divisions for measuring the length of the column of confined air. Have ready a long shallow dish containing ice water. Place the apparatus in the dish of ice water in as near a horizontal position as LABORATORY EXERCISES. 13 possible, and determine the length of the inclosed air column for 10 C., 20 C., and 30 C., heating the tube by the gradual addition of hot water to the water in the dish. From the data obtained compute the expansion per degree of the air per unit of volume as measured at C., between C. and 10 C., 10 C. and 20 C., etc. If the work is carefully done the result should be close to .00366, or 5 } y of the bulk per 1 C. At what temperature would 1 cc. of gas become 2 cc., the pressure being constant ? What would be the volume of 1 cc. measured at C., when cooled to 273 C.? What temperature is called absolute zero? NOTE. The student must understand that these quantitative experiments require the utmost care, and even then, with the crude apparatus used, and the limited experience in doing such work, only approximate results can be expected. The above experiments are intended to illustrate the fact that gases expand with the increase of heat, and vice versa, and therefore vary in the pressure they exert; and also the LAW OF CHARLES: "All true gases, when heated, expand 5^3 of their volume, measured at OP C.,for each increase oflC." Example i. Twenty cubic centime'ters of hydrogen were measured at 15 C. and heated to 35 C. What was the new volume ? Example 2. Five hundred cubic centimeters of nitrogen measured at 27 C. would become how many when cooled to 10 C.? EXERCISE VI. EFFECT OF PRESSURE ON VOLUME OF GASES. 20. For the Instructor. Procure a tube shaped like the letter J, the short arm being closed and about 25 cm. long, and the long arm at least 85 cm. in length and open at the top. There will be less liability to accident and consequent loss of mercury if the tube be fastened in an upright position, either by attaching it to a board set in a block for a base, or by tying it securely to the standard of a ring stand. Divide the short arm into four equal parts, indicating the divisions by means of rubber bands or gummed paper ; also make a mark near the base of the long arm exactly opposite the lowest mark on the short arm, which must be above the bend of the tube. Now pour mercury into the tube through a small funnel till the surfaces in the two arms stand at the lowest division in each arm. We have now entrapped some air, and the tension (pressure) in the short arm is the same as in the long arm. How is this indicated ? Under how much pressure per square inch is it ? Now pour mercury into the long arm, carefully inclining the tube so as to include as little air as possible, till the space 14 LABORATORY EXERCISES. in the short arm has been reduced one half. Measure the height of the column of mercury in the long arm above the mark near the base, and compare it with the reading of the barometer at the time of the experiment. Remembering what was the pressure of the confined air in the beginning, what is its tension now ? Under how many atmos- pheres' pressure is it? How is this shown? How does its former volume compare with that occupied now? Doubling the pressure has how affected the volume occupied ? How has it affected the tension of the gas? If the length of the tube will allow, add as much more mercury to the long arm. Under how much pressure is the confined air now? How does its volume compare with the original volume? Under about how much pressure per square inch is it ? EFFECT OF DECREASING THE PRESSURE. 21. In the previous experiment the pressure on the confined air was increased. Let us now reverse the condition and ascertain the effect of diminishing the pressure. Fill a barometer tube to within about 10 cm. of the top. Place the forefinger over the open end, thus inclosing a certain volume of air. Under what tension is the inclosed air ? Now invert the tube and allow the air to rise to the top. Keeping the finger over the end of the tube, measure the length of the air column. Place the mouth of the tube under mercury contained in a porcelain mortar and remove the finger, and the mercury column will fall, for the atmosphere outside cannot support the atmospheric pressure of the column plus the weight of the mercury column. As the mercury falls the imprisoned air expands and presses less on the mercury column, and a point is soon reached at which the pressure outside and inside the tube is equal. The inclosed air is under the atmospheric pressure (how many cm. of mercury ?) less the pressure due to the column of mercury in the tube. Call the original volume of inclosed air V l and the final volume V 2 ; the original tension P l and the final tension P 2 ; the results should give very close to the following equation, provided the work has been carefully done. Vj x P! = V 2 x P 2 . Repeat the experiment with a new volume of inclosed air, and see if the results agree. LABORATORY EXERCISES. 15 Express as a law the relation between volume and pressure as exhibited by these experiments. This is known as BOYLE'S LAW. The law is not rigorously correct for all gases, the variations depend- ing on the kind of gas used. It is sufficiently exact for all practical purposes, however. The law does not hold for gases when the measurement is taken too near their point of liquefaction. Example i. Twenty cubic centimeters of hydrogen were measured at 760 mm. pressure and afterward the pressure changed to 768 mm. What volume did the gas then occupy ? Example 2. Five hundred cubic centimeters of nitrogen measured at 768 mm. would become how many cubic centimeters at 755 mm. ? EXERCISE VII. WEIGHT AND DENSITY OP AIR. 22. Fit a stout Florence flask of about a liter capacity with a one- hole rubber stopper through which passes a tightly fitting glass tube carrying a stiff pinchcock. Make the apparatus air-tight by smear- ing the points with vaseline, if necessary. Be sure the flask is clean and dry. First, weigh the apparatus full of air and record the weight. Second, open the pinchcock and by means of an air pump exhaust the air from the flask as thoroughly as possible. Close the pinch- cock, reweigh the apparatus, and note the loss in weight. Third, measure the volume of air that was removed, as follows : Place the rubber tube bearing the pinchcock well beneath the surface of water in a pneumatic trough. Open the pinchcock, keeping the tube well under water all the time. The water will rush into the flask. Why ? Make the level of the water inside and outside the flask equal by lowering or raising the flask in the water. Close the pinchcock, remove the apparatus from the water, and make it dry. Ascer- tain the number of cubic centimeters of water in the flask by weigh- ing on a balance, remembering that 1 g. equals 1 cc. of water. From the data thus secured compute the weight of one cubic centi- meter of air under the conditions of the experiment, or the density of air. Define density. REMARK. The term density is often used as synonymous with specific gravity, but the latter expresses more properly the number of times a given volume of a certain substance is heavier than some other substance taken as a standard, while density refers to the weight of a unit volume. 16 LABORATORY EXERCISES. Referring to 19, 20, and 21, how would the density of air have been affected had you made the determination at C, and 760 mm. pressure, which are the standard conditions ? EXERCISE VIII. WATER. 34 to 37. 23. Hold a dry cold beaker for an instant over a Bunsen burner flame. Is any water deposited on it ? In a dry test tube heat gently a small piece of wood. What evi- dence do you obtain that water is given off? Try the same with a little sugar in a test tube. 24. For the Instructor. 1 Roll a piece of metallic sodium in a piece of wire gauze and by means of a pair of tongs thrust the cage under the mouth of an inverted receiver containing water. After the action has ceased close the mouth of the receiver, and turn it uppermost. Remove the plate and apply a lighted taper to the gas. Does it act like oxygen ? How does it differ ? 25. For the Instructor. 1 In a retort of convenient size distil water before the class. Call attention to the colorless nature of the steam. Place a few drops of the distilled water on a platinum foil and evapor- ate. Then do the same with a few drops of water remaining in the retort. *What effect has distillation on the purity of water ? Does well water contain solids in solution ? Why has distilled water not an agreeable taste? 26. Dissolve 2 or 3 g. of common salt in distilled water. Evapor- ate the solution slowly to dryness and compare the substance ob- tained with the original salt as to appearance, taste, and crystalline form. Dissolve 5 g. of sodium carbonate, Na 2 C0 3 , in dilute hydrochloric acid, HC1. Evaporate the solution as before, and compare the resi- due with the original salt as to appearance and taste, and by treating with hydrochloric acid. Explain the difference between physical and chemical solution. 27. Dissolve some ordinary alum in water (f ounce alum to 100 cc. water) by the aid of heat. Filter through a plaited filter and allow the filtered solution to cool. What takes place? 1 These may be performed in the lecture room. LABORATORY EXERCISES. 17 Which will hold in solution the greater amount of substance hot water or cold water ? Define crystallization. 28. Pour off the above liquid and place a few of the crystals on a piece of dry filter paper. After the water is all absorbed from them and they appear dry, put them in a dry test tube and heat gently. What evidence have you that water is contained in the crystals ? 29. 1 Some bodies give up their water of crystallization simply on contact with the air. Such bodies are said to be efflorescent. Put some crystals of sodium sulphate, Na 2 S0 4 10 H 2 0, on a watch crystal and expose them for an hour or more to the air of the room. Try also calcium chloride. Substances acting like the latter are said to deliquesce. Define efflorescence ; also deliquescence. 30. Take a glass tube 3 or 4 cm. in diameter, and close one end with a plug of plaster of Paris 1 or 2 cm. thick. Set the tube aside to dry until the next exercise. EXERCISE IX. ANALYSIS OF WATER. 31. For the Instructor. Fill a water decomposition apparatus with water acidulated with one-tenth its volume of sulphuric acid. Connect the platinum electrodes of the apparatus with the wire from a Bunsen battery. Start the current, and bubbles will collect in each of the two arms. Compare the volume of the two gases which collect in the two tubes in a given time. When one of the tubes is full hold a lighted match above the tube, open the stopcock, and ignite the gas. Is the gas air? How do you know? Does the gas burn ? Immerse a glowing splinter in the gas in the other tube. Is the action the same as in the former case? Is the gas in this tube air? Is it steam ? Does it fume ? How does it act toward the splinter? In this experiment the water has been separated into its constituent parts analyzed. Define analysis. The reverse of this would be synthesis. LAW OF DEFINITE PROPORTIONS. 32. For the Instructor. Fill a eudiometer tube with mercury and invert it over a dish of mercury. Now from a gasometer admit to the tube, which should be held in a slanting position, 10 cc. of hydrogen 1 This experiment should be arranged at the beginning of the exercise. LAB. EX. 2 18 LABOEATOEY EXBECISES. which is first passed through a wash bottle containing a solution of 2 g. of caustic potash in 10 cc. of water, and then through a second wash bottle containing strong sulphuric acid. Bring the tube to an upright position and secure the following data: (a) volume of gas, (6) temperature, (c) reading of barometer, (d) length in millimeters of the mercury column in the tube. Again slant the tube, and admit in a similar manner about 8 cc. of oxygen from the gasometer. Now bring the tube to an upright position, and again make readings as above. Press the eudiometer tightly against a leather or rubber washer on the bottom of the trough and clamp tightly in position. From an induction coil pass a spark through the gases by connecting wires from the coil to the electrodes of the tube. After the tube has cooled raise slightly and make the same readings as in the two instances above. Calculate all the volumes to the standard conditions as given under Remarks, Exp. 22. What became of the extra volume of oxygen? Test the gas remaining in the tube for oxygen, as under Exp. 12. Repeat the experiment, using an excess of hydrogen. 1 In what proportion did the two gases unite each time ? What lesson may be learned from these two experiments as to the proportion in which elements unite volurnetrically ? EXERCISE X. 33. Into a small tared beaker weigh just 5 g. of sodium carbonate crystals (they must not have effloresced) and dissolve them in water. Now add hydrochloric acid little by little as long as any effervescence takes place. Now evaporate the water and ascertain the weight of the salt remaining. Care must be exercised toward the last of the drying lest loss occur from spattering. Now repeat in exactly the same manner, except that after all effervescence has ceased an excess of the acid is added. Compare the weight of salt obtained in this case with the first. Did any different amount of the acid unite in the second case from that in the first ? What is effervescence ? . State your observation in the form of a law. 1 These gases should be admitted through a tube drawn out to a small jet. LABORATORY EXERCISES. 19 34. Conservation of Matter. One of the laws which the studeut must have impressed upon his mind early in his chemical study is that in the various changes which substances may undergo, nothing is lost. No matter is destroyed, nor is any made. This law is the basis of all physical science. This fact may be illustrated as follows : Make a saturated solution of calcium chloride by adding the salt to about 20 cc. of water as long as it will dissolve. In a like manner prepare an equal amount of a saturated solution of sodium sulphate. Fill a test tube a little less than half full of the first solution, and in another tube place an equal amount of the second solution. Place the two tubes in a beaker to keep them from overturning, and determine the combined weight. Pour one solution into the other, and, after shaking and observing the change, again ascertain the weight. Has weight either been gained or lost in this operation? State in the form of a law the fact here illustrated. EXERCISE XI. MULTIPLE PROPORTIONS. 35. For the Instructor. In the blue flame heat a small clean porcelain crucible supported on a pipestone triangle. Allow it to cool in a desic- cator containing either sulphuric acid, or some pieces of calcium chloride. When perfectly cold weigh the crucible on a delicate balance, and record the weighing. Place in the crucible a layer of dry copper oxide. Again obtain the weight. After placing the crucible again on the pipestone support, cover it with a cover having in the middle a small hole. Connect a bent porcelain tube with a Kipp's hydrogen gen- erator (or a gasometer), and turn on a current of hydrogen, directing it into the crucible by means of the bent porcelain tube through the hole in the cover. Heat the crucible with this current of hydro- gen passing in for about ten minutes. Allow the crucible to cool in the atmosphere of hydrogen, after which weigh. Take the difference in the two weighings as the weight of oxygen given up by the given weight of the copper oxide used. Calculate the weight of copper combined with eight parts of oxygen. Repeat the experiment, using red cuprous oxide in place of the black oxide, and calculate in the same manner. Compare the two proportions with each other. What does the experiment show? State the law of multiple proportions. Find in the text-book other examples of multiple proportions. 20 LABORATORY EXERCISES. EXERCISE XII. HYDROGEN. 37 to 41. The instructor will have prepared a gasometer filled with the gets from which it can be obtained by the students for studying its properties in Exp. 38. 36. Fill the plugged tube prepared in Exp. 30 with hydrogen and set it upright in a glass of water. Examine it from time to time during the laboratory period. Describe the result. 37. Into a cylinder or test tube put a few pieces of granulated zinc, and pour upon it enough ordinary hydrochloric acid to cover it. After the action has 'continued for a minute or two, apply a lighted match to the mouth of the vessel. Describe in full. Perform the same experiments, using sulphuric acid which has been diluted with four times its volume of water. 1 What is the result? Try iron and sulphuric acid. From what does the hydrogen come ? Ascertain if same result is obtained by substituting zinc oxide for zinc. 38. Carefully lift from the water pan a bottle completely full of hydrogen. Slowly carry the bottle, the mouth held downward, to a burning splinter of wood, and depress the bottle over this flame. After observing what happens, withdraw the taper slowly. Does the hydrogen burn at the surface or at the end of the splinter ? Does hydrogen support combustion ? What gathers on the inside of the jar when the hydrogen burns ? 1 To dilute ordinary concentrated sulphuric acid with water, the acid should be poured slowly into the water while the mixture is constantly stirred. If the water is poured into the acid, the heat evolved at the places where the two liquids come in contact with each other may be so great as to convert the water into steam and cause the strong acid to spatter. In experimenting with hydrogen, no light should ever be brought into con- tact with the contents of the bottle in which it is generated, or with any large quantity of the gas, until the purity of the sample, or rather its nonexplosive character, has been demonstrated by applying to a very small volume of the gas the test above described, LABORATORY EXERCISES. 21 EXERCISE XIII. COMPOUNDS OF OXYGEN AND NITROGEN. 57 to 64. 39. Place a small quantity of ammonium nitrate, NH 4 N0 3 , in a test tube, and heat. Hold a piece of cool glass near the mouth of the tube, and note what collects upon it. What is the first change that takes place ? What is the next ? Complete the equation, NH 4 N0 3 = 2H 2 +? 40. Collect over water a jar of nitrogen monoxid, N 2 0, made by heating about 5 g. of ammonium nitrate in a side-neck tube. Do not heat higher than is necessary to secure a regular evolution of the gas. Observe its taste, odor, and effect upon a burning stick. Explain the action of the burning gas toward the burning stick. 41. NITRIC OXIDE. For the Instructor. Fit a flask with both de- livery and safety tubes. Into this flask put pieces of copper turnings. Cover them with water. Now slowly add ordinary concentrated nitric acid. When enough acid has been added gas will be given off. If the acid is added quickly the evolution of gas takes place too rapidly, so that the liquid is forced out of the flask through the funnel tube. This can be avoided by not being in a hurry. Do not inhale the gas. Perform the experiments with nitric oxide where there is a good draught. What is the color of the gas in the flask at first ? What is it after the action has continued for a short time ? Collect over water two or three vessels full. Balance the following equation 3Cu + 8HN0 3 = 3Cu(N0 3 ) 2 + H 2 0+N0. 42. Turn one of the cylinders of nitric oxide with the mouth upward and uncover it. What takes place ? What element in the air is most likely to be the cause of this change ? Does the gas after exposure to the air resemble that in the generating flask at first ? Explain the presence of the colored gas at the beginning of Exp. 41 and the fact that it finally disappeared. How many oxides of nitrogen are described in your text-book ? Make a table of them, showing the relation N to 0, by weight. What fundamental law of chemical action may be derived from a consideration of these compounds ? 22 LABORATORY EXERCISES. EXERCISE XIV. NITRIC ACID. 64 to 70. 43. Instructor prepare nitric acid, HN0 3 , as per S. & L., p. 55, Exp. 33. The retort used must be a glass-stoppered one. Using the acid thus prepared (if sufficient for the entire class has not been prepared, then a portion of the class may demonstrate to the others), let the students ascertain the properties by the following experiments : 44. To 1 cc. of nitric acid of Exp. 43 add 10 cc. of water. Touch a drop of the mixture to the tongue. Dip a piece of blue litmus paper into the liquid. 45. With a stirring rod place a drop of the strong acid on the finger nail, and after a moment wash it off. Put a few pieces of white wool or worsted into a few drops of the strong acid in an evaporating dish and warm gently. What is the effect of the acid on organic substances ? 46. Try the effect of the acid on a bit of copper. Place in each of two test tubes a piece of zinc, Zn. To one add dilute hydrochloric acid and to the other a few drops of strong nitric acid diluted with an equal volume of water. What difference is noticed in the action ? Try the same with iron filings. How does nitric acid act upon metals ? 47. When nitric acid acts on metals nitrates are formed : e.g. nitric acid acting on copper. In such cases the hydrogen, H, of the acid is replaced by the metal. What becomes of the hydrogen ? These nitrates are good oxidizing agents. Melt in a tube a few pieces of potassium nitrate, and then drop in a bit of charcoal and heat. Repeat, using a small piece of roll sulphur instead of the charcoal. EXERCISE XV. BASES, SALTS. 70 to 72. 48. Test the acids that you find in the laboratory as to their effect upon litmus paper, and also as to their taste after diluting them with ten or fifteen times their volume of water. How do they compare in their volume action ? What does this experiment teach concerning acids? LABORATORY EXERCISES. 23 49. Into about 20 cc. of water put a few drops of potassium hydrate solution, KOH. Rub a little of the KOH solution between the fingers. Cautiously taste of the diluted solution. Immerse a piece of red litmus paper in it. Try the same with sodium hydrate, NaOH, and ammonium hydrate, NH 4 OH. What does this experiment teach concerning bases ? How do the bases act toward litmus as compared with the acids ? 50. To a solution of caustic soda, NaOH, add dilute hydrochloric acid slowly, examining the solution from time to time by means of a piece of paper colored blue with litmus. As long as the solution is alkaline it will cause no change in the color of the paper. The instant it passes the point of neutralization it changes the color of the paper red. When this point is reached, evaporate the water on a water bath to complete dryness, and see what is left. Taste the substance. Has it an acid taste? Does it suggest any familiar substance ? Is it an acid, an alkali, or is it neutral ? Treat a little of the material with sulphuric acid, H 2 S0 4 , and note its action. Treat a little common salt, NaCl, in a similar manner. How do the two substances compare in action ? Write the equation showing the formation of this salt. 51. Test with litmus paper the reaction of other salts in solution ; for example, ammonium sulphate, (NH 4 ) 2 S0 4 , and potassium chloride, KC1. What is the usual action of salts toward litmus ? Try, however, solutions of sodium carbonate, Na 2 C0 3 , copper sul- phate, CuS0 4 , sodium bicarbonate, NaHC0 3 , sodium bisulphate, NaHS0 4 . EXERCISE XVI. AMMONIA. 72 to 78. 52. Take in one hand a little dry quicklime, CaO, and in the other an equal bulk of pulverized ammonium chloride, NH 4 C1. Note that neither substance has an odor. Now rub the two together between the hands, and carefully note the odor of the gas given off. 53. To a solution of ammonium chloride in a test tube add a few drops of potassium hydroxide, KOH. Warm gently, and note the odor of the fumes, as well as their action toward litmus. Also do the same with a solution of ammonium nitrate, NH 4 N0 3 . What result ? Write equations for each reaction. 24 LABORATORY EXERCISES. Moisten a stirring rod with hydrochloric acid, and hold it in the escaping ammonia fumes. Try the same with nitric acid. The two gases unite in each case to form a solid. Write equations to show the union. What two classes of compounds are used to make ammonia ? 54. Fit a right-angled bend to a flask by means of a perforated stopper. Connect to this bend a second tube of the same shape, direct- ing the free end downward. To the lower end of this tube adjust a glass funnel, inverted. Allow the mouth of the funnel to dip under water held in a beaker. Into the flask put two parts of ammonium chloride to one of quicklime. Moisten slightly and apply gentle heat. After the gas has escaped in the water for a few minutes disconnect the funnel and direct the gas upward, by turning the tube, into a cylinder held over the end of the delivery tube. Into the gas thus collected insert a burning stick. Does the gas burn ? Does it support combustion ? Is the gas soluble in water ? Test with red litmus paper the water into which the gas has been passed. Is it acid or is it alkaline ? What is the liquid? What is the relation between ammonia and ammonium hydroxide? Make in your notebook a tabular statement of the physical and chemical properties of ammonia. 55. Place in a beaker a little dilute nitric acid ; now carefully add a solution of ammonia (ammonium hydroxide, NH 4 OH) until the acid is neutralized. Record how you tested. Slowly evaporate this solution to dryness and examine the salt. What is its name ? Write equation for its formation. What salt would have been formed if sulphuric acid had been used ? Write an equation to show its formation. EXERCISE XVII. HYDROCHLORIC ACID, HC1. 78 to 87. 56. Put a little common salt, NaCl, into a dry test tube, and pour on it a few drops of strong sulphuric acid, H 2 S0 4 . Note the character- istics of the gas given off. Test with a piece of blue litmus paper. Try the same with other chlorides, as KC1, NH 4 C1, CaCl 2 . Do you obtain HC1 in each case? Write an equation for each reaction. LABORATORY EXERCISES. 25 57. In a flask fitted with a delivery and a thistle tube generate HC1 from 20 gr. sodium chloride, NaCl, 15 cc. water and 10 cc. sul- phuric acid, H 2 S0 4 . The contents of the flask must be very gradually and moderately heated, else a violent frothing is liable to occur which would spoil the experiment. In your notes, describe the apparatus; and show the use of safety tubes. Let the delivery tube pass into a cylinder, keeping it covered as well as possible. After filling three cylinders by downward displacement let the gas pass into a beaker of water. 58. Test a cylinder of the gas with a lighted taper; notice also the white fumes which are formed when the gas comes in contact with moist air. 59. Moisten a piece of paper with ammonium hydroxide, NH 4 OH, and thrust it into a cylinder of the gas. Explain what takes place, writing an equation for the reaction. 60. Invert a cylinder of the gas, well covered with a glass plate, in a dish of water, and then remove the cover. Explain. 61. Now test the liquid in the beaker. Has it acquired acid prop- erties ? Compare the action of this liquid with that marked HC1 in the reagent bottle, as to the effect on litmus; on a piece of marble; on zinc or iron, and on a solution of silver nitrate, AgN0 3 ; of lead nitrate, Pb(N0 3 ) 2 ; mercurous nitrate, HgN0 3 . If you should evaporate to dryness the solution in the flask used in Exp. 57 what would the residue be ? If you should recover all the compound formed, how much would you obtain ? EXERCISE XVIII. CHLORINE. 62. Ascertain what happens when HC1 is heated with such sub- stances as manganese dioxide, Mn0 2 ; red lead, Pb 3 4 ; or potassium bichromate, K 2 Cr 2 7 . Be cautious in inhaling the vapors. 63. 1 In the work with chlorine extreme care must be exercised not to allow the fumes to escape into the room. Keep all receivers well covered with paper. If the gas is accidentally inhaled, the antidote is vapor of alco- hol inhaled from a handkerchief; or ammonia. Into a flask fitted with both a delivery and a safety tube put about 20 g. of manganese dioxide, Mn0 2 . Pour upon it enough ordinary 1 Unless there is excellent draught in the laboratory, the instructor is to perform experiments 63, 64, 65, and 66 in the lecture room only. 26 LABORATORY EXERCISES. concentrated hydrochloric acid to cover it completely. The delivery tube should be bent downward and reach nearly to the bottom of the receiver, and should pass through a hole in a paper cover for the receiver. Shake the contents of the flask well together and apply very gentle heat. Fill several bottles with the gas. You can tell when they are full by the color. Note the specific gravity of the gas. After the necessary amount of gas has been collected in the receiver, let the gas pass into water for a few minutes and note if it is soluble. Complete the equation : Mn0 2 + HC1 = MnCl 2 + 2 H 2 + ? 64. Into a jar of chlorine, Cl, thrust a burning taper or a bit of flaming paper. Is the gas combustible ? 65. Into one of the vessels containing chlorine introduce a little (as much as you can put on a ten-cent piece) finely powdered an- timony ; or heat a small piece of copper foil and introduce it into the jar. What takes place ? Equation. In what respects is this experiment like the one in which iron was burned in oxygen ? 66. Into a vessel put a piece of paper with some writing on it ; some flowers, and some pieces of colored calico which you have moistened, and also pieces of written and printed paper. What takes place ? Into a fourth vessel put a dry piece of the same material. What difference is there in the action of the chlorine on the dry and on the moist material ? Printer's ink is made of lampblack (carbon) and is not bleached. How could you distinguish between organic and inorganic colors ? State the theory of the chemistry of bleaching. 67. Put into a small beaker 5 g. bleaching powder, CaCl 2 + Ca(C10) 2 ; set this in a large beaker, and hang in the latter the sub- stance to be bleached. Cover the large one with pasteboard, through which passes a thistle tube into the smaller. Pour through the thistle tube 5 cc. dilute H 2 S0 4 . Add more if needed. Explain the action. 68. Try the effect of the chlorine water made in Exp. 63 on a solu- tion of litmus, indigo, or cochineal. Also to solutions of silver nitrate, AgN0 3 , and lead nitrate, Pb(N0 3 ) 2 , in separate test tubes, add a little of the chlorine water. Describe the results and write equations. LABORATORY EXERCISES. 27 How does the action in this case compare with that of HC1 on the same solutions ? (See Exp. 61.) 69. Drop into a test tube 3 or 4 crystals of KC10 3 . Add a few drops of HC1 ; hold in the flame for a minute, and when action begins add 5 or 10 cc. H 2 0. Cautiously take the odor. What has been liberated? To 2 cc. indigo solution in a test tube add a little Cl water. Is the color discharged ? To 2 cc. cochineal solution add a little Cl water. Is the solution bleached? Try also litmus solution. To 2 cc. K 2 Cr 2 7 solution add a little Cl water. Is this bleached? K 2 Cr 2 7 is a mineral pigment ; cochineal is of animal origin. Explain the results. This is the ordinary method of making chlorine water for laboratory uses. Compare chlorine with hydrogen and oxygen. EXERCISE XIX. BROMINE AND IODINE. 104 to 117. 70. From the instructor receive into a flask or bottle of 1 or 2 1. capacity 3 or 4 drops of bromine, Br. Cover the bottle loosely and leave it standing. Immerse a piece of moist litmus paper in the gas. What is the effect ? Pour the liquid into a beaker of water, and note the specific gravity. 71. Warm gently a few crystals of KBr with 0.2 g. Mn0 2 and 1 cc. H 2 S0 4 in a test tube and observe the vapor. (Under the hood.) KBr + Mn0 2 + H 2 S0 4 = NaHS0 4 + MnS0 4 + H 2 + ? Could Cl be made in a similar manner? Illustrate by writing an equation. Test the remaining liquid with litmus paper. 72. Make a solution of a few crystals of potassium bromide, KBr, in 3 or 4 cc. of water. Add a drop of chlorine water (see Exp. 69). Is Br set free? Write the equation. Add to the above solution two drops of carbon bisulphide and shake the tube. What effect is produced ? How could free bromine be detected in a solution ? Why does carbon bisulphide not become colored when shaken in a simple solution of KBr ? 28 LABORATORY EXERCISES. EXERCISE XX. IODINE (continued). 117 to 126. 73. Examine a small crystal of iodine, I. How does it act upon the fingers ? Is it soluble in water ? Try alcohol. What is tincture of iodine ? 74. Hold a dry test tube in the gas lamp by means of the wooden nippers, and warm it along its entire length, so far as this is prac- ticable. Drop into the hot tube a small fragment of iodine. What is the color of the vapor V 75. Prepare a quantity of thin starch paste by boiling 30 cc. of water in a porcelain dish and stirring into it 0.5 g. of starch which has previously been reduced to the consistency of cream by rubbing it in a mortar with a few drops of water. Note the change in the starch. 76. Pour 3 or 4 drops of the paste into 10 cc. of water in a test tube, and shake the mixture so that the paste may be equally diffused through the water; then add a drop of an aqueous solution of iodine. Heat the solution gently until the color disappears, and allow it to cool again. This action affords a delicate test for iodine when not in combination. 77. Dip a strip of white paper in the starch paste and suspend it, while still moist, in a large bottle, into the bottom of which 2 or 3 crystals of iodine have been thrown. What does the experiment show ? 78. To a portion of the starch paste made in Exp. 75, add a few drops of potassium iodide solution. Into the paste thus prepared dip strips of filter paper. This is " iodo-starch paper." 79. Repeat Exp. 76, using bromine water instead of an aqueous solution of iodine. 80. Put 1 g. Mn0 2 in a test tube, pour upon it 1 cc. of HC1, and warm gently. Now hold a piece of iodo-starch paper over the tube and notice the result. Explain the action. This affords a test for chlorine. 81. Dissolve a few crystals of potassium iodide, KI, in 3 or 4 cc. of water. Add to this a little starch paste. Does it stain as the element did? Why not? Now add some chlorine water. What effect ? Write equation. Which has the stronger chemism, chlorine or iodine? LABORATORY EXERCISES. 29 82. How could you detect the presence of starch ? How could you distinguish between the three elements, Cl, Br, and I ? Classify these elements in accordance with their properties. EXERCISE XXI. SULPHUR. 126 to 136. 83. Place 1 g. of sulphur in a dry test tube and pour upon it 5 cc. of carbon bisulphide, cork tightly, and shake for a few moments. Car- bon bisulphide is volatile and very inflammable. Have no lights near by. Pour a little of the clear solution upon a watch glass and allow the carbon bisulphide to evaporate under the hood. Examine the residue obtained, and note the shape of the crystals. 84. Put 10 g. S into a test tube and slowly melt it. Notice the yellow color, and see that the liquid is very thin. It is now somewhat above 100. Heat it more strongly till it becomes black. It is now very thick and cannot be poured (200). Apply more heat till it grows thin again (300) . Heat to boiling (over 400) ; note the color of the vapor, and any sublimate on the test tube. Pour the S into water. Knead it, and note its elasticity. See whether it afterward changes. What is the product of burning sulphur ? Write the equation. 85. Place in a test tube a little litharge, PbO, and pour upon it a few drops of the H 2 S solution. The brown litharge changes to black. Lead oxide changes to lead sulphide. Give the equation. Similar reactions take place with other metallic oxides when treated with H 2 S. 86. Mix sulphur and iron filings in the proportion of 3.2 g. of the former and 5.6 g. of the latter, and heat the mixture in a test tube. After the mass glows, allow it to cool, and remove it from the tube. Examine the mass carefully. Is it sulphur that remains? Is it iron? Try iron and some of this mass in dilute HC1. What is the mass ? What does the experiment illustrate? Into 5 cc. of a solution of lead nitrate, Pb(N0 3 ) 2 pour a few drops of the H 2 S solution. What takes place? Write the equation. Add a little H 2 S solution to each of the following solutions, describe results, and write the necessary equations, HgCl 2 , CuS0 4 , CdCl 2 , BaCl 2 , CaCl 2 . 30 LABORATORY EXERCISES. EXERCISE XXII. HYDROGEN SULPHIDE. 136 to 151. 87. Put into a large side-neck test tube 5 g. ferrous sulphide, 10 cc. water, and 5 cc. HC1. Equation. Adjust a delivery tube, and pass the gas for a minute or two into 5 cc. HgO. Have the bearings tight. See whether this solution is acid, alkaline, or neutral. Use both colors of litmus. Put a drop of the H 2 S solution on Ag and Cu coins. Reactions. Put a drop of Pb(C 2 H 3 2 ) 2 solution on paper, and hold it in the vapor of H 2 S. This is the characteristic test for H 2 S. Mix equal parts of CuCl 2 and BaCl 2 , add H 2 S, shake and filter. What is in the filtrate ? What is on the filter ? 88. Referring to Exp. 87, state how you could test for a sulphide. Try Na 2 S or CaS. Write equation. EXERCISE XXIII. SULPHUR DIOXIDE. 151 to 156. 89. Light a piece of sulphur in a deflagrating spoon, and suspend the latter in a cylinder full of air. Carefully observe the odor. Immerse a lighted taper in the gas obtained. 90. Place a piece of copper foil in a test tube, cover with concen- trated sulphuric acid, and gently warm. Observe the odor. Place over the mouth of the tube some flowers. What change takes place in the flowers ? Does sulphur dioxide act in the same way that chlorine does ? 91. Put eight or ten pieces of sheet copper, 1 to 2 inches long and about an inch wide, into a flask. Pour 15 to 20 cc. concentrated sulphuric acid upon it. Heat gently. The moment the gas begins to come off, lower the flame and keep it at such a height that the evolu- tion is regular and not too active. Pass some of the gas into a bottle containing water. Is it soluble in water ? Collect a vessel full by displacement of air. (It is more than twice as heavy as air.) See whether the gas will burn or support combustion. Is the gas colored ? Is it transparent ? Has it any odor ? Does it burn? LABORATORY EXERCISES. 31 92. Charge a bottle, of the capacity of a liter or more, with sulphu- rous acid by burning in it a bit of sulphur. Fasten a shaving, or, better, a tuft of gun cotton, upon a glass rod or tube bent at one end in the form of a hook ; wet the shaving in concentrated nitric acid, and hang it in the bottle of sulphurous acid. Interpret what you observe. Pour a little BaCl 2 solution into the bottle before beginning the experiment, and notice its condition after shaking the bottle. Write the equation. EXERCISE XXIV. SULPHURIC ACID. 156 to 164. 93. Place in a beaker 20 cc. water ; pour gradually into the water about the same volume of concentrated sulphuric acid, stirring the mixture. Note the change of temperature. Save the dilute acid. 94. Put one drop of strong H 2 S0 4 , and one from that just made, on writing paper, and evaporate them high over a flame, so as not to burn the paper. When it is dry, examine. 95. Put 2 cc. of strong H 2 S0 4 into a test tube, and dip into it a splinter. Wood and paper are mostly cellulose, C 18 (H 2 O) 15 . Explain the charring. 96. To 2 cc. sugar solution, C 12 (H 2 0) n , add 2 cc. H 2 S0 4 , and explain. Cover a fragment of starch, C 6 (H 2 0) 5 , with H 2 S0 4 in a test tube ; boil till it begins to blacken. Explain. 97. Pour a few drops of H 2 S0 4 into a test tube, and dilute with 20 cc. of water. Add a little BaCl 2 solution, and observe the effect. This is the test for H 2 S0 4 . 98. Dissolve a crystal of Pb (N0 a ) 2 in water in a test tube, and add a few drops of dilute H 2 S0 4 . Note the action. This is another test for H 2 S0 4 . EXERCISE XXV. VOLUMETRIC COMPOSITION OF HYDROCHLORIC ACID. 99. For the Instructor. Arrange an apparatus for the generation of hydrochloric acid gas as follows : On a ring of the stand place an ordinary 500 cc. Erlenmeyer flask, fitted by a two-hole rubber stopper, with a glass stopcock funnel passing nearly to the bottom of the flask, and a glass tube bent at a right angle, which tube connects with a 32 LABORATORY EXERCISES. similar tube leading to the bottom of a 500 cc. Erlenmeyer filter flask with a side neck, to which is connected about a foot of rubber tubing which can be closed air-tight by means of a stout pinchcock. Place concentrated sulphuric acid in the funnel (stopcock closed), and common salt in the generating flask. Allow the acid to drop slowly on the salt, thus generating HC1. Warm the generating flask occasionally. Completely fill the filtering flask with the gas by down- ward displacement. Now put into the filtering flask containing the HC1 3 cc. of magnesium powder, and stopper the flask air-tight, plac- ing the free end of the rubber tube attached to the side neck in a beaker of water. Very cautiously open the pinchcock on the tube and allow the water to pass into the flask to about half its capacity. Close the pinchcock again (make sure it is tight) and allow to stand 12 hours or more, to complete reaction. Invert the flask to bring the water into the neck, open the pinchcock, and make the level of the water the same in the flask and the beaker. Why? Set the flask upright and measure the water, and also ascertain the total contents of the flask, making the necessary corrections. Remembering that the flask was at first full of HC1, which was then absorbed by the water and finally decomposed by the magnesium, as follows : 2HCl+Mg = MgCl 2 +2H. Calculate the relative proportion of chlorine and hydrogen in the original HC1 by volume. 100. Repeat Exp. 99, using 2 volumes of hydrogen to 1 of oxygen. How many volumes of gas remains after the combination ? What is this gas ? What was the condensation ratio in this case ? Why was there not a condensation in the previous experiment ? The resulting volume is a constant one for the union of gases. Make a table of the nitrogen oxides, showing the condensation in each case. From the above what may be learned of how the space occupied by the compound molecule compares with that occupied with the unit volume ? This double volume is often called the product volume of a compound gas. EXERCISE XXVI. MOLECULAR WEIGHTS. For the argument on " Molecular Condition of Gases " the student is referred to Storer and Lindsay's " Manual of Chemistry," 168 to 173. LABORATORY EXERCISES. 33 DETERMINATION OF MOLECULAR WEIGHTS. Accepting Avogadro's Law, and the dependent fact that the vapor density equals one half the molecular weight, it becomes possible to calcu- late the molecular weight of any substance which is naturally a gas or which can be easily vaporized. This is done by the physical method. 101. Fit as large a flask as will ride conveniently on the laboratory balance with a one-hole rubber stopper, through which passes a tightly fitting glass tube projecting an inch above the stopper, and reaching to the bottom of the flask. Close the outer end of the glass tube by means of rubber tube and pinchcock. Be sure that all joints are per- fectly tight. Fill the flask to the pinchcock with water, and ascertain by measurement the volume of the flask. Clean and dry the flask, and loosen the stopper, but do not remove it. Now counterpoise the flask on the balance, or ascertain the weight, and record the same in your notebook. Ascertain both the temperature at the balance and the barometer reading. Now calculate according to Exp. 22 the weight of air in the open flask (1 cc. of air at standard conditions weigh .001293 g.). Generate oxygen, passing it first through water and then through concentrated sulphuric acid contained in wash bottles. Allow the gas to pass through the tube of the weighing flask till a glowing match held near the loose stopper shows the flask to be full of the gas. Insert the stopper, and close the pinchcock. Allow any excess to es- cape by opening for an instant only the pinchcock. Read the tempera- ture, the barometer, and find the weight of the flask. Repeat the filling of the flask till there is no further gain in weight. As soon as a constant weight has been reached, again take readings of tem- perature and pressure, using the latter readings for the computation. The gain in weight, plus the weight of the air the flask held, is the weight of the oxygen at the temperature and pressure observed. Cal- culate the weight of the same volume at standard conditions. The weight varies directly as the pressure and inversely as the absolute temperature. Why? Calculate the density as referred to hydrogen. From this calculate the molecular weight of oxygen. EXERCISE XXVII. 102. Many substances cannot be vaporized, hence it is impossible to use the physical method above illustrated. Recourse in such cases is had to chemical methods. Assuming that the molecular weight of LAB. EX. 3 34 LABORATORY EXERCISES. oxygen has been ascertained to be 32, we- can now proceed to deter- mine the molecular weight of a substance in which oxygen is a com- ponent, as, for instance, potassium chlorate, KC10 r For this experiment the perfectly pure and dry salt must be used. Weigh into a small porcelain crucible about 2 g. of the pulverized dry salt. The weight must be accurately known, and the crucible should be provided with a cover. By means of a Bunsen flame heat the chlorate in the crucible gently at first, taking care all the time to avoid spat- tering and foaming. Continue the heating till the mass becomes solid after melting. Apply heat by means of the blast lamp till the mass again melts, after which remove to a desiccator, let cool, and weigh. Repeat the heating with the blast lamp, cooling and weighing till there is no further loss. Oxygen has been driven off (see Exp. 13). Calculate the loss due to escape of oxygen. In this experiment the chlorate gives up all its oxygen, and the chemical changes in which it has entered shows it to have three atoms of oxygen, hence a molecule of the salt has given enough oxygen to make one and a half molecules of oxygen gas. Form the propor- tion : wt. of oxygen given off : wt. in grams of chlorate taken : : mol. wt. of 1^ molecules of oxygen : x (= mol. wt. of KClOs). EXERCISE XXVIII. PHOSPHORUS. 182 to 194. In handling phosphorus extreme care must be exercised. Do not touch it with the Jingers. Always cut it under water, holding it with pincers. Always have water at hand to extinguish it if premature ignition takes place. 103. Put a piece of P as large as a grain of wheat in a test tube and pour on it immediately 1 cc. of carbon bisulphide, CS 2 . Does the P dissolve ? Pour a little of the liquid on a piece of filter paper, leaving enough of the paper dry to hold it by ; wave it back and forth till the CS 2 has evaporated. The P ignites spontaneously. Does the paper burn? Why not? Why does not the P on a match burn spontaneously ? What is the product formed when P burns ? Write equation. Why is P kept under water ? Why does it not combine with the of the water ? LABORATORY EXERCISES. 35 104. In an evaporating dish place a grain of P as large as a grain of rice, cover it with a beaker, and ignite it by means of a warm glass rod. Note the white, flaky product formed. What is it ? Breathe on it. What result ? Wash the inside of the beaker with a little water and test the water with litmus paper. In this experiment metaphosphorous acid is formed. Write the equation. 105. In an evaporating dish put about 2 g. of " glacial " phosphoric acid, HP0 3 , add about 3 oc. of water. After a little of the acid has dissolved pour the liquid into a test tube and add a solution of silver nitrate, AgN0 3 . White precipitate. Write equation. To the HP0 3 remaining in the dish add about 30 cc. water. Boil a few minutes. Test a little of the liquid with AgN0 3 ; then add 1 or 2 drops of water. Complete the equation : HP0 3 + H 2 = H 4 P 2 7 Write equation for the action of AgN0 3 on H 4 P 2 7 . Now fill the dish with water and boil for about 30 minutes. Test again with AgN0 3 , and the smallest drop of NH 4 OH. Are the precipitates the same in each case ? Complete the equations : H 4 P 2 7 + H 2 = H 3 P0 4 H 3 P0 4 + AgN0 3 = ? * 106. To 10 cc. of Na 2 HP0 4 solution add BaCl 2 in excess. Filter the mixture, and test the solubility of the precipitate in ammonia water and in dilute hydrochloric acid. 107. (I) To 2 cc. of Na 2 HP0 4 add 10 cc. of water, 2 cc. of NH 4 C1, and 1 cc. of ammonia water. Now add slowly 3 cc. of a mixed solu- tion containing MgS0 4 and NH 4 C1. Observe carefully under the microscope the character of the precipi- tate formed. 108. (II) To 5 cc. of Na 2 HP0 4 , add AgN0 3 solution in slight excess. Note the color of the precipitate. Filter and test its solubility in ammonia water, dilute nitric acid, and dilute hydrochloric acid. Save the silver residue. 109. (Ill) To a few cc. of Na 2 HP0 4 add an equal volume of a solution of ammonium molybdate in nitric acid. The last three experiments are the usual tests for phosphoric acid and phosphates. 36 LABORATORY EXERCISES. 110. Boil a little bone ash in water in a tube and filter the liquid. Test the liquid for phosphates as in Exp. 107. The bone ash con- sists of tricalcium phosphate, Ca 3 (P0 4 ) 2 , which is nearly insoluble, so there will be very little, if any, trace of phosphates in the liquid. Now place a little bone ash (or bone meal) in a test tube and add two or three drops of sulphuric acid ; warm the mixture. After the tube becomes cool, add water and shake the mixture ; filter. Test the filtrate as above. What has been the effect of the sulphuric acid on the bone ash? How do you know? The change is represented as follows : Ca 3 (P0 4 ) 2 + 2 H 2 S0 4 = CaH 4 (P0 4 ) 2 + ? This experiment illustrates the manufacture of superphosphates for fertilizing purposes. EXERCISE XXIX. SILICA AND CARBON. 215 to 229. 111. To a concentrated solution of " water glass " in an evaporating dish add enough concentrated hydrochloric acid, HC1, to render the solution acid. A jellylike mass of silicic acid, H 4 Si0 4 , will separate. Evaporate the contents of the dish to dryness, slowly, and then heat the residue gently over a lamp. After allowing to cool add water. Examine the contents of the dish for a fine white powder, Si0 2 . Give the reaction representing the change from H 4 Si0 4 to Si0 2 . 112. Put into a tube about 12 cm. long, enough soft coal to fill it about one third full. 1 Fit to the tube a delivery tube, and support the apparatus on the ring stand. Heat the coal and collect the gas. Test the gas with a flame. It consists of a mixture of carbon com- pounds. It is illuminating gas. It contains NH 3 . How could this be removed ? Remove the coke from the tube and examine it. See if it will burn, and note in what manner. 113. Repeat Exp. 112, using wood shavings or saw dust. Collect the gas and test it. After driving off the gas, remove the end of the tube from the water, plug it to prevent air entering, and allow the appa- ratus to cool. Finally remove the contents of the tube and examine. How does the charcoal burn ? 1 A clay pipe whose bowl is filled with the coal and sealed with plaster of Paris answers well. LABORATORY EXERCISES. 37 114. Over a burning candle hold the bottom of an evaporating dish or glass plate. Note the collection of lampblack (carbon). Explain why the smoke accumulates. 115. Mix on paper and put into a tube 10 parts CuO to 1 part C (powdered charcoal), by weight 5 g. in all. The tube should not be over one third full. Pass the gas into limewater contained in a test tube. What result? Carbon dioxide turns limewater milky. What is the appearance of the substance left in the tube? Does it suggest the metal copper, Cu ? Treat a little with concentrated nitric acid, HN0 3 . What should take place if the substance is metallic copper ? What does take place ? Write equation. What is the reaction which takes place between the copper oxide and the charcoal? Write equation. Compare the action of hydrogen with that of carbon on copper oxide. In what respects are they alike, and in what respects do they differ? EXERCISE XXX. CARBON (continued). 229 to 243. 116. Prepare a solution of H 2 S in a receiver with 20 cc. H 2 0. Notice the odor. Put into the receiver 5 g. powdered charcoal, and shake the mixture well. Pour the whole on a filter, collect the filtrate in a clean receiver, and see whether any odor remains. If so, use more charcoal and filter again. 117. Make a filter of boneblack by fitting a paper filter into a funnel 8 to 10 mm. (3 to 4 inches) in diameter at its mouth. Half fill this with boneblack. Pour a dilute solution of indigo l through the filter. What effect does this have on the color of the solution? Do the same thing with a dilute solution of litmus. If the color is not completely removed by one filtering, filter the solution again. Try also a solution of potassium bichromate, K 2 Cr 2 7 . As indigo and litmus are organic coloring matters, and K 2 Cr 2 7 is mineral, state any inference. 1 Prepared by treating | g. of powdered indigo for some time with 4-5 cc. of warm concentrated sulphuric acid and diluting with a liter of water. 38 LABORATORY EXERCISES. 118. Mix 4 g. of potassium nitrate with 2 g. of powdered charcoal. Place the mixture upon an iron plate and touch it with a lighted stick. What does it act like ? EXERCISE XXXI. CARBON DIOXIDE. 243 to 247. 119. Put a short piece of candle on the desk and invert over it a dry, wide-mouthed bottle. What forms on the inside of the bottle ? Remove the bottle and pour into it a small quantity of limewater and shake. What is the effect on the limewater? What is formed? Write the equation for the product of combustion and for the action of the limewater. 120. Into a beaker of limewater blow the breath through a glass tube. Describe the result. Continue to blow till the solution finally clears. What is the effect of 'a saturated solution of C0 2 on calcium car- bonate? 121. In a flask, arranged as for the generation of hydrogen, place 10 or 12 gr. of calcium carbonate (limestone), CaC0 3 . Cover the end of the thistle tube with water, and add by small successive portions strong hydrochloric acid. Collect several bottles of the gas. Try to collect it by downward displacement. Thrust a lighted taper into a bottle of the gas. Note the result. From a large bottle of the gas pour a quantity on a lighted candle. Is the gas heavier or lighter than air ? 122. Fill a bottle half full of the gas. Cork under water and shake. Lower the mouth of the bottle into the water again, remove the stop- per, and note if the water rises on the inside. To what is this due ? What is " soda water " ? What other gases will dissolve in water? 123. Pass carbon dioxide into a solution of 2 gr. of caustic potash (potassium hydroxide) until it will absorb no more. Add any dilute acid in 2 or 3 cc. of water in a test tube to the solu- tion thus obtained. What gas is given off when the acid is added? How do you know ? Write the equations expressing the reactions which take place on passing the carbon dioxide into the caustic potash solution, and on adding an acid to the solution. 124. Put about a gram of sodium carbonate, Na 2 C0 3 , into each of four test tubes, and then add to one tube about 4-5 cc. of dilute hydro- chloric acid, to a second add the same quantity of dilute sulphuric LABORATORY EXERCISES. 39 acid, to a third the same quantity of dilute nitric acid, and to the fourth the same quantity of dilute acetic acid. What takes place? Is a gas given off ? Pass it through limewater. What is it ? Perform the same experiment with small pieces of marble. What gas is given off? What conclusions can you draw from these observations? How can you easily detect carbon dioxide ? EXERCISE XXXII. FLAME. 257 to 267. 125. Examine the structure of a Bunsen burner (unscrew the top), make a drawing to show the orifices, and state what use each sub- serves. Light the gas at the base for a minute. Replace the top, and light the gas at the top of the burner. Hold the flame in front of a dark object, examine the parts, make a drawing, give a brief description, and state the color of each part. Put the flame in direct sunlight, and study the parts from its shadow, to confirm your results. Make a careful examination of the parts and colors of a candle flame, and make a drawing to show them. Move it slightly in the air to show the outer flame. This is best seen in a dark room. 126. Light the gas of a Bunsen burner. Put a stick across the base of the flame for an instant, and notice what parts are burned. Make a sketch. Hold a stick just above the inner blue cone of the flame. Press quickly down on the flame with a paper (remove before it burns) and notice the shape of the charred part. Sketch. Press down on the flame with a fine wire gauze, and observe by the glowing of the wire where the heat is most intense. Test the heat of the inner cone with the end of a platinum, Pt, wire. Notice that it glows when near the top, but not elsewhere in this cone. Put one end of a small tube into the inner blue cone, and try to light the gas at the other end. From the above, state what takes place in each of the two chief parts of the flame. 127. Observe the light of a Bunsen flame, and its color. Sprinkle a very little charcoal dust in the flame, and note any change of light or color. 40 LABORATORY EXERCISES. Close the orifices at the base of the burner, and explain the change of light. Hold an evaporating dish in the upper part of this closed flame for a minute, and notice deposit. Now open the orifices and persistently try to burn off a little of the deposit from the evaporating dish. What is the cause of light in a flame ? 128. Ignite the gas and hold a fine wire gauze 3 or 4 cm. above the burner. Why does it not burn above the wire? Extinguish, then relight, the gas above the gauze. Result. Gradually lift the wire till the gas will not burn. Again light the gas above the gauze, and hold another gauze above the flame, so as to confine it above and below. From this experiment define kindling point, and state three condi- tions of combustion. 129. Put a fragment of lead, Pb, not larger than a pea on a piece of charcoal, slightly hollowed out to hold it. Insert the metallic tube in a Bunsen burner, and with a mouth blowpipe direct the oxidizing flame strongly and steadily against the Pb for 4 or 5 minutes. As you stop blowing, notice the yellow vapor that escapes from the pellet of Pb; note also, as it cools, the yellow coating of lead oxide, PbO, on the coal. Write the reaction. 130. Put \ gr. PbO on a piece of charcoal. With the blowpipe di- rect the reducing flame steadily against it for some time, or until a metallic pellet is obtained. What is it ? Equation. EXERCISE XXXIII. METHANE AND ILLUMINATING GAS. 131. Mix together 2 g. of crystallized sodium acetate, 4 g. of caustic soda, and 8 g. of slaked lime. Heat the mixture gently upon an iron plate, until all the water of crystallization of the acetate has been expelled, and the mass has become dry and friable. Charge an igni- tion tube 20 cm. long with the dry powder, heat it above the lamp, and collect the gas at the water pan. Marsh gas is evolved from the mixture, at a temperature below redness, and a residue of sodium car- bonate is left in the ignition tube. The purpose of the lime is to ren- der the mass porous and infusible, or nearly infusible, so that the tube may be heated equably. The reaction may be represented as follows : NaC 2 H 3 2 + NaOH = CH, + LABORATORY EXERCISES. 41 132. Fill a tall bottle of at least one liter capacity with warm water, invert it over the water pan, and pass marsh gas into it, until a little more than one third of the water is displaced ; cover the bottle with a thick towel, to exclude the light, and then fill the rest of the bottle with chlorine. Cork the bottle tightly, and shake it vigorously, in order to mix the gases together, keeping the bottle always covered with the towel. Finally, open the bottle, and apply a light to the mixture. Ignition takes place, hydrochloric acid is produced, while the sides and mouth of the bottle become coated with solid carbon in the form of lampblack. The presence of the acid may be proven by the smell, by its reaction with moistened blue litmus paper, and by the white fumes which are generated when a rod moistened with ammonia water is brought in contact with the escaping acid gas. 133. Fill an ignition tube a third full of bituminous coal. Hold it steadily over a lamp to heat it, meantime trying to ignite the escap- ing gas. Note the color of the flame, and see whether any soot is deposited on porcelain held in the burning gas. Break and examine the tube for a tarry residue arid for coal. Put the coal on an iron plate, and bring a Bunsen flame in contact with it, noting whether it burns with a flame or only glows. Only gas burns with flame. Write no reactions, but state how many and what products you observed in this experiment. EXERCISE XXXIV. ALCOHOL. 134. Introduce 20 cc. of molasses into a flask of 200 cc., fill it with water to the neck, and put in half a cake of yeast. Fit to this a delivery tube, and pass the end of it into a test tube holding a clear solution of limewater. Leave in a warm place for two or three days. Then look for a turbidity in the limewater, and account for it. Is the liquid remaining in the flask sweet? What has become of the sugar? This is a fermented liquor. 135. Pour off one half of the fermented liquor, and reserve it in a loosely covered dish ; with the remainder proceed as follows : Support the flask on the iron lamp stand, and by means of a delivery tube con- nect it with a second flask capable of holding one third of the liquid, and placed on a water bath. From this second flask a delivery tube is carried to a small flask kept cool by immersion in cold water. Heat the liquid in the largest flask, so that it just boils. The vapor of alco- 42 LABORATORY EXERCISES, hoi, together with a certain amount of steam, passes into the second flask, which is kept just below the boiling point of water by being supported on the water bath in which the water barely boils. At this temperature a considerable portion of the alcohol, together with some water, passes over into the third flask, where it is condensed. Con- tinue the operation until about one third of the liquid has passed out of the large flask. The liquid obtained in the third flask is a dilute alcohol ; the odor of alcohol is distinctly perceptible, but the alcohol may not be strong enough to burn. In that case support the third flask on the wire gauze over the lamp, and connect it by means of a delivery tube with another small flask, which is kept cool. Heat the contents of the flask gently until they just boil, and transfer the first teaspoonful of the liquid which condenses in the cooled flask to a porcelain dish. If the experiment has been successfully conducted, the alcohol thus obtained will be strong enough to take fire if a flame be brought into contact with it. Taste and smell the distillate. Try to ignite it. 136. Put a little of the white of egg into a beaker ; cover it with strong alcohol and note the effect. Strong alcohol has the same coagu- lating action on the brain and on the tissues generally, when taken into the system, absorbing water from them, hardening them, and contracting them in bulk. EXERCISE XXXV. ETHER. The student should never attempt to perform any experiment requiring more than a very minute quantity of ether, since it is highly dangerous to work with this substance on account of its great volatility and ready inflammability. 137. Into a small test tube put 10 drops of ordinary alcohol and as much strong sulphuric acid, and heat the mixture gently over the lamp. Ether will be formed, and may be recognized by its peculiar odor. 138. Pour a small quantity of ether into the palm of the hand, and observe the rapidity with which it evaporates, and also the cold pro- duced by this evaporation. 139. Into a tumbler or other very wide-mouthed vessel put a few drops of ether. Cover the vessel loosely, and allow to stand for a few moments ; then bring a lighted match to the mouth of the vessel : the heavy vapor of ether will have displaced the air in the vessel, and will take fire at the mouth of the vessel with a sudden flash. LABORATORY EXERCISES. 43. 140. Into a small test tube put 10 drops of ordinary alcohol, and the same amount of strong sulphuric acid. Add a crystal of sodium acetate as large as a small pea, and heat the mixture gently. Acetic ether, ethyl acetate, is formed, and may be recognized by its peculiar odor. Examine the alcoholic liquid preserved from Exp. 135 by taste, smell, and litmus. The alcohol has changed to what ? Preserve the liquid for subsequent work. EXERCISE XXXVI. ACETIC ACID AND ACETATES. 141. To the acid liquid of the previous experiment or to 40 or 50 cc. of common vinegar, add powdered chalk (calcium carbonate) as long as the addition causes effervescence. Calcium acetate is formed and remains dissolved in the liquid. Filter the solution, and evaporate the nitrate to dryness at a gentle heat. The solid residue is an impure calcium acetate. Place a portion of this calcium acetate in a small test tube, and heat gently with a few drops of strong sulphuric acid. Acetic acid will be set free, and may be recognized by its pecul- iar odor. If ordinary vinegar be used in this experiment, it will be better to decolorize the solution of calcium acetate by mixing it with powdered bone black before filtering. SOAP. 142 (a). Dissolve 15 g. of solid caustic soda in 120 cc. of water. When the suspended impurities have settled to the bottom of the solution, pour off one half of the clear liquor into a deep iron or porcelain dish of at least 500 cc. capacity, add an equal bulk of water, and 50 g. of beef tallow. Bring the mixture to boiling and boil it steadily for three quarters of an hour, supplying from time to time the water .lost by evaporation; then add the remainder of the solution of caustic soda, and continue to boil steadily for an hour or more, allowing the liquid to become somewhat more concentrated toward the end of that time ; then add 20 g. of fine salt, boil for a minute or two, and allow the liquid to cool. A part of the mass becomes solid, and rises to the top ; it is hard soap. The chemical action is thus explained: when tallow (glyceryl stearate and oleate) is boiled with sodium hydroxide, there is formed sodium stearate (and oleate) and glyceryl hydroxide. When common salt is added, the soap (sodium stearate and oleate), being insoluble 44 LABORATORY EXERCISES. in the saline liquid, separates as a solid. The liquid remaining con- tains in solution the excess of sodium hydrate employed, as well as the salt and the glycerine. (&). Soap may be made more quickly by using castor oil in- stead of beef tallow. Mix 100 cc. of castor oil and 100 cc. of caustic soda solution prepared as above, and boil for 30 minutes. Then add 150 cc. of water, bring to a boil, and add 20 g. of salt. The soap rises to the top and may be removed when cold. Castor oil is mainly glyceryl ricinoleate; the chemical change is similar to that just described. 143. Heat some of the soap with soft water. A nearly clear solu- tion will be obtained if the decomposition of the tallow or oil was complete. Add dilute hydrochloric acid until the solution is decidedly acid. The liquid will become turbid, and on standing will become covered with a layer of a fatty substance, which is a mixture of stearic and oleic acids (or mainly ricinoleic acid if castor oil was used). The sodium chloride formed will be held in solution by the liquid. 144. Dissolve some of the soap above made in water, and render the water hard by adding 10 cc. of a clear solution of calcium or magnesium sulphate. Shake and note any change. An insoluble soap has been formed. f stearate 1 Sodium j palmitate t + calcium sulphate = ? I oleate J EXERCISE XXXVII. SUGAR. 145. Heat cautiously a small quantity of white sugar in a porcelain dish until it melts. Allow the pasty liquid to cool rapidly. The product is barley sugar. Heat again to a higher temperature, but not too high ; the sugar turns brown, froths, and gives off pungent vapors, and there remains a dark brown mass. This is caramel. 146. Into a flask of 250 cc. capacity introduce 100 cc. of water. Add 1 cc. of strong sulphuric acid, and heat the mixture to boiling. In a porcelain mortar rub 10 g. of starch with enough water to make a cream, and pour the mixture, little by little, into the boiling liquid, taking care not to interrupt the boiling. The starch dissolves without forming a paste. Boil for three or four hours, replacing from time to time the water lost by evaporation, and then add powdered chalk (calcium carbonate) until the liquid is no longer acid. When the LABORATORY EXERCISES. 45 mixture has become cold, filter off the insoluble calcium sulphate formed by the action of the sulphuric acid on the calcium carbonate, and evaporate the solution at a gentle heat to a sirupy consistency. The solution contains dextrose, which, on long standing, may separate from the liquid in crystals. EXERCISE XXXVHI. SOLUBILITY. 147. Solubility in water. Weigh three portions of 1 gr. each of powdered Na 2 S0 4 , CaS0 4 , and PbS0 4 . Take three test tubes and place in each 10 cc. of water. Into one of the tubes pour one portion of Na 2 S0 4 and shake, and, if this dissolves, add the second ; and, if the second dissolves, add the third. Repeat the process with the CaS0 4 and PbS0 4 in the other tubes. If, however, the first lot fails to dissolve, ascertain whether any has dissolved by filtering some of the mixture very carefully and evaporating a few drops of the filtrate on a clean porcelain. 148. Solubility in hot and cold water. Weigh four portions of 1 g. each of powdered Ba(N0 3 ) 2 . Place 10 cc. of water in a test tube and add one portion and shake. Note the result. Warm slowly, and as often as the salt is entirely dissolved add a new portion of 1 g. Finally bring the liquid to boiling. What does the experiment show ? 149. Use of different solvents. Compare the solubility of iodine in water, carbon disulphide, and alcohol. Apply no heat. Use a very small quantity of I, and 3 or 4 cc. of the liquid. Note the rapidity with which the I disappears, and judge in which liquid it is the most soluble. 150. Compare the solubility of Nad in H 2 0, concentrated HC1, and alcohol. EXERCISE XXXIX. SOLUBILITY (continued). 151. Solubility in mixtures. Try to dissolve BaCl 2 in concentrated HC1. Finally add 10 cc. H 2 0. 152. Dissolve 2 gr. Na 2 S0 4 in 5 cc. H 2 and add an equal volume of alcohol 153. Neutralization of the solvents. Dissolve calcium phosphate in warm dilute HC1, then add NH 4 OH to alkalinity. Repeat, using barium oxalate. 154. Solution of liquids. Test the solubility of the following liquids in water : alcohol, ether, olive oil, glycerol, carbon bisulphide. 46 LABORATORY EXERCISES. Proceed in each case as follows : Take 5 cc. of water in a clean test tube; pour 1 cc. of the liquid to be tested upon the water in the tube. Shake several times, and then observe the depth of the liquid layer above or below the water, if any such layer there be. Ether and car- bon bisulphide are very inflammable and very volatile, and must not be brought near aflame. 155. Physical and chemical solution. Take two portions of 5 gr. each of salsoda. Dissolve one portion in 10 cc. of water, evaporate to dryness slowly, and compare the substance obtained with the original salt in appearance, crystalline form, and taste. Dissolve the second portion in dilute hydrochloric acid, evaporate, and compare. 156. Place 5 cc. of alcohol in a test tube and add 1 cc. of HC1, and shake. Drop a small piece of fused potassium carbonate into the tube. Place 5 cc. of water in a test tube and add 1 cc. of HC1. Drop a small piece of fused potassium carbonate into this tube. Explain. EXERCISE XL. CRYSTALLIZATION. 157. Crystallization by solution and evaporation. Place 5 gr. of oxalic acid in a test tube, and add 10 cc. of water. Heat until the acid has dissolved, and then allow to cool. 158. Crystallization by sublimation. Heat a small lump of benzoic acid gently in a dry test tube. 159. Crystallization by precipitation. To a few drops of concen- trated common salt solution add 5 cc. of alcohol. Examine the pre- cipitate under a microscope. 160. Purification by crystallization. Weigh 25 gr. of soda ash. Place in a beaker, and pour upon the ash 50 cc. of water. Heat until the ash has dissolved, filter the turbid solution while hot, and allow to cool. Remove some of the crystals, place in an evaporating dish, and heat until they are reduced to a fine powder. Compare this powder with the original powder. LABORATORY EXERCISES ON METALS. EXERCISE XLI. SODIUM AND POTASSIUM. 161. With a pair of pincers remove a small piece of metallic sodium, Na, from the oil, and dry on a filter paper. Scrape off the outside coating, handling the metal all the time with the pincers. Note whether the metal is hard or soft, also its color and relative f v- i > * * A? T> f or THE r \ I UNIVERSITY LABORATORY EXERCISES. V P OF 47 weight as compared with water. When the sodium is placed on water what takes place ? What is the color of the flame ? When action ceases test the solution remaining with litmus paper, and draw conclusion as to class of compounds to which it belongs. Write the equation for the action of sodium on water. 162. Perform the same experiments as described in Exp. 161, sub- stituting metallic potassium, K, for sodium. Compare results. Write the equation for the action of K on H 2 0, for the combustion of H, and for the combustion of K. 163. Test some potassium carbonate, K 2 C0 3 , solution with litmus paper. Also add to a little of the salt a few drops of dilute hydro- chloric acid. Is a gas given off? What is it? Treat a pound of wood ashes with water, and filter off the liquid, leaching the same ashes several times with the nitrate. Examine in the same manner as the above solution. Evaporate to dryness, and test the solid remaining with a little dilute hydrochloric acid. Does it act like potassium carbonate? Save the material for use in Exp. 165. 164. Examine a piece of caustic potash or soda, and note its solu- bility. Note its action on litmus paper. Place the solution in an evaporating dish and add slowly, while stirring with a glass rod, dilute hydrochloric acid till the solution is slightly acid. Evaporate to dryness. What is the powder left ? Why is it not caustic potash ? EXERCISE XLII 165. Tests for potassium. Insert a piece of perfectly clean platinum wire in a solution of potassium chloride, and hold it in the non-lumi- nous flame of a Bunsen burner. Note the color imparted to the flame. Try the same with a sodium chloride solution. Is the same color imparted to the flame ? Mix a little of the solutions and repeat the test. Could you tell that potassium was present in the mixture by means of the naked eye? Repeat each of the experiments, using blue glass through which to view the flame. Describe the results. To about 2 cc. of a solution of potassium chloride add a like quan- tity of platinic chloride. What is the result? Try the same with a sodium solution. 48 LABORATORY EXERCISES. From the above experiments explain how you could detect the dif- ference between a sodium and a potassium compound. Try the tests on a solution of the material obtained in Exp. 163. EXERCISE XLIII. 166. A laboratory use of borax. What is the formula for borax, and what is its other chemical name ? Make a loop in a piece of platinum wire as large as a capital here printed. Heat the loop red hot in the flame, and thrust it while still hot into some powdered borax, a quantity of which will adhere to the wire. Reheat the loop in the oxidizing flame, and fuse the borax to a clear glass. When a good bead has been obtained touch it while still hot to a very small particle of maganese dioxide. Reheat in the oxidiz- ing flame till the particle is seen to dissolve and diffuse through the bead. Look through the bead toward the light. What is the color ? Make a similar experiment, using copper oxide instead of manganese dioxide. Ascertain if it makes any difference whether the oxidizing or the reducing flame be used. For what purpose might borax be used by the chemist as indicated by the above experiment ? EXERCISE XLIV. 167. Name the principal sodium compound. From this compound show by what chemical reactions there could be obtained sodium sul- phate, sodium carbonate, sodium hydroxide. Start with 20 g. of this principal salt, and using such other material as may be necessary, make some potassium chloride. Prove that this is a potassium compound. EXERCISE XLV. AMMONIUM. 168. Bring near to each other two vessels, one containing a little strong ammonium hydroxide and the other a little concentrated hy- drochloric acid. Explain the result. Try strong nitric acid instead of hydrochloric acid. Try sulphuric acid. How do you explain the difference? 169. Re-read Exp. 53 and the account of your work on that experi- ment. How could ammonium salts be detected? LABORATORY EXERCISES. 49 170. In a dry small test tube heat a few small pieces of ammonium chloride. What collects in the upper part of the tube ? Such action is called sublimation. Prove by some test that it is an ammonium salt. Prove that it is ammonium chloride by some test. (See Exp. 61.) Ammonium salts of volatile acids are volatile. Ammonium salts of non-volatile acids lose their ammonia on heating. All ammonium salts are decomposed by heating. 171. Dissolve a small piece of alum in a test tube and add am- monium hydroxide till after shaking the solution smells slightly of ammonia. What is the precipitate? Write the equation. What is meant by a precipitate ? Why are substances thus precipitated? Judg- ing from this experiment, for what purpose could ammonium hydroxide be used in the laboratory ? EXERCISE XLVI. CALCIUM. 172. Dissolve about 10 gr. of marble in hydrochloric acid. Evapo- rate to dryness, and allow some of the residue to be exposed to the air for a time. Does it grow moist? What is the compound you have made and which thus absorbs the moisture? Write the equation to show the reaction. If you should add concentrated sulphuric acid to some of this resi- due, what would you expect to be the result? Try it. What is the solid residue from the sulphuric acid treatment? Is it soluble or insoluble ? How could you tell whether a substance is potassium chloride, sodium chloride, calcium chloride, or ammonium chloride. 173. Review the results obtained in Exp. 120, and explain how they indicate the manner of cave formation. 174. On a piece of quicklime (what is the formula?) in a beaker pour some warm water. The lime slakes and falls to pieces. Is heat developed? Does the lime increase in bulk? After standing for some time the undissolved lime will settle out. The clear liquid is called limewater. 175. Shake a little powdered plaster of Paris (what is the for- mula?) with water for a few minutes. Filter, and divide the nitrate into four parts, using one part for each of the following: (a) To one add a few drops of a solution of barium chloride. A turbidity of barium sulphate shows the presence of sulphuric acid. LAB. EX. 4 50 LABORATORY EXERCISES. (&) To another part add a little ammonium oxalate solution. A white precipitate shows lime. The two tests together show that calcium and sulphuric acid are present. If both are present, in what form must they be? (c) To a third portion add a little soap solution. A scum or pre- cipitate indicates a hard water. (d) Boil the fourth part. No precipitate shows permanent hard- ness. What is the cause of temporary hardness ? (See Exp. 120.) How can it be removed ? EXERCISE XLVH. MAGNESIUM. 176. Characteristic action of magnesium. Pour 5 cc. of magnesium chloride into each of two test tubes. To one add a little ammonium hydroxide, and to the other add an equal volume of ammonium chlo- ride and then a little ammonium hydroxide. Repeat the experiment, but use ammonium carbonate instead of the hydroxide. What is the precipitate obtained in each case ? Why is a precipitate not obtained when the ammonium compounds are present? 177. To 3 cc. of a magnesium chloride solution add about 5 cc. ammonium chloride and an equal amount of ammonium hydroxide. Dilute with an equal volume of water, and add about 3 cc. disodium phosphate, Na 2 HP0 4 , solution and allow to stand. A crystalline pre- cipitate thus obtained is a test for magnesium. The precipitate is am- monium-magnesium phosphate. Write the formula for it. The precipitate is soluble in acids, and reprecipitated by ammonia. 178. Zinc. Try the effect of nitric, sulphuric, and hydrochloric acids, dilute, on a piece of metallic zinc. Describe the results. 179. In a solution of lead acetate (or copper sulphate) immerse a piece of sheet zinc. What collects on the zinc ? 180. Place 5 cc. zinc sulphate solution in a test tube and add a few drops of caustic soda solution; finally add a larger quantity. In excess of alkaline hydroxides soluble zincates are formed. EXERCISE XLVIII. LEAD. 181. Examine a piece of sheet lead. Observe its color and softness. Heat a small piece on charcoal with the blowpipe. Does it melt easily ? Try its solubility in nitric, hydrochloric, and sulphuric acids. What is the compound formed in each case ? Is each soluble ? LABORATORY EXERCISES. 51 182. Heat any lead compound with soda on charcoal. What do you obtain ? How do you know ? 183. To a solution of a lead salt add dilute hydrochloric acid till no more precipitate forms. Write equation. What is the precipitate? Filter off the precipitate, and to one third of nitrate add dilute sul- phuric acid; to the other third add hydrogen sulphide; add to the re- mainder potassium chromate solution. Write equation for each test. Ascertain if the original precipitate is soluble in hot water. If the precipitate dissolves, ascertain if soluble chlorides and sul- phates coul,d be substituted for their respective acids as tests for lead. From the above experiment answer the following: Is lead thrown out of a solution completely by hydrochloric acid or a soluble chloride ? How could lead be detected in an unknown solution ? EXERCISE XLIX. MERCURY. 184. Examine the action of the common mineral acids on mercury after noting the physical characteristics of the metal. Heat a little in a glass tube closed at one end and note the result. Drop a frag- ment of zinc into a little mercury. What change takes place ? Define amalgam. In what special physical property does this metal differ from the others you have studied ? 185. Put into each of two test tubes 4 or 5 cc. of water. Add to one a little mercurous chloride and to the other mercuric chloride, and boil each. Which dissolves? 186. In one of two test tubes place about 5 cc. of mercuric nitrate solution, and in another a like amount of mercurous nitrate solution, and add to each hydrochloric acid as long as a reaction is noted. Explain the result. In the same way try the action of these salts with caustic soda. What is formed in each case ? 187. Compare the action of the two salts toward hydrogen sulphide. 188. To a solution of mercurous nitrate add a solution of stannous chloride. What is the precipitate ? Try the action of the same re- agent on mercuric chloride. The precipitate is at first white, then gray. This is a characteristic reaction. EXERCISE L. COPPER. 189. Examine some pieces of copper. Try their solubility in the common mineral acids. What is the best solvent? 52 LABORATORY EXERCISES. 190. With the blowpipe heat on charcoal a little copper chloride with an equal amount of sodium carbonate. Explain the result in full. 191. To a solution of copper chloride add sodium carbonate solu- tion so long as a precipitate forms. Boil till the precipitation turns black. What chemical changes take place? Filter, dry, and save the powder, and make a borax bead test as in Exp. 166. What is the color of the bead ? This is a characteristic action of copper compounds. 192. To a solution of copper sulphate add some caustic potash solution. What is the result ? After observing the characteristics of the precipitate, heat, and notice the change. Explain the action and write equations for the reaction. 193. To a solution of copper sulphate add ammonia water, drop by drop, shaking the solution after each addition. Note all changes. This is a delicate test for copper. 194. In a solution of copper sulphate dissolve a few crystals of tar- taric acid and then add an excess of sodium hydroxide. Finally add a small piece of glucose, and boil the solution. What takes place ? What is the composition of the precipitate ? What are the formulae and names of the oxides of copper? 195. Acidify 5 cc. of copper sulphate solution with acetic acid, and add a few drops of potassium ferrocyanide solution. This test is a more delicate one for copper than that in Exp. 193. 196. To a solution of copper sulphate add a little hydrochloric acid. Is a precipitate formed ? Pass sulphuretted hydrogen through the solution. What is the result? Filter and wash thoroughly. Treat the precipitate with nitric acid, and boil with water. Test for copper as indicated in the previous experiments. How could you detect copper in a solution ? EXERCISE LI. SILVER. 197. Fill three test tubes one third full of water and pour into each a few drops of silver nitrate. Add to one about 3 cc. of hydrochloric acid, and shake the tube violently. To the second tube add a like quantity of potassium bromide solution and shake the tube. To the third^tube add a like quantity of potassium iodide solution. Describe the results in each case and note the color of the precipitates. Name each precipitate and write the equation for each reaction. Withdraw a little of the precipitate from each tube and test the solubility in nitric acid. LABORATORY EXERCISES. 53 Test also the solubility of the precipitates in ammonium hydroxide and in a solution of sodium thiosulphate. Heat some of the precipitate from one tube with a little sodium carbonate on charcoal before the blowpipe. The action may be con- sidered typical of that of the other silver compounds. 198. Precipitate from a silver solution some of the chloride by means of hydrochloric acid (could a solution of sodium chloride be used to do this?). Throw the precipitate on a filter and wash it thoroughly with water. Open the filter and spread the precipitate evenly over the filter, and place it in direct sunlight. Upon the facts illustrated in these two experiments depends the art of photography. EXERCISE LII. ALUMINUM. 199. Examine the physical properties of a piece of aluminum. Try the action of the common mineral acids on the metal as in the .pre- ceding experiments with the other metals. 200. Dissolve a small crystal of aluminum sulphate in water, in a test tube. Place one half of the solution in another test tube, and add drop by drop ammonium hydroxide, till after shaking the odor of ammonia is permanent. Write the equation for the reaction. Place a little of the moist hydroxide in another tube and try to dissolve it in ammonium hydroxide. What result? In a similar manner try its solubility in sodium hydroxide. Sodium aluminate is formed. 201. With the remainder of the solution made in Exp. 200, ascer- tain if other alkaline hydroxides will precipitate aluminum hydroxide. Ignite a small quantity of the precipitate on charcoal before the blow- pipe. Moisten the solid obtained with a drop of cobalt solution ; heat again and note the color of the mass obtained. This is a test for aluminum. 202. Dissolve a little " alum " and test for aluminum as indicated in the above experiments. How could you show that alum contains potassium ? That it is a sulphate ? That it contains water of crystallization ? 203. To a small quantity of an organic coloring matter (cochineal is good), add a little potassium sulphate and finally a little ammonium hydroxide. A colored precipitate is thrown down, which is known as a " lake " color. If cochineal is used, it is a carmine lake. It is simply colored aluminum hydroxide. 54 LABORATORY EXERCISES. 204. Make a solution of aluminum acetate by adding to a solution of lead acetate a solution of alum and filtering off the lead sulphate. In this solution soak a piece of cotton cloth and put it away till the next exercise. EXERCISE LIE. 204. (Continued). Treat the cotton cloth prepared at the last exercise as well as a piece of ordinary cotton cloth, with a solution of logwood, and observe the relative amount of color imparted. To what is the different result due ? What is a material that will act thus called ? IRON. 205. Dissolve a few small pieces of iron (tacks will answer) in about 10 cc. of dilute sulphuric acid. When action ceases, dilute with an equal bulk of water. What is given off during the action ? What is in solution ? Write the equation for the action. Add to the solution ammonium hydroxide till the smell of am- monia is permanent, and collect the precipitate on a filter. What is the precipitate? What is its color? Write the equation for the reaction. Ascertain if the precipitate is soluble in hydrochloric acid. 206. Ascertain the action of the hydroxide on the borax bead in a loop of platinum wire. What is the color of the bead ? 207. To a solution of ferrous sulphate add ammonium hydroxide in excess. Observe color of precipitate. Allow it to stand some time, stirring frequently, and observe any change of color. Explain the reaction. 208. Dip a piece of cotton into a solution of nutgalls or tannin, and dry it. Then dip it into a solution of copperas and dry again. Finally try to wash out the color. 209. Soak a piece of cotton cloth in a solution of ferric sulphate and then immerse it in an acidulated solution of potassium ferro- cyanide. Prussian blue is precipitated in the fibers of the cloth. EXERCISE LIV. FERROUS AND FERRIC COMPOUNDS. 210. Ferrous salts. Dissolve a little iron in dilute hydrochloric acid. Place half the solution in another beaker or test tube for use in the next experiment. To a part of the solution reserved for this experiment add a few drops of potassium ferricyanide solution ; to LABORATORY EXERCISES. 55 another part add a few drops of potassium ferrocyanide solution. Notice the difference. To a third portion of the solution add a few drops of potassium sulphocyanide solution. 211. Ferric salts. To a portion of the solution made in the preced- ing experiment and reserved for this experiment, add a few drops of nitric acid, and boil. Test this solution, as in the preceding experi- ment and record the difference in action as compared with the results obtained above. Explain the action of the nitric acid. 212. To a solution of a ferric salt add stannous chloride solution, and warm. Ascertain whether the iron is in the ferric or the ferrous condition. How may ferrous salts be changed to ferric salts, and vice versa ? In what ways can ferric and ferrous salts be distinguished from one another ? What is the most delicate test for each class of these compounds ? EXERCISE LV. ANALYSIS; QUALITATIVE. Substances may be subjected to analysis with two ends in view. 1. To ascertain the kind of matter present (Qualitative). 2. To ascertain the quantity of each kind of matter (Quantita- tive). Review the experiments recorded in your notebook and make a list of those in which you have analyzed qualitatively any substance. Do the same with experiments in which quantitative results have been obtained. In qualitative analysis both physical and chemical changes may be used as tests. (a) Physical tests. Such tests are illustrated in Exp. 165. Try as there described for sodium, a solution of a lithium salt, also of a strontium salt. State the difference. (6) Chemical tests. Such tests have been illustrated several times during the study of the metals. Look over your previous work and select at least three such cases. These may be further illustrated by the following : 213. In a test tube place about 5 cc. of silver nitrate solution and a like quantity of copper nitrate solution. Shake well. Proceed now as in 196. What is the precipitate in the first case? What remains in the filtrate ? What is the precipitate in the second case ? How could you have distinguished a lead precipitate from one of silver in the first instance ? (See 183.) 56 LABORATORY EXERCISES. 214. A Qualitative Analysis. Put a ten-cent silver coin into an evaporating dish, and pour over it a mixture of 5 cc. HN0 3 and 10 cc. H 2 0. Warm it till all or nearly all has dissolved. Remove any that is undissolved, and pour the liquid into a test tube. Add HC1 as long as a precipitate continues to form, then filter. AgCl is the residue. Give the reaction. Add a drop or two of HC1 to the filtrate, and, if a precipitate falls, add more, and filter again, to remove all the Ag. Evaporate the filtrate to a few drops in an evaporating dish (to remove any free HN0 3 ), then add H 2 and pass H 2 S gas into the filtrate so long as a precipitate forms. This is CuS. Write the reaction. Filter. The coin is thus found to contain Ag and Cu. EXERCISE LVI. A QUANTITATIVE ANALYSIS. In the previous experiments what substances have been shown to be present in common salt? In what experiments and how was this shown ? By what kind of analysis was it shown ? 215. In a perfectly dry and clean mortar pulverize some pure so- dium chloride. Put it in an evaporating dish and heat it over the Bunsen flame about 5 minutes. Cool, and weight out accurately on the balance 1 g. of the salt. Dissolve it in a beaker by means of about 25 cc. of distilled water. Warm the solution and add as long as a precipitate is formed a silver nitrate solution, pouring it gradually and with constant stirring. The contents of the beaker should be protected from the direct sun light as much as possible. Allow the precipitate to settle, and decant the supernatant liquid through a well-dried and weighed filter held in a funnel. Care should be taken to make sure that all the chlorine has been precipitated. Finally, transfer the whole of the precipitate to the filter paper, being careful not to lose even the least particle. Wash the precipitate on the filter thoroughly. Dry on the filter paper to constant weight, being careful not to scorch the paper. The increase in weight is the weight of the silver chloride. Calculate now the per cent of chlorine in silver chloride, AgN0 3 , and, using this as a factor, calculate the per cent of chlorine in the sodium chloride used, and also the per cent of sodium. APPENDIX. THE METRIC SYSTEM OF WEIGHTS AND MEASURES. 1 The metric system, employed in the affairs of everyday life by most of the nations of continental Europe, and by scientific writers throughout the world, is based upon a fundamental unit, or measure of length, called a meter. This meter is defined as the 40-millionth part of the circumference of the earth, or, in other words, of a " great circle," or meridian. Its length was originally determined by actual measurement of a considerable arc of a meridian, but the various measurements heretofore made of the length of the earth's meridian differ slightly from each other ; and it is to be expected, and indeed hoped, that the steady improvements of methods and instruments will make each successive determination of the length of the meridian better than, and therefore different from, the preceding. It is on this account necessary to define the standard of length by legislation to be a certain rod of metal, deposited in a certain place, under specified guaranties, and to secure the uniformity and permanence of the standard by the multiplication of exact copies in safe places of deposit. From this single quantity, the meter, all other measures are deci- mally derived. Multiplied or divided by 10, 100, 1000, and so forth, the meter supplies all needed linear measures ; and the square meter and cubic meter, with their decimal multiples, supply all needed measures of area or surface on the one hand, and of solidity or capacity on the other. From the unit of measure to the unit of weight, the transition is admirably simple and convenient. The cube of the one hundredth of the linear meter is, of course, the millionth of the cubic meter : its bulk is about that of a large die of the common backgammon board. This little cube of pure water is the universal unit of weight, a gram, which, decimally multiplied and divided, is made to express all weights. The numbers expressing all weights, from the least to the greatest, find direct expression in the decimal notation ; the weights used in different trades only differ from each other in being different iFrom Manual of Chemistry, by Storer and Lindsay. 67 58 LABORATORY EXERCISES. decimal multiples of the same fundamental unit ; and, in comparing together weights and volumes, none but easy decimal computations are ever necessary. The nomenclature of the metrical system is extremely simple ; one general principle applies to each of the following tables. The Greek prefixes for 10, 100, and 1000, viz., deca, hecto, and kilo, are used to signify multiplication; while the Latin prefixes for 10, 100, and 1000, viz., deci, centi, and milli, are employed to express subdivision. Of the names thus systematically derived from that of the unit in each table, many are not often used ; the names in common use are those printed in small capitals. Thus, in the table for linear measure, only the meter, kilometer, centimeter, and millimeter are in common use, the first for such purposes as the English yard subserves, the second instead of the English mile, the third and fourth in lieu of the fractions of the English foot and inch. Divisions Unit . . Multiples Divisions Unit . . Divisions Unit . . Multiples LINEAR MEASURE. Meters. f MILLIMETER 0.001 or 1-1000 of a meter. j CENTIMETER 0.01 or 1-100 [ Decimeter 0.1 or 1-10 METER 1. f Decameter 10. \ Hectometer = 100. I KILOMETER = 1000. SURFACE MEASURE. f Millimeter square = 0.000,001 of a meter square. | Centimeter square = 0.000,1 " " I Decimeter square = 0.01 " " METER SQUARE = 1. f Cubic Millimeter = j Cubic- Centimeter = [ Cubic Decimeter = CUBIC METER = C Cubic Decameter = I Cubic Hectometer = CUBIC MEASURE. Cubic Meters. 0.000,000,001 0.000,001 0.001 1. 1,000. 1,000,000. i Cubic Kilometer =1,000,000,000. APPENDIX. 59 The table for land measure we omit, as having no connection with our subject. For the measurement of wine, beer, oil, grain, and simi- lar wet and dry substances, a smaller unit than the cubic meter is desirable. The cubic decimeter has been selected as a special stand- ard of capacity for the measurement of substances, such as are bought and sold by the English wet and dry measures. The cubic decimeter (1000 cubic centimeters) thus used is called a liter. CAPACITY MEASURES. Liters. Cubic Meter. f Milliliter = 0.001 = 0.000,001 = 1 cubic centimeter. Divisions . | Centiliter = 0.01 = 0.000,01 [Deciliter = 0.1 =0.000,1 Unit . . LITER = 1. = 0.001 = 1 cubic-decimeter. f Decaliter = 10. =0.01 Multiples . ] HECTOLITER = 100. =0.1 iKiloliter =1,000. =1. = 1 cubic meter. The table of weights bears an intimate relation to this table of capacity. As already mentioned, the weight of that die-sized cube, a cubic centimeter, or milliliter, of distilled water (taken at 4, its point of greatest density), constitutes the metric unit of weight. This weight is called a gram. From the very definition of the gram, and from the table of capacity measure, it is clear that a liter of dis- tilled water at 4 will weigh 1,000 grams. WEIGHTS. Grams. f MILLIGRAM = 0.001 Divisions . | CENTIGRAM = 0.01 [ DECIGRAM = 0.1 Unit . . GRAM = 1. = 1 cubic centimeter of water at 4. f Decagram = 10. Multiples . I Hectogram = 100. I Kilogram = 1,000. = 1 cubic decimeter of water at 4. The simplicity and directness of the relations between weights and volumes in the metric system can now be more fully explained. The chemist ordinarily uses the gram as his unit weight, and, for his unit of volume, a cubic centimeter, which is the bulk of a gram 60 LABORATORY EXERCISES. of water. For coarser work, the kilogram becomes the unit of weight, and the corresponding unit of measure is the liter, which is the bulk of a kilogram of water. In commercial dealings, in manu- facturing processes, and, above all, in scientific investigations, these simple relations between weights and measures have been found to be an inestimable advantage. The numerical expressions for metric weights and measures may always be read as decimals. Thus, 5.126 meters will be read, " five meters and one hundred and twenty-six thousandths," and not " five meters, one decimeter, two centimeters, and six millimeters." The expression " 10.5 grams " is read " ten and five tenths grams " ; just as we say one hundred and five dollars, not ten eagles and five dollars; or sixty-five cents, not six dimes and five cents. All computations under the metric system are made with decimals alone. The abbreviations commonly met with in chemical literature are : mm. for millimeter. cm. for centimeter, m. for meter. cc. or cm 8 for cubic centimeter, g. or grm. for gram. k. or kgm. or kilo, for kilogram. 1. for liter. The equivalents in English weights and measures, of those metric weights and measures which are used in chemistry, can be readily found by the aid of the table on page 63, which is available not only for grams, centimeters, and liters, but, by mere change of the position of the decimal point, for all decimal multiples or subdi- visions of these quantities. One cubic meter = 35.31660 cubic feet. One cubic decimeter (a liter) = 61.02709 cubic inches. One cubic centimeter = 0.06103 cubic inch. One liter = 0.22017 imperial gallon. One liter = 0.88066 imperial quart. One liter = 1.76133 imperial pint. One liter = 0.26427 U. S. gallon. One liter = 1.05708 U. S. quart. One liter = 2.11415 U. S. pints. One gram = 15.4323 grains. One meter = 39.3708 inches. APPENDIX. 61 One pound Avoirdupois = 7,000 grains One pound Troy = 5,760 grains One ounce Avoirdupois = 437.5 grains One ounce Troy = 480 grains One grain One English imperial gallon = 277.274 cu. in. One U. S. standard gallon = 231 cu. in. One U. S. quart One fluid ounce One foot One yard One inch 453.59 g. 373.24 g. 28.35 g. 31.10 g. 64.80 mg. 4.54 1. 3.78 1. 0.95 1. 29.56 cc. 0.3048 m. 0.9144 m. 2.54 cm. LABORATORY EXERCISES. TABLE FOB THE CONVERSION OF DEGREES ON THE CENTIGRADE THERMOMETER INTO DEGREES OF FAHRENHEIT'S SCALE. C. = (F. - 32) f ; and F. = f C. + 32. CENT. FAHR. CENT. FAHR. CENT. FAHR. o -50 -58.0 17 62.6 o 60 140.0 -45 -49.0 18 64.4 61 141.8 -40 -40.0 19 66.2 62 143.6 -35 -31.0 20 68.0 63 146.4 -30 -22.0 21 69.8 64 147.2 -25 -13.0 22 71.6 65 149.0 -20 - 4.0 23 73.4 66 150.8 -19 - 2.2 24 75.2 67 152.6 -18 - 0.4 25 77.0 68 154.4 -17 + 1.4 26 78.8 69 156.2 -16 3.2 27 80.6 70 158.0 -15 5.0 28 82.4 71 159.8 -14 6.8 29 84.2 72 161.6 -13 8.6 30 86.0 73 163.4 -12 10.4 31 87.8 74 165.2 -11 12.2 32 89.6 75 167.0 -10 14.0 33 91.4 76 168.8 - 9 15.8 34 93.2 77 170.6 - 8 17.6 35 95.0 78 172.4 - 7 19.4 36 96.8 79 174.2 - 6 21.2 37 98.6 80 176.0 - 5 23.0 38 100.4 81 177.8 - 4 24.8 39 102.2 82 179.6 - 3 26.6 40 104.0 83 181.4 - 2 28.4 41 105.8 84 183.2 - 1 30.2 42 107.6 85 185.0 32.0 43 109.4 86 186.8 + 1 33.8 44 111.2 87 188.6 2 35.6 46 113.0 88 190.4 3 37.4 46 114.8 89 192.2 4 39.2 47 116.6 90 194.0 5 41.0 48 118.4 91 195.8 6 42.8 49 120.2 92 197.6 7 44.6 60 122.0 93 199.4 8 46.4 61 123.8 94 201.2 9 48.2 62 125.6 95 203.0 10 60.0 63 127.4 96 204.8 11 51.8 64 129.2 97 206.6 12 53.6 55 131.0 98 208.4 13 55.4 66 132.8 99 210.2 14 57.2 67 134.6 100 212.0 15 69.0 58 136.4 16 60.8 69 138.2 APPENDIX. 63 -E - -E - -- -- ^rt __ L g __ -- i^ -- -- Cl _ _ _ -E - -) a STORER AND LINDSAY'S Elementary Manual of Chemistry BY F. H. STORER, S.B., A.M., and W. B. LINDSAY, A.B., B.S. Cloth, 12mo, 453 pages. Illustrated. Price, $1.20 This work is the lineal descendant of the " Manual of Inorganic Chemistry" of Eliot and Storer, and the "Ele- mentary Manual of Chemistry " of Eliot, Storer and Nichols. It is in fact the last named book thoroughly revised, rewritten and enlarged to represent the present condition of chemical knowledge and to meet the demands of American teachers for a class book on Chemistry, at once scientific in statement and clear in method. The purpose of the book is to facilitate the study and teaching of Chemistry by the experimental and inductive method. It presents the leading facts and theories of the science in such simple and concise manner that they can be readily understood and applied by the student. The book is equally valuable in the classroom and the laboratory. The instructor will find in it the essentials of chemical science developed in easy and appropriate sequence, its facts and generalizations expressed accurately and scientifi- cally as well as clearly, forcibly and elegantly. ' ' It is safe to say that no text-book has exerted so wide an influence on the study of chemistry in this country as this work, originally written by Eliot and Storer. Its distinguished authors were leaders in teaching Chemistry as a means of mental training in general edu- cation, and in organizing and per- fecting a system of instructing students in large classes by the experimental method. As revised and improved by Professor Nichols, it continued to give the highest satisfaction in our best schools and colleges. After the death of Pro- fessor Nichols, when it became necessary to revise the work again, Professor Lindsay, of Dickinson College, was selected to assist Dr. Storer in the work. The present edition has been entirely rewritten by them, following throughout the same plan and arrangement of the previous editions, which have been so highly approved by a generation of scholars and teachers. " If a book, like an individual, has a history, certainly the record of this one, covering a period of nearly thirty years, is of the highest and most honorable character." From The American Journal of Science. Copies cf this book will be sent prepaid to any address^ on receipt of the price^ by the Publishers : American Book Company New York Qefli*i^. * Chicago (161) * UNIVERSITY *;~2Lmi YB 16873 107459