m \ vv>at4 14 :. jy UNIVERSITY OF CALIFORNIA DEPARTMENT OF EDUCATION GIFT OF THE PUBLISHER /?<** No ' * Received LIBRARY OF THE UNIVERSITY OF CALIFORNIA. GIFT OF Class An Elementary Experimental Chemistry BY JOHN BERNARD EKELEY, A.M. * Science Master at St. Paul's School, Garden City SILVER, BURDETT & COMPANY NEW YORK. BOSTON CHICAGO COPYRIGHT, 1900, BY SILVER. BURDETT & COMPANY PREFACE THERE are two things that the study of a science should accomplish for the student : first, the development of the powers of observation ; second, a knowledge of the relations of the principles and facts that underlie that science. The first may be obtained to a greater or less degree from the study of any science by the experimental method ; the lat- ter however may often be only imperfectly acquired even by those who are able to make accurate observations, but who fail to draw the correct or logical conclusions. It is the aim of this book to aid the student in accomplishing both these things. The author has found by experience (as, in- deed, what teacher has not ?) that all beginners in the study of chemistry are by no means natural adepts at making the necessary observations, but that by a sufficiently prolonged course of experiments their powers in this direction are easily augmented. The difficulty comes, however, in ena- bling the, student, after the observations have been made, to form the correct conclusions, and, finally, when both these things have been done throughout the subject, to round it off into a symmetrical whole. To accomplish these things, the author believes that it is necessary to help the student considerably by emphasizing what must especially be marked in the experiment, and to outline the course of thought that must be followed, in order that the student may feel satisfied at the close of the experi- 235384 IV PREFACE ment that he has made some definite progress at least in the process of building up his knowledge of the subject. In Parts I. and II., this method is followed. Part I. is made up entirely of experiments dealing with The Prepara- tion and Properties of the Principal Elements and Compounds. No teaching of chemical theory whatever is attempted in this part. Qualitative equations are given, showing merely what kinds of matter have been concerned in each chemi- cal change. Words, and not chemical symbols, are used, the conventional chemical equation being left for Part II. It is hoped that the student, when he has performed the ex- periments of Part I., may have a good idea of the qualitative composition of the principal compounds, and of how he may distinguish them from one another. Each experiment gives him a knowledge of one or more elements or compounds ; in doing this, no compound is used that has not been studied in some previous experiment. Thus the student's knowledge is built up step by step by the inductive method, and in a logi- cal manner, until he has made, and studied the properties of those elements and compounds of inorganic chemistry with which he is most concerned. The author believes that the study of chemical theory is most successfully carried on after the student has prepared the elements and compounds and has studied their proper- ties. Hence Part II. is concerned entirely with The Laws and Theories of Chemistry. Special emphasis is laid upon the difference between laws and theories. The laws are illus- trated by experiments; and the reasoning upon which we base our belief in the respective explanatory theories, which most students find so difficult, is given in full. This gives an accurate idea of theoretical chemistry instead of the hazy conceptions that are so often obtained when the theory is scattered throughout a series of merely descriptive experi- PREFACE V ments. The use of chemical symbols and the writing of reactions are now taken up in detail, and numerous exam- ples are given to illustrate the stoichiometrical relations of the elements. The plan of deferring the writing of reactions until late in the course is in line with the most recent sug- gestions of the best teachers of the science. Throughout the book the revised spelling of chemical names is used. The illustrations of apparatus in Parts I. and II. will save the teacher valuable time that would otherwise be used in answering questions on manipulation. The author wishes to acknowledge his indebtedness to Prof. J. F. McGregory of Colgate University, and to Dr. Albert C. Hale of the Boys' High School, Brooklyn, for many valuable suggestions, and to his colleague at St. Paul's, Mr. Arthur De L. Ayrault, for his painstaking criticism of the English. JOHN B. EKELEY. INTRODUCTION TO THE TEACHER To use this book, especially Parts I. and II., to the best advantage, the order of the experiments, in most cases, should not be changed. Great care has been taken to arrange them so that they follow one another in logical order. If there is a lack of time to do them all, some (the teacher will easily see what ones) can be omitted without destroying the continuity of the whole. Students should be examined by means of recitations or individual examinations after each ten experiments of Part I. At St. Paul's, each student meets the instructor for a " quiz " as soon as ten consecutive experiments have been finished. With Part II., however, it seems best to supple- ment the experimental work with frequent recitations. TO THE STUDENT Procure a blank book of about one hundred and fifty pages, with durable covers (leather back and corners). Begin your notes on the sixth page, leaving the first five pages blank for a future " table of contents." On the left-hand page, record with a hard pencil the notes on your observations as taken at the time of the experiment. Do not fall into the habit of writing your original notes on slips of paper, but record them immediately in your note book. When your Vlll INTRODUCTION original notes have been criticised by your instructor, study them carefully. After you are satisfied that you thoroughly understand the experiment, write out in ink on the right- hand page in your best English a detailed description of the experiment and the conclusions that you have reached. Before attempting an experiment, study carefully the de- scription as given in the text, so that you are sure you un- derstand what you are going to do, and what the object of the experiment is. Keep your laboratory desk and apparatus clean. You will often be tempted not to do this, but you will get the best results from your work, and get the most enjoyment from it, if you avoid slovenly habits of experimentation. CONTENTS PAGE Introduction , . ... vii PART I. Preparation and Properties of Elements and Compounds . . . 3-85 Exp. i. Three Conditions of Matter ... 3 " 2. Physical and Chemical Changes 3 " 3. Mechanical Mixture and Chemical Compound .... 4 " 4. Copper and Copper Oxid ... 5 " 5. Mercury 6 " 6. Oxygen 7 " 7. Phosphorus 10 8. Carbon 12 " 9. Sulfur 13 " 10. Sodium 17 " n. Potassium 18 " 12. Zinc .... 19 " 13. Magnesium 20 " 14. Iron . . . . ... 20 " 15. Hydrogen 21 " 1 6. Hydrogen and Oxygen 25 " 17. Zinc and Magnesium Oxids, and Sulfuric Acid . ... 27 " 1 8. Neutralization 28 " 19. Carbonic Acid and Carbon Dioxid with Sodium and Potassium Hydroxid 30 Calcium 30 Analysis of Marble 32 22. Chlorin ... 34 23. Hydrochloric Acid from Sodium Chlorid and Sulfuric Acid 36 X CONTENTS PAGE Exp. 24. Analysis of Hydrochloric Acid 37 " 25. Hydrochloric Acid with Metals, Hydroxids, and Car- bonates 38 " 26. Carbon Dioxid from Marble and Hydrochloric Acid . . 39 " 27. Preparation of Chlorin 40 " 28. Bromin 41 " 29. lodin 43 " 30. Potassium lodid 44 " 31. Calcium Fluorid and Hydrogen Fluorid ...... 45 " 32. Sulfids 47 " 33- Other Compounds of Carbon 49 " 34. Nature of Flame 54 " 35. Hard and Soft Water 56 " 36. Nitrogen 57 " 37. Nitric Acid 59 " 38. Neutralization of Nitric Acid 62 " 39. Nitric Oxid 63 " 40. Elements in the Nascent State 65 " 41. Ammonia 66 " 42. Ammonium Chlorid . .' 67 " 43. Ammonia from Ammonium Chlorid 68 " 44. Ammonium Amalgam 69 " 45. Neutralization of Acids with Ammonium Hydroxid . . 69 " 46. Nitrous Oxid 70 " 47. Analysis of Nitrous Oxid 72 " 48. Arsenic 72 " 49. Antimony 74 " 50. Bismuth 75 " 51. Cadmium 76 " 52. Mercury 77 " 53- Lead 77 " 54- Tin 79 " 55. Aluminum 79 " 56. Iron * 80 " 57. Nickel 8 1 " 58. Barium 81 " 59. Strontium 82 " 60. Silver 83 " 61. Gold 84 " 62. Platinum 85 CONTENTS XI PART II. PAGE Laws and Theories of Chemistry 89-156 Exp. i. The Sum of the Weights of the Factors in a Chemi- cal Change equals the Sum of the Weights of the Products 90 " 2. Law of Definite Proportions by Weight 92 " 3. Law of Multiple Proportions 93 " 4. Quantitative Analysis of Sodium Chlorid 94 " 5. Barometer 98 " 6. Law of Boyle , 99 " 7. Law of Charles 101 " 8. Weight of a Liter of Air 104 " 9. Weight of a Liter of Carbon Dioxid ...... 105 " 10. Weight of a Liter of Hydrogen 106 " 1 1. Density 107 " 12. Density of a Liquid in the form of Vapor 107 " 13. Chemical Equivalence 110 ' 14. Electrical Equivalents 112 " 15. Specific Heat of Lead 114 ' 16. Specific Heat of Iron 1 16 ' 17. Law of Definite Proportions by Volume 119 " 1 8. Molecular Weight of Potassium Chlorate 123 " 19. Molecular Weight of Potassium Chlorid 124 " 20. Writing of Reactions 141 " 21. Heat of Chemical Action . . 147 " 22. Heat of Neutralization 149 " 23. Heat of Solution and of Hydration 150 " 24. Dissociation (Gaseous) 152 " 25. Dissociation In Liquids 153 " 26. Burning of Organic Matter ; Dry Distillation .... 154 " 27. Alcohol 155 " 28. Saponification . 156 Xll CONTENTS PART III. PAGE History, Occurrence and Industrial Applications of the Principal Elements and Compounds 159-223 Qualitative Analysis 227-239 Appendix I 240 Appendix II 241 PART I. PREPARATION AND PROPERTIES OF ELEMENTS AND COMPOUNDS AN ELEMENTARY EXPERIMENTAL CHEMISTRY PART I. PREPARATION AND PROPERTIES OF ELEMENTS AND COMPOUNDS EXPERIMENT i Three Conditions of Matter TAKE a piece of ice. Note its properties. After crush- ing, heat in a 100 cc. flask. When all is melted, note its properties. Heat the water obtained to boiling. Does any- thing escape from the mouth of the flask ? If so, state its properties. What are the three conditions of matter ? EXPERIMENT 2 Physical and Chemical Changes a. Dissolve a little sugar in water. Evaporate until all the water is gone. What remains ? b. Heat a little sugar in a porcelain dish. Examine the residue. Is it still sugar ? 3 .. XPERIMENTAL CHEMISTRY c. Heat an iron wire in a flame. Examine it again after it is cool. Is it still iron ? d. Have ready a few small iron nails. Pour 10 cc. of water into a 100 cc. flask, and then add 5 cc. of sulfuric acid. In diluting sulfuric acid, always pour the acid into the water. If you pour the water into the acid, it will spatter. Dangerous. Place the- nails in the diluted acid; and, after all action has ceased, filter, evaporate, and allow to crystallize. Is the resulting product iron ? State what you conceive to be the difference between physical and chemical change. EXPERIMENT 3 Mechanical Mixture and Chemical Compound An element is a substance that has nof been separated into two or more dissimilar substances. Sulfur and iron are elements. a. Mix equal portions of flowers of sulfur and fine iron filings. Can you distinguish the particles of each element on looking at the mixture through a strong magnifying glass ? Pass a magnet through the mixture, and tap it lightly on some solid body. Place a little of the mixture in a test tube, and add a little carbon disulfid. After shaking well for a few moments, pour off the liquid and examine the residue. Is it iron ? Evaporate the liquid. What remains ? b. Place the remainder of the mixture in a small test tube, and hold the bottom of the tube in the flame of i Bunsen burner. After the action has ceased, break thj ELEMENTS AND COMPOUNDS 5 tube, and examine the contents by the same tests as before heating. Is it iron ? Is it sulfur ? State what you conceive to be the difference between a mechanical mixture and a chemical compound. EXPERIMENT 4 Copper and Copper Oxid a. Examine some copper in the form of wire and sheet. State as many of its properties as you can ; i.e., color, hard- ness, luster, weight, tenacity, fusibility, volatility, etc. b. Place about 5 grms. of fine copper filings in a porce- lain crucible, and weigh carefully to i c. grm. Heat over a Bunsen burner, stirring occasionally with a clean iron rod. After it is cool, weigh again. Let us explain the in- crease in weight. Since matter is the only thing that has weight, some kind of matter must have been added to the con- tents of the crucible. It has not been added to the crucible itself; for, by trying, you would find that the weight remains unchanged. This matter could not have come from the heat, since heat is not material ; so it must have come from the air. This something chemists call oxygen. Examine the contents of the crucible ; you will see that it is no longer copper. A chemical change has taken place, and a new substance called copper oxid has 6 AN ELEMENTARY EXPERIMENTAL CHEMISTRY been formed. It is thus named to indicate the chemical union of copper and oxygen. When elements unite chemically, the resulting substance is called a compound. The process of combining two or more elements chemically is called synthesis. We may ex- press what has taken place by the following : *+<* Understand that, when the names of elements are written one under another, the combination signifies that the ele- ments are chemically combined. EXPERIMENT 5 Mercury a. Properties. Examine a small quantity of the element mercury. State as many of its properties as you can dis- tinguish. In heating mercury, be careful not to breathe any of the vapor. By heating mercury a long time in a flask, it is possible to obtain a red powder that weighs more than the original mercury. This red powder is called mercury oxid, and is formed by the union of mer- cury and oxygen just as copper oxid is formed by the union of copper and oxygen. b. Analysis of Mercury Oxid. Place a small quantity of mercury oxid in a hard glass tube of about 15 mm. diameter ELEMENTS AND COMPOUNDS 7 and closed at one end. Fit the tube with a one-holed stopper through which extends a glass exit tube of about 4 mm. bore. By means of a piece of rubber tubing, con- nect this with a glass delivery tube. Clamp the hard glass tube on a standard so that it is in a slightly inclined posi- tion. Have ready a pneumatic trough with a small beaker full of water inverted on the shelf. Now heat the mercury oxid gently. Gradually increase the heat, and allow the bubbles to escape under the inverted beaker of water. What forms on the sides of the tube ? If mercury oxid is composed of mercury and oxygen, what is the substance in the inverted beaker ? This process, the separating of a compound into two or more elements, is called analysis. Mercury _ = Mercury + Oxygen Oxygen EXPERIMENT 6 Oxygen a. Properties. Examine the substance in the jar made in the last experiment. Place a ground-glass plate under the jar, and remove it from the water. Has the gas formed color, odor, " or taste ? Light a pine splinter, blow out the flame so that a spark remains, and plunge the splinter into the jar. What happens ? \Vhat is the substance present in the air, that makes fuel burn ? Why is it that a fire is smothered when air is kept away from it ? b. Preparation from Manganese Dioxid. Perform an experiment similar to Experiment 5 b, using manganese dioxid (composed of manganese and oxygen) instead of mercury oxid. Is a metal left as in experiment 5 b ? 8 AN ELEMENTARY EXPERIMENTAL CHEMISTRY c. Preparation from Potassium Chlorate. Perform a sim- ilar experiment, using potassium chlorate instead of mer- cury oxid. Heat gently at first, beginning at the top of the chlorate, and be careful to keep the hands from under the tube in case it should break ; for melted potassium chlorate makes a serious wound. After the action has ceased, dis- solve the remaining substance in hot water, and evaporate to dryness. Preserve this residue, and, when you have per- formed Experiment 25 c, compare. d. Preparation from a Mixture of Potassium Chlorate and Manganese Dioxid. Perform a similar experiment, using 8 grms. of potassium chlorate and 2 grms. of manganese dioxid. Mix the two thoroughly. Begin heating at the top of the mixture. When action has ceased, add hot water and transfer to a beaker. Be sure to get all the residue. Weigh a filter paper carefully. Filter the hot liquid, being careful to throw it all and the black residue upon the filter paper. Wash with hot water from a wash bottle five or six times. Dry the filter in an oven at a tem- perature of about 100. Weigh again. Subtract the weight of the filter paper from the weight of the filter paper and black residue. How does the result compare with the ori- ginal weight of the manganese dioxid ? Has the manganese dioxid undergone any change chemically ? Evaporate the filtrate (the liquid that ran through the paper) to dryness in a weighed porcelain dish. Weigh, and account for the loss of weight of the chlorate. e. Preparation by Electrolysis of Water. Have ready a vessel fitted with two platinum electrodes. Ino this pour sufficient water acidified with sulfuric acid (i part acid to 20 of water),* so that it will somewhat more than cover the electrodes. Fill two glass tubes graduated in cubic centi- * The acid is only added to aid the water in conducting the current. ELEMENTS AND COMPOUNDS 9 meters with some of the water, and invert them in the vessel so that the mouth of each tube covers an electrode. Place the vessel in circuit with a battery of two bichromate cells. Note the evolution of gas at each electrode. After a few minutes, disconnect the cells, and measure the relative vol- umes of the gases. Test the lesser one with a glowing splinter. What is it ? Apply a lighted match to the greater, and note the result. Here we have a new gas, which chemists have been unable to decompose into anything else. It has been named hydrogen. We now have experimental proof as to the composition of water. IO AN ELEMENTARY EXPERIMENTAL CHEMISTRY Remark. If a silent discharge of electricity is passed through dry oxygen, the gas will decrease in volume one- third, and will acquire greater chemical activity. Some change, then, has occurred in the gas without the addition of any other kind of matter. When elements can thus be changed by any means, they are said to exist in allotropic forms. This second form of oxygen is called ozone. It has a faint peculiar odor, often noticed near electric machines when operating, or just after a discharge of lightning. EXPERIMENT 7 Phosphorus Phosphorus occurs in two allotropic forms, the yellow and the red. The yellow must be stored under water, as it ignites at a temperature of 40 in the air. a. Properties. Examine both red and yellow phosphorus. Be very careful in handling the yellow variety, since it ignites so easily, and a burn from it heals with difficulty. Never handle it with the fingers, and always cut it under water. State as many properties as you can of each kind. b. Preparation of Phos- phoric Oxid by Burning Phosphorus in air. Have ready a quick-sealing fruit jar. The rubber washer should be greased with vaseline. On a support in the jar, hang a deflagrating spoon con- taining a piece of ignited* phosphorus about the size of * The phosphorus may be placed in the jar, and then ignited by means of a strong lens. ELEMENTS AND COMPOUNDS I I a pea. Close the jar air tight. Note what forms in the jar. This product is called phosphoric oxid. Open the jar under water. Was any of the air in the jar used up ? Does the oxid dissolve in water ? Phosphorus + Oxygen - Remark. Another oxid of phosphorus exists, called phos- phorous oxid. It is formed when phosphorus is burned in an insufficient supply of air. It is a white powder with an odor resembling garlic. When heated in the air, it becomes the oxid you have just made ; this shows that it differs from phosphoric oxid, in that it contains less oxygen. c. Preparation of Phosphoric Oxid by burning Phosphorus in Oxygen. Perform a similar experiment, using a jar of oxygen instead of air. In this case, it is better to ignite the phos- phorus in the jar by means of a burning glass. How does the product compare with that formed in b ? d. Phosphoric Acid. Dissolve the oxid prepared in c in a small amount of water. Taste a drop of the solution. Also note its effect upon blue litmus paper. This product is called phosphoric acid ; and, since it is made of phosphoric oxid and water, it can contain only the elements hydrogen, phosphorus, and oxygen. Hydrogen Phosphorus * + ' = Phosphorus Oxygen Oxygen Remark. When phosphorus is burned in an insufficient supply of moist air, another acid of phosphorus is formed which has a faint odor like garlic. This is called phospho- rous acid. It differs from phosphoric acid in that it contains less oxygen. Note the signification of the endings ic and ous. 12 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Definition of an Acid. For the present it will be sufficient to define an acid as a compound of hydrogen, or else of hydrogen and oxygen, with a non-metallic element, which compound has usually a sour taste and turns blue litmus paper red. EXPERIMENT 8 Carbon a. Properties. Examine, and state the properties of, the following allotropic forms of carbon: charcoal, gas carbon, bone black, graphite, and soot. b. Carbon Dioxid. Burn a piece of charcoal in a defla- grating spoon in a jar of oxygen, as in Experiment 7 b. * The charcoal must be well ignited before being placed in the jar. Open the jar under water. Does the cover stick ? What can you say of the volume of the resulting carbonic oxid compared with the original volume of oxygen ? Close the jar, allowing a little water to enter, and shake vigorously. Open again under water. Is carbonic oxid soluble in water ? In the same way, prepare another jar of the gas. Plunge a lighted splinter into the gas. Does the gas burn? Does the splinter continue to burn ? The name carbon dioxid is given to this gas for reasons that will appear later. - Carbon Carbon + Oxygen = Qxygen c. Carbonic Acid. Taste the solution formed in b, and test it with blue litmus paper. Here we have a second acid, which has been named carbonic acid. Transfer to a test tube and heat to boiling. Taste, and test with litmus paper again. Is carbonic acid a stable compound, i. e., is it hard to de- compose ? What has become of the carbon dioxid ? ELEMENTS AND COMPOUNDS 13 Carbon Hydrogen = Oxygen Oxygen Remark. In Experiment 7 b, and 8 , we have observed the phenomenon of combustion. Combustion is the chemical union of substances, accompanied by the evolution of light and heat. The substances that unite may be any substances whatever, but ordinarily we apply the term to the union of substances with the oxygen of the air. The substance that unites with the oxygen is called the combustible, while the oxygen is said to support combustion. The ease with which elements unite with oxygen varies. Some, such as phospho- rus, require only a slight rise in temperature, w r hile others, like carbon, require a comparatively high one. The temper- ature at which a substance takes fire in air is called its kindling temperature. This temperature is constant for any particular substance. EXPERIMENT 9 Sulfur a. Properties. After examining, state as many properties as you can of roll sulfur. Place about 10 grms. in a test tube, and heat over a Bunsen burner, heating gently at first and gradually increasing the temperature. Note carefully the changes that the sulfur passes through, especially in re- gard to its color and consistency. Note also the ease with which sulfur takes fire. b. Allotropic Forms of Sulfur. i. In a beaker put about 60 grms. of roll sulfur, and heat gently until all is melted. When all is melted, allow it to 14 AN ELEMENTARY EXPERIMENTAL CHEMISTRY cool. As soon as crystals have just covered the surface of the liquid, break a hole in the crust, and quickly pour out the melted sulfur. Note the shape and color of the crystals that remain in the beaker. Allow it to stand for a few days, and examine again. 2. Pulverize a small lump of roll sulfur, and shake it up with carbon bisulfid in a test tube. Keep away from a flame, since carbon bisulfid is very inflammable. When it is dissolved, pour it into an evaporating dish and allow it to crystallize. Note the color and shape of the crystals formed. 3. In a test tube, heat to boiling about 15 grms. of roll sulfur, and quickly pour the liquid into a beaker of cold water. State the properties of the product. Allow it to stand for a few days, and examine again. c. Sulfur Dioxid. Burn a small amount of sulfur in a deflagrating spoon in a jar of oxygen. Note the color of the flame. After action ceases, open under water. Does the cover stick? What can you say of the volume of the gas compared with the original volume of oxygen ? Get its odor and color, if any. Allow a little water to enter the jar, seal again, and shake. Open again under water. Does the cover stick ? Is the gas soluble in water ? We shall call this gas sulfur dioxid for reasons that will appear later. Sulfur Sulfur + Oxygen = _ Oxygen d. Sulfurous Acid. Taste the solution made in b, and test it with blue litmus paper. This is called sulfurous acid. Sulfur Oxygen Oxygen e. Sulfur Trioxid. Have ready a gas holder full of oxy- gen. Apply to the instructor for an apparatus for obtaining ELEMENTS AND COMPOUNDS 15 a steady stream of sulphur dioxid. (See g.) By means of rubber tubing and a glass Y tube, allow the oxygen and sulfur dioxid to come together. In order to dry the mixed gases, pass them by means of rubber tubing through a catch bottle containing concentrated sulfuric acid, which has the property of extracting moisture from gases as they bubble through it. Then pass the dry gases through a hard glass tube, about 15 cm. long and i cm. bore, containing platinized asbestos. Do not pack the asbestos in the tube, but place it in loosely. Have the tube containing the asbestos at- tached to a test tube fitted with a two-holed stopper contain- ing an entrance and an exit tube. Let the entrance tube extend down to within two centimeters of the bottom. Pack the test tube in a mixture of ice and salt. Now heat the asbestos to redness. The red-hot platinized asbestos simply aids in the chemical union of the two gases. After three or four minutes, remove the test tube from the ice. What is 1 6 AN ELEMENTARY EXPERIMENTAL CHEMISTRY the nature of the crystals formed ? Note especially the effect of warmth upon them. This product is called sulfur trioxid. Since it is formed by the chemical union of sulfur dioxid and oxygen, it differs from sulfur dioxid in that it contains more oxygen. Sulfur Sulfur Oxygen + Xygen = Oxygen / Sulfuric Acid. Allow a few drops of water to fall upon the oxid made in e. Get the properties of the resulting compound, especially its taste (dilute a little with considerable water be- fore tasting), and action upon litmus paper. This compound is called sul- furic acid, and evidently differs from sulfurous acid in that it contains more oxygen. Examine a little com- mercial sulfuric acid. Oxygen Oxygen Oxygen g. Su Ifu r Dioxid from Sulfuric Acid and Copper. In a 250 cc. ] flask, place about5ogrms. of sheet cop- per clippings. Fit the flask with a two-holed stopper containing an exit tube and a funnel tube that extends down to the copper. Pour concentrated sulfuric acid into the flask till it somewhat ELEMENTS AND COMPOUNDS I/ more than covers the end of the funnel tube. Heat with a Bunsen burner, and collect a jar full of the gas by means of a glass tube passing to the bottom of the jar. Note that it is the same gas as you obtained on burning sulfur in oxygen, i.e., sulfur dioxid. EXPERIMENT 10 Sodium The element sodium must be stored under naphtha, to pre- serve it. When you wish to use any, apply to the instructor for it. Never handle it with the fingers. Use forceps. a. Properties. Examine a piece of clean sodium. Get its properties, hardness, color, etc. Is it a metal ? Make a fresh cut on it with a knife, and note the appearance of the clean surface for a few moments. Throw a small piece upon water, and note the result. b. Sodium Oxid. In a crucible, heat a piece of clean so- dium about the size of a pea. As it burns, note the color of the flame. Get the properties of the resulting sodium oxid. Sodium Sodium + Oxygen = c. Sodium Hydroxid. When the oxid made in b is cooled, allow a few drops of water to fall on it. Note the sputtering sound. Put a drop of the solution on your finger, and note the greasy feeling. Test the liquid with red litmus paper. Evaporate it to dryness, and allow it to stand for some time, after which note that moisture has collected upon the sub- stance. Since sodium oxid has united with water, the resulting compound must be composed of sodium, oxygen, and hydro- 1 8 AN ELEMENTARY EXPERIMENTAL CHEMISTRY gen. It is called sodium hydroxid or sodium hydrate. This is the first of a class of compounds called bases, which you will study. Sodium Sodium Hydrogen Definition of a Base. For the present it is sufficient to de- fine a base as a compound of a metallic element with oxy- gen, or with oxygen and hydrogen. Bases that turn red litmus paper blue are called a/kalis. d. Sodium Amalgam. Weigh out 20 grms. of mercury, and place it in an iron pan under a hood. Heat it to about 200. The temperature may easily be observed by means of a thermometer. Take 2 grms. of clean sodium, and by means of a pair of long forceps drop it into the mercury. Step back instantly, since mercury and sodium unite with great violence, and poisonous vapors of mercury are evolved. The product is called sodium amalgam. EXPERIMENT n Potassium Potassium is stored in a way similar to sodium, and must be handled with equal care. a. Properties. Note the properties of potassium as you did those of sodium. Compare the metal with sodium. b. Potassium Oxid. Perform an experiment with a piece of potassium, corresponding to Experiment 10 b. Note the color of the flame. Potassium Potassium + Oxygen = _ Oxygen ELEMENTS AND COMPOUNDS IQ c. Potassium Hydroxid. Perform an experiment with potassium corresponding to Experiment 10 c. How does the compound here made differ in chemical composition from that made in Experiment 10 ^? Here we have another compound belonging to the class of bases. It is called potassium hydroxid. Potassium Potassium . Hydrogen _ Oxygen + Oxygen = <*?* Hydrogen EXPERIMENT 12 Zinc a. Properties. Examine, and note the properties * of, the metal zinc in its various forms, sheet, stick, granular, and dust. b. Zinc Oxid. Heat a small piece of zinc in a crucible, stirring with an iron rod. Note the color of the flame as the zinc burns, the color of the oxid when hot and when cold, and the peculiar woolly appearance that it assumes. Zinc + Oxygen = _ Oxygen c. Zinc Oxid and Water. Try the effect of water on a little oxid of zinc, and test it with litmus paper turned red with carbonic acid. Filter, and evaporate the filtrate. Is zinc oxid soluble in water ? Does zinc oxid form an acid or a base ? * Zinc is peculiar in that its physical properties vary considerably with heat. At the ordinary temperature, zinc is rather brittle, but, if heated to 100, it becomes malleable and can be rolled into sheets. At 205 it becomes so brittle that it can be powdered. It melts at 400, and boils at about 1000. Granular zinc is formed when the molten metal is allowed to fall drop by drop into cold water. Zinc dust is ob- tained when the vapors of boiling zinc are suddenly condensed in the absence of air. 2O AN ELEMENTARY EXPERIMENTAL CHEMISTRY EXPERIMENT 13 Magnesium a. Properties. Examine the element magnesium both in the form of ribbon and powder (flash light powder). State its properties. b. Magnesium Oxid. Oxidize a little magnesium in a cru- cible, noting the color of the oxid when hot and when cold. How could you distinguish magnesium oxid from zinc oxid ? Magnesium Magnesium + Oxygen = Oxygen c. Magnesium Oxid and Water. Treat magnesium oxid with water as you did zinc oxid in Experiment 12 c. Does mag- nesium oxid form an acid, or a base ? EXPERIMENT 14 Iron a. Properties. Examine, and state the properties of, iron both in the form of nails and in the form of powder (iron by hydrogen). b. Iron Oxid. (Burning iron in air.) Place about 10 grms. of " iron by hydrogen " in a crucible, and heat, stir- ring occasionally. The resulting compound is evidently iron oxid. Iron Iron -f Oxygen = _ Oxygen c. Iron Oxid and Water. Try the effect of water upon iron oxid. d. Iron Oxid. (Burning iron in oxygen.) Fill a jar with ELEMENTS AND COMPOUNDS 21 oxygen gas. Pierce a hole through a flat cork large enough to cover the mouth of the jar. Through this hole, insert a watch spring from which the temper has been taken by heating in the flame of a Bunsen burner. Around one end of the spring, wind a little cotton string, and soak this in melted sulfur. Ignite the sulfur, and plunge it into the jar of oxygen, making the cork cover the mouth of the jar. As the spring is consumed, feed it through the hole until action ceases. What compound is formed ? EXPERIMENT 15 Hydrogen a. Hydrogen by Electrolysis of Water. Properties. Make hydrogen again by electrolysis of water, as in Experiment 6 e. Note its properties. b. Hydrogen from Steam and Hot Iron. Place about 25 or 30 grms. of dean iron filings in the middle of a piece of half- inch gas pipe about two feet long. Be careful that the filings do not stop up the pipe. Fit the ends with one-holed stop- pers containing glass tubes about 5 mm. bore. To one end of the pipe connect a 500 cc. flask about half full of water, and to the other a rubber tube fitted with a glass leading tube. Heat the gas pipe red hot with a blast lamp, and then boil 22 AN ELEMENTARY EXPERIMENTAL CHEMISTRY the water in the flask. Collect the gas that issues from the leading tube in test tubes over water. Compare the gas with the hydrogen obtained in a. Be careful to have no flame near the leading tube, since hydrogen mixed with air is an explosive compound. Let us see what has taken place. We know from Experi- ment 6 e that water is composed of oxygen and hydrogen, and from Experiment \\b we know that iron will unite chemically with oxygen. Examine the substance left in the gas pipe, and recognize it as iron oxid. Hydrogen , Iron Oxygen + Ir0n = Oxygen + =**<*" c. Hydrogen from Sodium and Water. In a cage made of fine wire gauze and fitted with a handle, place a small piece of metallic sodium about the size of a pea. Have ready a pneumatic trough with an inverted beaker full of water on the shelf. Hold the cage containing the sodium under water. After a few bubbles have escaped, hold it under ELEMENTS AND COMPOUNDS 23 the beaker and collect the gas evolved. When action has ceased, recognize the gas as hydrogen. Into a small beaker full of water, throw a few small pieces of sodium, and test the liquid afterwards with red litmus paper. Sodium has so strong an attrac- tion for oxygen that it takes oxygen out of water, leaving the hydrogen. The resulting oxid unites with water and forms the hydroxid. If there is an insufficient supply of water, the sodium bursts into the characteristic yellow flame. To show this, place a small piece of sodium on a wet piece of filter paper, and note the result. , Hydrogen Sodium (i) Sodium + = 0xygen + Hydrogen Sodium Hydrogen dlu Oxygen Oxygen d. Hydrogen from Potassium and Water. Perform an ex- periment similar to c, using potassium instead of sodium. Hydrogen Potassium (i) Potassium + J + Hydrogen Oxygen Oxygen Potassium Potassium Hydrogen Oxygen Oxygen e. Hydrogen from Zinc and Sulfuric Acid. In a 500 cc. flask fitted with a two-holed stopper containing an exit and a funnel tube, place about 100 grms. of granular zinc. Add enough sulfuric acid (i part acid and 5 parts water) to cover the zinc. Catch the gas evolved over water, and recognize as hydrogen. Be sure to have no flame near the flask. 24 AN ELEMENTARY EXPERIMENTAL CHEMISTRY After no more zinc will dissolve, filter the contents of the flask ; then, after the solution has evaporated somewhat, allow it to crystallize. Note the properties of the crystals. Place a dry crystal in a glass tube closed at one end, and heat gently in the flame of a Bunsen burner. Notice the water that collects on the upper, cool part of the tube. Does the crystal remain intact ? This water, then, seems in some way to be necessary to the crystal, and is called " water of crystallization." Since sulfuric acid is composed of hydrogen, sulfur, and oxygen, and since in the experiment the zinc disappears and mu^ ^ m fSii'i'l ^Afri^Ki^iMii^ : -"-' I Hw^" ^^^^^^^^^^^PBBBBHBRl^ hydrogen is evolved, it is evident that the zinc takes the place of the hydrogen in the sulfuric acid. The crystals are therefore composed of zinc, sulfur, and oxygen, together with the water, which holds the combined elements in crystalline form. When this water is driven off, of course only zinc, sul- fur, and oxygen are left, in the form of a white powder. This is called zinc sulfate. Hydrogen Zinc Zinc + Sulfur = Hydrogen + Sulfur Oxygen Oxygen ELEMENTS AND COMPOUNDS / Hydrogen from Iron and Sulf uric Acid. Perform a simi- lar experiment, using iron nails instead of zinc. What is the color of the crystals obtained ? Of what are they composed ? What would you naturally name them ? Hydrogen Iron + Sulfur Oxygen Iron = Hydrogen + Sulfur Oxygen EXPERIMENT 16 Hydrogen and Oxygen a. Synthesis of Water by Burning Hydrogen in Air. Set up an apparatus as in Experiment 15 e. To dry the hydrogen, pass it through a catch bottle containing concentrated sulfu- ric acid, and have the delivery tube drawn to an opening of about i mm. diameter. After the gas has been escaping 26 AN ELEMENTARY EXPERIMENTAL CHEMISTRY for a few minutes, collect a test-tubeful over water. If, on ignition, no explosion is heard, it is safe to light the hydro- gen as it escapes from the apparatus. Never light hydrogen from a generator without taking this precaution. Hold a dry bell jar over the flame, and keep the jar cool with a cloth wrung out in cold water. Notice the water forming on the inside of the jar. Here we have further proof of the com- position of water. In Experiment 6 ^Lr powder in a mortar. Place this in a hard glass tube closed at one end. Fit the tube with a one-holed stopper con- taining an exit tube extending a centi- meter beneath the surface of a little lime water in a test tube. Heat the * The production of carbon dioxid by the processes of respiration and combustion depletes the supply of oxygen in a room. The oxygen must be renewed from the air outside, and the injurious respiration and combustion products must be removed by proper ventilation. This is especially true of churches, theaters, schoolrooms, and all in-door places where many persons are assembled at the same time. ELEMENTS AND COMPOUNDS 33 hard glass tube with a blast lamp. After the air has been driven out, notice the effect of the bubbles upon the lime water. Continue heating it until the milkiness of the lime water disappears. Then take away the lime-water tube (before taking away the lamp), and heat it over a Bunsen burner. Does the milkiness return ? We therefore know that one constituent of marble is carbon dioxid (see Experi- ment 20 d). Take another lump of marble about the size of a pea, and, holding with a pair of forceps over a porcelain plate, heat it with a blast lamp for about five minutes. Collect the particles that have fallen on the plate, and. together with what is left of the lump, place them in a very small quantity of water (only enough to mois- ten them) on a watch glass. Test this with moist litmus and turmeric paper. Now add more water (about i cc.), and filter into a test tube. Add a little carbonic acid, and notice the milkiness formed. On adding more carbonic acid, note that it disappears, and reappears on heating. Hence we see that another constituent of marble is calcium oxid (see Ex- periment 20 d), and that marble is composed of calcium, carbon, and oxvgen. Its chemical name is calcium carbonate. Calcium Carbon = Calcium Oxygen Carbon Oxygen Oxygen To make a complete analysis, we must take two more steps. First take a jar of carbon dioxid, and, after igniting 34 AN ELEMENTARY EXPERIMENTAL CHEMISTRY a ribbon of magnesium, plunge it into the jar. Note that the magnesium burns to the white oxid, and that carbon is deposited on the sides of the jar. Carbon , Magnesium + Magnesium == 4- Carbon Oxygen T Oxygen Secondly take about 5 grins, of calcium oxid, and grind it to a powder in a mortar. Mix with this 2.5 grms. of metallic magnesium powder, place the mixture in an iron crucible, and heat it over a Bunsen burner. While the contents of the crucible are still warm, hold it by means of a pair of long forceps under a beaker full of water and inverted on the shelf of a pneumatic trough. Recognize the gas collected as hydrogen. Since neither magnesium, magnesium oxid, nor calcium oxid decomposes water with the evolution of hydrogen (see Experiments 13 and 20), it must be that metallic calcium was formed in the crucible and, on being placed in the water, decomposed it, giving hydrogen. / N Calcium Magnesium 1 i ) _ + Magnesium = _ ' -f Calcium Oxygen Oxygen Calcium (2) Calcium + y r gen = Oxygen + Hydrogen xygen Hydrogen EXPERIMENT 22 Chlorin a. Properties. Apply to the instructor for a jar of chlorin. In working with this element, perform your experiments under a hood with a strong draught, since chlorin taken into the lungs is very dangerous. Note the color and odor (you can ELEMENTS AND COMPOUNDS 35 hardly avoid smelling it). Moisten a little blue and red litmus paper, also a colored flower, if possible, and allow them to hang in the gas for a few moments. Pour about 100 cc. of water into the jar ; seal, shake, and open it under water. Is chlorin soluble in water ? b. Hydrogen Chlorid (Hydrochloric Acid) by Burning Hy- drogen in Chlorin. Have ready a hydrogen generator (see Experiment 16 a\ When the air is driven out, ignite the jet and allow it to burn in a jar of chlorin until no green color remains. Blow your breath across the mouth of the jar, and notice that the gas fumes in the presence of mois- ture. Test with moist litmus paper. Taste the gas by allowing a little to enter the mouth. This is called hydrogen chlorid, or hydrochloric acid gas. c. Union of Hydrogen and Chlorin by means of Light. To be performed by the instructor. In a dark room with ruby light (a photographic dark-room), fill a 250 cc. flask one- half full of chlorin and one-half full of hydrogen, collected over warm water. Fit a rubber stopper to the flask, and 36 AN ELEMENTARY EXPERIMENTAL CHEMISTRY fasten it securely with wire. Wrap a towel or thick cloth around the flask so that no light can enter. Tie a long string to a corner of the towel, and carry all out of doors. At a safe distance, drag off the towel by means of the string so that the sunlight may strike the flask. What caused the action, and what is formed ? d. Sodium Chlorid. Apply to the instructor for a jar full of dry chlorin. Into this, drop a thin slice of metallic sodium about half a square centimeter in area, and seal. Allow it to stand until the next day. Then open the jar, and remove the white compound. Crush it, be sure there is no metallic sodium left, and then taste it. What is it ? Being formed by the synthesis of sodium and chlorin, it is called sodium chlorid. Note that, in cases where two elements unite, the name always ends in id. Sodium Sodium + Chlorin = _. , Chlorin EXPERIMENT 23 Hydrochloric Acid from Sodium Chlorid and Sulf uric Acid In a 250 cc. flask fitted with a two-holed stopper contain- ing funnel and exit tubes, place about 10 grins, of sodium chlorid. Have the exit tubes connected with a glass tube reaching to the bottom of a fruit jar covered with a piece of cardboard. Into the funnel tube, pour 15 cc. of concen- trated sulfuric acid diluted with 4 cc. of water. Pour the acid into the water. Collect two or three jars of the gas. Note that it is the same gas as that obtained by burning hy- drogen in chlorin. To one of the jars, add about 25 cc. of water, close and shake it. Open it under water. Is hydro- chloric acid soluble in water ? Into another jar, pour a little ELEMENTS AND COMPOUNDS 37 water, shake and test ttie liquid with blue litmus paper. Di- lute the contents of the flask with water after all the gas pos- sible has been allowed to escape. Evaporate and crystallize. Recognize the crystals as the same as those obtained in Ex- periment 1 8 a, i.e., sodium sulfate. Hydrogen Sodium Sodium _' Hydrogen Chlorin Chlorin Oxygen Oxygen EXPERIMENT 24 Analysis of Hydrochloric Acid a. By Electrolysis. Have ready an apparatus like that of Experiment 6 e, except that the electrodes are of gas carbon instead of platinum. Fill it with concentrated hydrochloric acid solution. On turning on the current, note the immediate evolution of hydrogen at the negative pole. Since you know that chlorin is soluble in water, you can easily account for the fact that chlorin does not immediately appear at the positive pole. Allow the current to pass through the solu- tion for some time ; then, when it has taken up all the chlorin it can hold, hydrogen and chlorin will appear in equal vol- umes at the negative and positive poles respectively. Hydrogen CMor . n Chlorin b. By Means of Sodium. Fill a dry quick-sealing fruit jar with hydrochloric acid gas dried by being passed through concentrated sulfuric acid. Be sure to have the rubber washer well greased with vaseline, so that the jar will be air tight when closed. Prepare about 25 grms. of sodium amal- gam (10 parts mercury to i part sodium), and powder it in a 38 AN ELEMENTARY EXPERIMENTAL CHEMISTRY mortar. Drop the powdered amalgam into the jar, seal quickly, shake vigorously, and note the action. Open it under water in a glass or porcelain dish, and, allowing the mercury to fall out, note the volume of the remaining gas. Remove the jar from the water, and recognize the gas as hydrogen. If the experiment is successful, you will have one-half a jar of hydrogen, showing that hydrochloric acid is composed of equal volumes of hydrogen and chlorin. EXPERIMENT 25 Hydrochloric Acid with Metals, Hydroxids, and Carbonates a. Try the effect of hydrochloric acid on the following metal.5 that you have studied, iron, magnesium, zinc, and copper, and verify the following statements : Hydrogen Iron (1) Iron+_;\ - + Hydrogen Chlorin Chlorin Hydrogen Magnesium (2) Magnesium + = + Hydrogen Hydrogen Zinc Chlorin = Chlorin Name the salts formed. b. Neutralize a small amount of hydrochloric acid, first with sodium hydroxid, second with potassium hydroxid, and third with sodium carbonate, and verify the following statements : , ^ Hydrogen _ Sodium Hydrogen W Oxygen + Chlorin ~ chlorin + Oxygen Hydrogen ELEMENTS AND COMPOUNDS 39 Potassium Hydrogen Potassium Hydrogen (2) Oxygen + Chlorin = chlorin + Hydrogen . Hydrogen _ Sodium Carbon Hydrogen + Chlorin ~ Chlorin + Oxygen Oxygen Oxygen Name the salts formed. c. Potassium Chlorid. Compare the salt formed in b i by neutralizing potassium hydroxid and hydrochloric acid, with the residue obtained in the hard glass tube in Experi- ment 6 c. We now see that potassium chlorate is composed of potassium, chlorin, and oxygen, and that the action in Experiment 6 c, can be expressed by the following : Potassium Potassium Chlonn = Chlorin + 0x ysen Oxygen EXPERIMENT 26 Carbon Dioxid from Marble and Hydrochloric Acid The best way to obtain a stream of carbon dioxid for use in the laboratory is by the action of hydrochloric acid upon marble (calcium carbonate). In a 500 cc. flask fitted- with funnel and exit tubes, place a number of lumps of marble, and then add dilute hydrochloric acid (i part acid to i part water). Carbonic acid is first formed, which, being very un- stable (see Experiment 8 c], breaks up into carbon dioxid and water. <; al c ) grms. of chlorin in the silver chlorid obtained. But this chlorin came from the a grms. of sodium chlorid taken ; therefore the percentage of chlorin in sodium chlorid is ^73_I -I, Your result should be about 60.6%. Atomic Theory. As soon as the preceding facts, i.e., the indestructibility of matter, the law of definite proportions, and the law of multiple proportions, had been established, men of science immediately began to look for some explana- tion of them. It remained for Dalton in the early part of this century to propose an hypothesis that is accepted to the present day. We are compelled to believe in its truth, for we can think of no other hypothesis that will account for the facts. Dalton 's hypothesis is this. All matter is composed of minute particles, which cannot be divided by any chemi- cal means. Each of these particles of any simple substance (i.e. element) is like every other one of that sub- LAWS AND THEORIES OF CHEMISTRY Q/ stance, both in properties and weight. These particles are called atoms. Therefore, an atom is the smallest particle of matter that can enter into chemical combination. These like or unlike atoms may unite with each other, and form other particles which are called molecules. Therefore, a molecule is the smallest particle of matter that can exist and still have all the properties of the substance, and is made up of atoms chemically united* This, briefly stated, is Dalton's Atomic Theory. It must be borne in mind that it is simply an hypothesis or guess formulated to explain known facts. If the hypothesis is the actual truth, then it follows that, in the case of copper and oxygen, when their atoms are chemically united to form copper oxid, the proportions will always be the same. In the case of two compounds of the same ele- ments, as for instance the two oxids of copper, it follows that the amount by weight of oxygen in the first must bear a simple ratio to the amount by weight of oxygen in the second. For evidently twice as many atoms of oxygen would unite with a definite number of atoms of copper to form the second oxid as would unite with the same number of atoms of copper to form the first oxid. In all cases, the ratio must be a simple one, since, if matter is made up of atoms, chemical combination must take place between the whole atoms. Atmospheric Pressure. Barometer. In order to understand the full meaning of the atomic theory, it will be necessary for the student to perform a number of experiments with gases. To do this, he must understand the meaning of atmospheric pressure, and the effects of changes of pressure and of temperature on the volumes of gases. First let us take up the subject of atmospheric pressure. * Some molecules are made up of single atoms. 98 AN ELEMENTARY EXPERIMENTAL CHEMISTRY EXPERIMENT 5 Barometer Take a glass tube of uniform bore (5 mm.) about a meter long, and close one end. Fill it with mercury, and, by means of a wire, remove what air bub- bles cling to the sides of the tube. Invert the tube in a trough of mercury, and measure the height of the column. Such an in- strument as this is called a ba- rometer. Now since the mercury remains in the tube, some force must be exerted to hold it there, and the only thing that can do so is the air pressing down upon the surface of the mercury in the trough. Therefore the weight of the mercury in the tube balances a column of air of equal cross section extending to the top of the atmosphere. If you should perform this experi- ment on different days, you would observe that the column stands at various heights. Hence we see that the pressure or weight of the atmosphere resting upon a certain part of the earth's surface varies at LAWS AND THEORIES OF CHEMISTRY 99 different times. The standard pressure is the pressure that will hold up a column of mercury 760 mm. high, when the thermometer stands at o C. This is about 15 Ibs. to the square inch or 1033.6 grms. to the square centimeter. Boyle's Law. Now we are ready to investigate the effect of pressure upon the volume of a gas. We all know by every day experience that gases can be reduced to a smaller volume by increasing the pressure upon them, and that they expand when the pressure is removed ; but most of us are perhaps ignorant, unless we have studied physics, of the exact effect of the pressure upon the volume. The follow- ing experiment will show that the volume of a gas varies inversely as the pressure exerted upon it ; i. e., if you double the pressure, you divide the volume by two, or if you treble the pressure, you divide the volume by three. Expressing the law algebraically, we have V : V : : P' : P in which V is the volume at the pressure P, and V is the volume at the pressure P'. This is called Boyle V Law. EXPERIMENT 6 Law of Boyle Have ready a clean dry glass tube of about 6 mm. bore, closed at one end and bent so as to form a narrow letter J, the hook being about 30 cm. long. Let the long arm be about 100 cm. in length. Into this, pour a little clean mer- cury, so that it stands about the same height in both arms. It will do no harm if it stands a little higher in the long arm. Now fasten the tube to a perpendicular support, and meas- ure from the base of the support to the top of the bore in IOO AN ELEMENTARY EXPERIMENTAL CHEMISTRY the short arm, reading the scale to tenths of millimeters. The edge of a visiting card held across the tube will aid the eye in taking the reading. Then measure the height of the mercury in the long and short arms. Do not touch the column of con- fined air with the hand ; for, in that case, the warmth of the hand will cause the air to expand and thus give incorrect readings on the scale. Now take the reading of the barometer. Four into the tube enough mer- cury to make the column in the long arm stand about 10 cm. higher. .To remove any air bubbles that may be held im- prisoned by the mercury, push a long iron wire down into the mercury ; then withdraw it, at the same time turning it around in the tube. Take the meas- urements as before. Do this as many times as the length of the tube permits. LAWS AND THEORIES ' OF"CHEM/STk% ' : /TQI Arrange your numbers in a table thus : ^S i ^ ^ S 1 3 1 ai ^ *E ; i~ * |c| < ri- "Sj^ . jl yit ~~ tc- ^ ^= c c^ 1 ^+ i== ll! ! II rt 1 til" = 52 I 1 Hll |! 1 Compare the volume at any pressure with the volume at any other pressure, and verify the proportion V : V : : P' : P. EXPERIMENT 7 Law of Charles Have ready a tube of i mm. bore about 100 cm. long, bent at right angles at a distance of about 50 cm. from the closed end, and containing a column of dry air confined by means of mercury, the mercury extending somewhat beyond the bend. Insert this through a cork stopper in a hole in the end of a shallow tin tray containing melting ice. Allow the air column to remain in the ice for a few minutes, keeping the ends of the mercury column at the same level. ' AN' ' EI5EMENTAHY EXPERIMENTAL CHEMISTRY Mark the end of the air column, and measure with a meter rod. Now insert the confined column of air in a steam jacket containing a thermometer ; then allow steam to pass through the apparatus. Keeping the two ends of the mer- cury level, again mark the end of the air column. Remove it from the jacket, and measure it. The difference between the lengths in ice and in steam will be the amount by which the air column has expanded in being heated 100. To find the amount one centimeter expands for one de- gree, divide this result by the original length x the number of degrees heated. This should give a number near .00366, which reduced to a fraction equals about ^\^. We there- fore see that, for every degree it is heated at a constant pres- sure, a volume of air expands -%\^ of what its volume was at o C. This is known as the Law of Charles. LAWS AND THEORIES OF CHEMISTRY IO3 Theoretically then, the volume of a gas at 273 below zero would become nothing. Of course this is impossible, but the point 273 is taken as the absolute zero of temper- ature. Then a in the ordinary scale w r ould be a -f- 273 in the absolute scale. Deduction of the Formula. Let A be the volume of a gas at o, then A H --- will be the volume at i 273 A + will be the volume at 2 273 A + 3 will be the volume at 3 273 A + will be the volume at t 273 and A + - - will be the volume at t' 273 Representing the volume at t by V, and the volume at t' by V, we have V = A + , andV' = A + 273 ' 273 Dividing and canceling the A's, we have V v~ 2 73 +r Combining this with the formula for pressure, V : V' : : P' : P, we obtain VP V'P' 273 -ft 7 2 73 +t' Boyle's and Charles s Laws are General. In the experi- ments we have used air, but it will be found by trial that other gases act in the same way as air. Thus we see that differ- ent kinds of matter in the gaseous state have a property in common. IO4 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Examples, i. The volume of a gas at 765 mm. pressure and 20 C. is 450 cc. ; what will its volume be at o and 7 60 mm. ? Ans. 422. -|- 2. At what temperature, pressure remaining the same, will a gas be double its volume at o ? Ans. 273. 3. Reduce the following volumes with the annexed tem- peratures and pressures to volumes at second temperatures and pressures : V t P V t' P' 500 100 800 mm. ? 20. 650 mm. 1000 70 900 mm. ? o. 200 mm. EXPERIMENT 8 Weight of a Liter of Air Fit a strong glass bottle of about one-liter capacity, with a one-holed rubber stopper containing a glass tube. Wire on to the tube a short piece of antimony-rubber tubing fitted with a pinch cock.* Make all connections air-tight with vase- line. Weigh the bottle and connections carefully to one centigram. Attach an air pump to the rubber tube, pump out as much air as possible, close the tube with the pinch cock, and weigh the bottle carefully again. Open the tube under water of about the same temperature as the room. After the water has ceased rushing in, hold the bottle so that the water stands at the same level both inside and out- side, close the pinch cock, remove the bottle, and weigh it again. Take the readings of thermometer and barometer. Arrange your calculations thus : Wt. of bottle full of air = a Wt. of bottle, air pumped out = b Wt. of air pumped out = a b Wt. of bottle with water = d Wt. of water = vol. air pumped out = d a = e. LAWS AND THEORIES OF CHEMISTRY 105 Therefore e cc. of air at the temperature and pressure on the day of the experiment weigh a b grams. Calcu- late the weight of one liter. You should obtain with this apparatus somewhere near 1.2 or 1.3 grams. The correct weight of a liter of air at o and 760 mm. pressure is 1.293 grams, which number we shall hereafter use. EXPERIMENT 9 Weight of a Liter of Carbon Dioxid at o and 760 mm. Fit a dry one-liter flask with a two-holed rubber stopper. In the holes, insert glass tubes bent at right angles, one reaching to the bottom of the flask and the other just through the stopper. Fit with rubber tubes and pinch cocks. Make all joints air-tight with vaseline, and weigh carefully to one centigram. Allow carbon dioxid, dried by being passed through sulfuric acid, to pass through the flask until it is com- pletely rilled. Close the tubes and disconnect them. Open IO6 AN ELEMENTARY EXPERIMENTAL CHEMISTRY one pinch cock for a moment to relieve the pressure, close it again, and weigh the flask. Take the temperature near the flask and the reading of the barometer. The increase in weight will be the amount by which the weight of the flask full of carbon dioxid exceeds the weight of 'the flask full of air. Mark the capacity of the flask, fill it with water to the mark, and weigh. Since i cc. of water, weighs one gram, the number of grams of water in the flask equals the number of cubic centimeters in the volume of the flask. Calculate what this volume of air at the temperature and pressure at the time of the experiment would be at o and 760 mm. Knowing that one liter of air at o and 760 mm. weighs 1.293 grams, find the weight of the air. Add to this the in- crease in weight due to the carbon dioxid, and obtain the weight of the same number of cc. of the gas. From this result, obtain the weight of 1000 cc., i. e., i liter of carbon dioxid. You should obtain a number somewhere near 1.97. EXPERIMENT 10 Weight of a Liter of Hydrogen Perform a similar experiment with hydrogen gas. Use a 100 cc. flask, and weigh it on a delicate balance. Keep the flask inverted when filling, and be sure the hydrogen is dry. To insure this, pass the hydrogen not only through sulfuric acid but also through a U tube containing pieces of granu- lar calcium chlorid. Make the hydrogen with zinc and sul- furic acid (one part acid to five of water). In this case, the weight of the flask full of hydrogen is less than that of the flask full of air ; hence you must subtract the decrease from the calculated weight of the flask full of air at o and 760 mm. You should obtain a result somewhere near .09 grams for i liter. (See illustration on page 107.) LAWS AND THEORIES OF CHEMISTRY IO/ Density. By the density of an element or compound is meant its weight in gaseous form compared with the weight of an equal volume of hydrogen, the temperature and pres- sure being the same. EXPERIMENT n Calculate the density of air and carbon dioxid compared with hydrogen, and show that they are 14.43 an( ^ 22 respec- tively. EXPERIMENT 12 Density of a Liquid in the form of Vapor Have ready a clean dry 100 cc. flask fitted with a one- holed rubber stopper containing a glass tube drawn out to a diameter of about i mm. Make it air-tight with a little vase- line, and weigh it carefully on a delicate balance. Take the reading of the thermometer and barometer. Place about 20 cc. of alcohol in the flask, and put the flask into boiling water up to its neck. When the alcohol has boiled away,* * Do this gradually so as not to blow out the stopper. IO8 AN ELEMENTARY EXPERIMENTAL CHEMISTRY ignite the vapor issuing from the tube, and the moment that the flame dies away, seal the tube with a blowpipe. Now weigh the flask again. Fill the flask with water by breaking the end of the tube under water. If the flask does not fill almost completely, the experiment must be repeated. The weight of the flask full of water minus the weight of the flask gives the volume. The increase in weight will be the weight by which the flask full of vapor of alcohol at 100 exceeds the weight of the flask full of air at the temper- ature of the room. Find what the volume of the air in the flask would become at o and 760 mm. Knowing the weight of a centimeter of air, find the weight of this volume of air. Add the increase in weight due to the alcohol vapor, and obtain the weight of the alcohol vapor at 100 and the pressure at the time of the experiment. Calculate the volume of the air in the flask at 100. Let us call this M. Then if V is the volume of the flask at the temperature of the experiment, the weight of the air left in the flask at 100 would be - times its M weight at the temperature of the experiment. We now have the weights of the flask full of air and alcohol at 100. Hence the density of alcohol referred to air is easily found. But air is 14.43 times as heavy as hydrogen; therefore, to find the density of alcohol compared with hydrogen, multiply by 14.43. Weights of the Atoms. The student will remember that, in the atomic theory, it is assumed that the atoms of each ele- LAWS AND THEORIES OF CHEMISTRY 1 09 ment are of the same weight. Under the present condition of methods of investigation, it is impossible to determine the absolute weight of these atoms with any satisfactory degree of accuracy. However, as soon as the atomic theory was accepted, men began to try to find out the weight of these atoms in terms of some atom taken as a standard. Hydro- gen being the lightest element known, its atom was taken as this standard. It required a number of years and the estab- lishment of auxiliary principles to fix even these numbers with any degree of certainty. As time went on, investigation brought out the principles of chemical, electrical, thermal, and isomorphic, equivalence of elements. The Law of Definite Properties by Volume, Avogadro 's Law, and the Periodicity of the Elements, were also established : and with the aid of these it has been possible to determine the relative weights of the atoms to a remarkable degree of accuracy. We will take up these subjects in turn. Chemical Equivalence. By experiment, it is found that on causing hydrogen and chlorin to unite, the ratio of the parts by weight are i : 35.4. That is, the chemical value of chlorin compared with hydrogen is 35.4. In like manner, 8 parts by weight of oxygen unite with i part by weight of hydrogen. We therefore say the chemical equivalence of oxygen is 8. This is also called its combining number. To find the chemical equivalent of an element in terms of another, it is necessary only to find the amount by weight of the first that unites with a fixed amount of the second, or the amount which replaces a fixed quantity in a compound. In some cases, that is, where there is more than one com- pound between the same element, it is difficult to decide which number to take. If we knew the exact number of atoms of each kind that unite to form a compound, or the exact number of atoms of an element replaced in a compound IIO AN ELEMENTARY EXPERIMENTAL CHEMISTRY by one atom of another element, it would be easy to deter- mine the atomic weight directly from the chemical equiva- lence. Thus, if we knew that one atom of hydrogen united with one atom of oxygen to form water, we could be sure that the atomic weight of oxygen was 8. Or, if we knew that two atoms of hydrogen united with one of oxygen, we could be sure of the number 16 as the atomic weight of oxygen. Not knowing this, however, we see that our data is insufficient, and we must look farther. EXPERIMENT 13 Chemical Equivalence of Zinc Have ready a clean 100 cc. flask containing a dilute solution of c.p. hydrochloric acid (10 cc. of water to 20 cc. of acid). Fit the flask with an air-tight one-holed rubber stopper fitted with a delivery tube. Weigh out as nearly as possible 1.25 grms. of c.p. zinc, recording the exact weight. Place in a pneumatic trough. a 500 cc. flask inverted and full of water. Have the delivery tube clamped so that its end is under the mouth of the flask. Remove the stopper from the 100 cc. flask, drop the zinc into it, close it quickly, and catch all the evolved hydrogen. When all the zinc is dissolved, hold the flask containing the hydrogen so that the water stands at the same level both inside and out- side. Do not touch with the hand that part of the flask con- taining the hydrogen. Take the temperature near the flask LAWS AND THEORIES OF CHEMISTRY II I and the reading of the barometer. Mark the point on the flask at which the gas stands, remove it, and fill it with water to this point. Then weigh the flask and water. From this, obtain the volume of the gas. To obtain the best results, it will be necessary to take into account the fact that, besides the hydrogen, there is vapor of water in the flask, which helps bear the pressure of the air upon the surface of the water in the trough. This varies for different temperatures. The following table is near enough for all practical purposes. DEGKKUS. X H X H x p Gfl 1 jj^l 1 Hi a a > j *jj i X MlLI.IMICTKKS OF MliKCDKY. IO 9-2 16 13.6 22 19.7 28 28.1 II 9 .8 17 14.4 =3 20.9 29 2 9 .8 . I2 10.5 18 154 24 22.2 3 31-5 13 II, 19 16.4 =5 23-5 3 1 33-4 14 II.9 20 17.4 26 25.0 3 2 35-4 15 12.7 21 18.4 27 26.5 33 37-4 Subtract the number of millimeters of mercury in the table at the temperature corresponding to the temperature observed near the flask, and you will have the true pressure to which the hydrogen is subjected. 112 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Calculate by means of the formula V P V P' 273 + t 273 + t' what the volume of the hydrogen would be at o and 760 mm. Knowing that i cc. of hydrogen, at o and 760 mm., weighs .00009 grins., calculate the weight of the hydrogen ; then obtain from this the number of grams of zinc that will replace one gram of hydrogen. You should obtain some- where near the number 32.5. Should time allow, it will be interesting for the student to determine the number for iron, using piano wire instead of zinc. The number in this case is 27.9. Electrical Equivalents. Michael Faraday found that, when a current of electricity was passed through a substance which it decomposed, the quantity of the substance decomposed varied with the strength of the current. If the same current passes through two substances, such as a solution of copper sulfate and a solution of silver nitrate, the amount of copper deposited is to the amount of silver deposited as the chemical equivalent of copper is to the chemical equivalent of silver. But it is found that the same difficulty exists as with finding the chemical equivalents ; namely, that, if there are two or more different compounds between the same elements, then as many different electrical equivalents are obtained. Thus copper gives 31.6 and 63.2, and mercury gives 99.9 and 199.8. So we have still the same question as to which num- ber shall be taken as the atomic weight. EXPERIMENT 14 Electrical Equivalents Have ready two Daniel cells. In a beaker, put a solution of copper sulfate, and in another, put a solution of nickel LAWS AND THEORIES OF CHEMISTRY 113 sulfate. Cut out two copper plates three centimeters wide by six long, making a hole in each, and weigh one care- fully. Cut out two similar nickel plates of the same size, and weigh one. Place the copper plates in the copper sulfate, and the nickel plates in the nickel sulfate. Connect the cop- per plate that was not weighed with the weighed nickel plate by means of a wire, and then put the two solutions in circuit with the two Daniel cells, so that the weighed plates are the cathodes. Allow the current to pass for about twenty minutes ; then disconnect, dry, and weigh the plates again. Arrange your calculations as follows : Wt. cathode copper plat3 : Wt. after Difference Wt. cathode nickel plate = Wt. after Difference Compare the weights of the metals deposited, and see whether they are in the same ratio as the chemical equiva- lents of the two metals, 63.2 and 58.6. Specific Heat. We have noticed in our every day expe- rience that it requires a different amount of heat to raise the temperature of various substances the same number of degrees. Since heat is a form of energy, it can be measured. To measure a quantity of heat, it will be necessary to use some definite quantity as a unit of measurement. The unit used is the calorie. A calorie is that quantity of heat that is used up in raising one gram of water one degree centigrade* 114 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Therefore to raise one gram of water 2, or two grams i requires two calories. The amount of heat, measured in calo- ries, necessary to raise one gram of a substance i C., is called the specific heat of that substance. Below is given the specific heat of a number of the elements. Iron .114 Copper .094 Silver .057 Gold .032 Nickel .108 Zinc .095 Tin .055 Lead .031 Apparatus used in Determining Specific Heat. To find the specific heat of elements experimentally, two pieces of apparatus are necessary ; first, a vessel, called a calorimeter, to hold the substance. An ordinary nickel-plated lemonade shaker will answer very well. Secondly, we must have some sort of apparatus in which the substance can be heated to a constant temperature. Such an apparatus (see illustra- tion, p. 115) can be obtained at small cost from dealers in laboratory supplies, or one can be set up as follows. Have ready a large beaker of about 1000 cc. capacity, one-quarter full of water. Then cut a thin board in the shape of a ladle, large enough to make the broad part cover the large beaker. Cut a hole in the center of the broad part large enough to hold another beaker of about 200 cc. capacity. EXPERIMENT 15 Specific Heat of Lead To determine the specific heat of lead with any degree of accuracy by means of the apparatus here used, we shall assume that the specific heat of the material of the calorim- eter is known. In the case of brass it is .094, in the case of glass .2. LAWS AND THEORIES OF CHEMISTRY I I $ Weigh out 500 grms. of fine shot, and place the shot in the small beaker. Set the small beaker in the large beaker, and heat the water in the large beaker to boiling. Cover the small beaker with a piece of cardboard. Meanwhile weigh the calorimeter, and add exactly 100 grms. of water cooled about 8 below the temperature of the room. Stir the shot occasionally and thoroughly with a thermometer, and note when the temperature becomes constant. Then pour the shot quickly into the calorimeter, stirring with another thermometer, and note the temperature of the mixture. (The illustration shows apparatus mentioned on p. 114.) Let us see what has taken place. The amount of heat lost by the shot (we neglect what little heat has been radiated off) must be equal to the amount of heat gained by the water and the calorimeter. Il6 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Arrange your calculations as follows : Weight of the shot = s = Weight of the calorimeter = c = Weight of the water = w = Temperature of the shot = t s = Temperature of the water = t w = Temperature of the mixture = tm = Specific heat of substance of cal. = h = .094 Amount of heat lost by shot = amount of heat gained by water -(- amount of heat gained by the calorimeter. Let x = the specific heat of lead. Then, since i gram of lead loses x calories in falling i, s grams of shot will lose s (t a / m ) x calories in falling 4 / m . Since one gram of water gains one calorie in being heated one degree, w grams of water will gain w (t m t w ) calories in being raised / w t w . Since one gram of the substance of the calorimeter gains h calories in being raised one degree, c grams will gain eh (t m 4,) calories in being raised t m t w . Therefore we have the original statement expressed algebraically. s (ts t m ) x=w (tm tw) + ch (t m tw) Substituting your observed values, you should obtain, by solving for x, a number somewhere near .03 1 as the specific heat of lead. EXPERIMENT 16 Specific Heat of Iron In the same way, find the specific heat of iron, using iron filings free from grease or oil. You should obtain some- where near the number .114. LAWS AND THEORIES OF CHEMISTRY 117 Thermal Equivalents. In 1819, two French chemists, Dulong and Petit, noticed that there existed a simple rela- tion between the chemical equivalents of the elements and their specific heats ; namely, that the product of the two numbers always approximates 6.4, or half that number. In most cases it is 6.4. The following table shows this remark- able relation. The chemical equivalents are doubled where it is necessary to obtain approximately the product 6.4.* ELEMENT. U H ll SPECIFIC HEAT. PRODUCT. ELEMENT. ATOMIC WEIGHT. SPECIFIC HEAT. PRODUCT. Iron 55.6 .114 6.4 Silver lOJ.II 057 6.1 Nickel 58.24 .108 6-3 Tin 118.15 055 6. 5 Copper 63.12 .C 94 6.0 Gold 195-74 .032 6-3 Zinc 64.91 .095 6.1 Lead 205.36 .031 6.4 The product of the atomic weight and the specific heat varies somewhat, of course ; but it is only reasonable to sup- pose that, if our numbers were absolutely correct, this product would always be the same. The quotient of the number in the last column divided by the specific heat is called the thermal equivalent. Explanation. Here we have a method of aiding us in determining what numbers we shall take for the atomic weights. The only way we can explain the relation noted in the last paragraph is by saying that all atoms have the * This product will probably be somewhat smaller than 6.4 as the atomic weights. are more and more accurately determined. I 1 8 AN ELEMENTARY EXPERIMENTAL CHEMISTRY same capacity for heat ; that is, it requires the same amount of heat to raise every atom, no matter of what kind, one degree in te?npe?'ature. For example, the atomic weight of lead is 205 microcriths (a microcrith being the weight of an atom of hydrogen), and that of iron is 56 microcriths. If we assume that it requires the same amount of heat to raise an atom of each one degree, it will take 205 as much to raise one microcrith of lead one degree, and ^ as much to raise one microcrith of iron one degree. But the atomic weights must bear to each other the same ratio as their mass weights, therefore the amount of heat necessary to raise one gram- of lead one degree must be to the amount of heat ne- cessary to raise one gram of iron one degree as ^J^ is to g 1 ^, or as 56 is to 205. The specific heats of lead and iron, .031 and .114 respectively, are to each other in this ratio. We are then justified in believing that our assumption that every atom requires the same amount of heat to raise it one degree in temperature is true. We have therefore considerable evidence to warrant us in believing that atomic weight x specific heat = a constant, i.e., 6.4 From this we have 6.4 atomic weight = specific heat or 6.4 specific heat = atomic weight From Experiment 13, we find that 32.5 grams of zinc re- place one gram of hydrogen. Suppose we take 32.5 for the atomic weight, and divide 6.4 by this number. We obtain .193 for the specific heat. Now let us double the number 32.5, and divide 6.4 by the result. We then obtain .096 for the specific heat. Since the specific heat of zinc is LAWS AND THEORIES OF CHEMISTRY I IQ found by experiment to be .0955, we are justified in believ- ing that the atomic weight is 65, just twice its chemical equivalence. Isomorphic or Crystallographic Equivalents. A short time after Dulong and Petit made their discovery, Mitscherlich, a German chemist, found that certain elements can replace others in a compound without changing the crystalline form. Such elements are said to be isomorphous. The replace- ment always takes place in definite quantities. These quantities are called the Crystallographic equivalents. The chemical equivalence of silver is 107. In a crystal, 107 parts by weight of silver can be made indirectly to re- place 65 parts of zinc; hence we say that 107 parts of silver is the Crystallographic equivalent of 65 parts of zinc. But the chemical equivalence of zinc is 32.5. The only explana- tion we can give of this is that one atom of zinc replaces two atoms of hydrogen; and hence the atomic weight of zinc is not the same as its chemical equivalence, but twice 32.5, or 65. This is evidently in harmony with what we learned in regard to the atomic weight of zinc from the Law of Dulong and Petit. In like manner, the atomic weights of other elements have been investigated, and it has been found that this principle is a great aid in determining the right number to use in many cases. It has not however been found to be universally true. EXPERIMENT 17 Law of Definite Proportions by Volume Refer to Experiment 6 N 1 oo' 1 & s, K. 28 * o 1 H g 3 & h 13 > > c N N Thallium Thorium Thulium C H Titanium Tungsten Uranium Vanadium Ytterbium Yttrium u c N Zirconium 5 s 8 I oo 1 00 8 1 1 CT^ >)' . 00 ? oo 1 .8 s, 3 PS CO CO c^ CO 60 g CO CO H H e 'a " 3 1 E 3 1 E 3 'i c o W 3 rontium J3 a illurium E 3 X* CO CO CO CO CO Cfl CO CO H h f-H i oo ? 1 8. 1 1 i 1 i fx. oo" % z }5 6 o Pi 0, ^ ^ JS JO 11 13 c 1 Osmium 1 o Palladium Phospho- rus 3 C rt s Potassium Praseody- mium Rhodium Rubidium q 1 vO vO 2 s i ? 1 1 4 S 1 c? ffi c KH i-, V fe J J3 PH 3 I 3 c 00 ffi Hydrogen Indium fi | Iridium c Lantha- num 1 Lithium is rt C Mercury Molybde- num oo rt 1 q V? 1 00 '? 00 ? o 1 D u O U 3 W tu s 0) C - 3 Chlorin c 1 e 1 U Jl y 1 o U E 3 w Flu or in Gadolin- ium Gallium Germa- nium Glucinum T3 O P 1 S *? CO I ro R ? f c> 5" o ^ JO CO < 03 2 22 03 s U rt U CJ CJ 1 Aluminum Antimony '2 . .5 1 1 Bismuth C o 02 C a 1 Cadmium I Caesium Calcium 1 Carbon 1 E 1 LAWS AND THEORIES OF CHEMISTRY The table on page 128 gives a list of the elements with their symbols and atomic weights.* Determination of Molecular Formulae. How do we know that sulfuric acid is H 2 SO 4 , that copper sulfate is CuSO 4 , that alcohol is C 2 H 6 O, etc. ? First we must determine by analysis the percentage com- position of the compound. Let us take as an example the last one above mentioned, alcohol. When analyzed, alcohol is found to contain Carbon 52.17% Hydrogen 13.05 % Oxygen 34 .78 % IOO.OO % We find that the ratios 52.17:13.05:34.78 can be ex- pressed in any of the following ways : 4 I 2.66 8 2 : 5-32 12 3 7.98 16 4 10.64 20 5 13-30 2 4 6 15.96 and the last of which is practically 24:6:16. Therefore 24 grams of carbon, 6 grams of hydrogen, and 16 grams of oxygen unite, forming 46 grams of alcohol. Or, expressing the quantity of each element in microcriths, 24 microcriths of carbon, 6 microcriths of hydrogen, and 16 microcriths of oxygen unite, forming 46 microcriths of alcohol. But, since an atom of carbon weighs 12 microcriths, an atom of hydro- gen i microcrith, and an atom of oxygen 16 microcriths, we see at once that the simplest composition of a molecule of alcohol would be two atoms of carbon, six atoms of hydrogen, and one atom of oxygen, i.e., C 2 H 6 O. * From report of American Chemical Society, 1899. I3O AN ELEMENTARY EXPERIMENTAL CHEMISTRY We might just as well use the ratios 48 : 12 : 32 Or 64 : 18 : 48 We should then have obtained either the formula C 4 H 12 O 2 or C 6 H 18 O 3 . The only way we have of deciding which is cor- rect is by comparing the molecular weights which these for- mulae give, with the density of alcohol in the state of vapor. C 2 H 6 gives a molecular weight of 24 + 6 + 16 = 46 C 4 Hi2 2 gives a molecular weight of 48 + 12 + 32 = 92 Ce HIS Oa gives a molecular weight of 72 + 18 + 48 = 138 By weighing alcohol in the state of vapor and comparing it with hydrogen, we found in Experiment 12, Part II. that its destiny is 23. Therefore its molecular weight must be 2 x 23, or 46. Hence we take the formula C 2 H 6 O as the correct one. In the case of compounds whose density can- not be found, we take the simplest formula. The simplest way numerically to find the formula when the percentage composition is given is as follows : Carbon 52.17 Hydrogen 13.05 Oxygen 34.78 Rule. Divide the percentage of each element by its atomic weight. Find what the resulting ratios become, if the smallest number be taken as unity. If there is only one atom of that element, the ratios will then be expressed by whole numbers. If the resulting ratios are not whole numbers, try successively what the ratios would become, if the smallest number were taken as 2, j, etc. Examples. Deduce the formulae for the following sub- stances : LAWS AND THEORIES OF CHEMISTRY I. 2_ S- 6. f r I r I r 1 r j ^ L 1 Hydrogen Oxygen Density, 9 Hydrogen Oxygen II. 12 88.88 IOO.OO 5-88 94.12 IOO.OO Manganese 72.05 Oxygen 27.95 Silicon Oxygen Iron Oxygen Carbon Hydrogen IOO.OO 46.67 53-33 IOO.OO 70.01 29-99 IOO.OO 92.3 7-7 100.00 Density, 13. f Copper 1 Sulfur 7*1 Oxygen f Hydrogen ~ J Phosphorus j Oxygen f Potassium J Nitrogen j Oxygen f Nitrogen | Hydrogen 10. -| Carbon Oxygen 39-62 20.13 40.25 IOO.OO 3.06 31.64 65-30 IOO.OO 45-95 16.45 37-6o IOO.OO 29.17 8.33 12.50 50.00 100.00 Valence. Let us suppose that we know the molecular for- mulae for all the known compounds. Then, if we arrange a few of the simplest thus : I. II. III. IV. HF H 2 H 3 N H 4 C H Cl Ho S H 3 P Hi Si H Br H 2 Se H 3 As HI H 2 Te H 3 Sb we shall see that an atom of each element possesses the power of uniting with a certain number of hydrogen atoms. Again, let us make a table thus : Acids Salts HC1 AgCl CuCl 2 A1C1 8 HN0 3 AgN0 3 Cu (N 3 ) 2 Al (N Os) 3 H 2 S 4 Ag 2 S 4 CuS0 4 A1 2 (S 4 )3 H 3 P0 4 Ag 3 P 4 Cu 3 (P 4 ) 2 A1P0 4 132 AN ELEMENTARY EXPERIMENTAL CHEMISTRY We here see that an atom of each element possesses the power of replacing a certain number of hydrogen atoms in a compound. This property of an atom is called its valence. We might define valence as the quantity of combining power an atom has compared with that of an atom of hydrogen. Atoms that combine with or replace an atom of hydrogen are called monovalent, or monads. Those which combine with or replace two atoms of hydrogen are called divalent, or dyads. There are also trivalent, tetravalent, etc., atoms. For some of the elements, the number representing the valence varies. For instance, nitrogen has a variable valence, being some- times triad and sometimes pentad. The following table shows the valence of the more common elements. Na i Ni 2 2 P 3,5 Pt 4 Pb 2 S 2,4,6 Si 4 Sn 2,4 Sr 2 Sb 3,5 Zn 2 Graphic Formulae. Although we do not know the nature of chemical affinity, still, as we have seen, we can express it quantitatively. We can represent monad, dyad, triad, etc., valence by lines ; a monad element being written H - , a I dyad, - O - , a triad, B - , and so on. Let us start with the compound H 2 O, and write it H - O- H, Al 3 Cr 3 As 3, 5 CU 2 Agi F i B 3 Fe 2,4 Ba2 H i Bi 3, 5 Hg2 Br i (3, 5, 7) I i(5) C 4(2) K i Ca 2 Li i Cl i (3, 5, 7) Mg2 Cd2 Mn 2 Co 2 N 3,5 LAWS AND THEORIES OF CHEMISTRY 133 TT or _ O, understanding by this that one atom of oxygen rl ^ with a valence of two is united with two atoms of hydrogen, each having a valence of one. Suppose now that we replace the monad hydrogen atoms by two monad sodium atoms. We shall then have for sodium oxid the formula Na Na or JJ* > In like manner for potassium oxid, we should have K O K. If we replace the two monad hydrogen atoms by a single dyad atom, such as calcium, we should have Ca = O for calcium oxid. Using the triad aluminum atom, we should have to replace the six monad hydrogen atoms of three molecules of water by two aluminum atoms. H - o H- ..^0 H-" In like manner other graphic formulae of other oxids may be written. The corresponding hydroxids, of course, will be n /O H Na H, K H, Ca<5~JJ, Al H the metals replacing only one hydrogen atom from each water molecule. In the case of an acid, we write H Cl for hydro- chloric, H Br for hydrobromic, etc. We then have : SODIUM SALTS CALCIUM SALTS ALUMINUM SALTS / Cl Na Cl Ca aluminum sulfate Al 0\ Q ^0 ^0/ S =-0 Similarly phosphoric acid, H 8 PO 4 , contains three hydroxyl groups, thus : H 0-, H P H 0^ from which we have the salts, Na 0^ Na P 0, sodium phosphate Na O-' Ca < ^ calcium phosphate P Al P 0, aluminum phosphate Acids that contain one hydroxyl group are called mono- basic, those containing two such groups are called dibasic, those containing three are called tribasic, and so on. When all the hydrogen of the hydroxyl groups is replaced by a metal, the resulting salt is said to be normal. Evidently it is possible that only the hydrogen of one hydroxyl group in a dibasic or higher acid may be replaced, thus giving rise to acid salts. Hence two different salts containing the same metal may be obtained from a dibasic acid, three from a tri- basic acid, and so on. We may illustrate this as follows : Q f Na -0 ^0 S ^ , acid sodium sulfate H O^ ^ % u ^ Na - ^ "^ S ^ , normal sodium sulfate 136 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Na ^ H P 0, dihydrogen sodium phosphate H -0 ^ Na ^ Na P 0, hydrogen disodium phosphate H ^ Na ^ Na P 0, normal sodium phosphate Na ^ At present little is known concerning the molecular struc- ture of basic salts. Some think that to form these salts, additional molecules of the base replace part of the oxygen of the acid ; while others think that they are combinations of the normal salt molecules with extra base molecules. ^H The molecular formula for ammonia is NH 3 , or N H ^H We have formed ammonium chlorid, NH 4 C1, synthetically by the union of ammonia and hydrochloric acid, giving of course ^H Cl-N < and showing the pentad valence of nitrogen. The group NH 4 is evidently a monad group, and as such appears in H ^ TJ ^ N H, ammonium hydroxid 11 ^ H ^ H _ Nj , ammonium nitrate Q, ammonium sulfate LAWS AND THEORIES OF CHEMISTRY 137 This group is called the ammonium group, and acts very much like a metal. Thus the student will see the possibility of representing through the eye relations that would otherwise be incom- pletely realized. It must not, however, be thought that these structural formulae are intended to represent the actual position of the atoms of a molecule with respect to each other. The intention is merely to emphasize certain relations that experiment and reason have shown must exist between the atoms in the molecule. Examples. Write the graphic formulae for silver nitrate, aluminum hydroxid, calcium carbonate, and ammonium phos- phate. Positive and Negative Elements. Chemical compounds can be decomposed by electricity. In all such decompositions, part of the atoms appear at the negative pole, and part at the positive pole. Those which appear at the negative pole are called electro-positive, or simply positive, elements ; and those which appear at the positive pole are called electro- negative, or simply negative. One element, however, need not always appear at the same pole. When liberated from its union with one element, it may appear at the negative pole ; while, when liberated from another, it may appear at the positive pole. The non-metals are negative with re- spect to the metals, but it is readily seen that, compared with one another, they may be one or the other according to what elements they are. The strength of chemical com- bination depends upon this, the rule being that the more electrically remote the elements are, the stronger their union. Naming of Compounds. Binary compounds are those composed of two elements. All others are called ternary. The names of binary compounds, the names of ternary com- 138 AN ELEMENTARY EXPERIMENTAL CHEMISTRY pounds that do not contain oxygen, and the names of com- pounds of positive elements with hydrogen and oxygen, end in id. Examples respectively, NaCl, sodium chlorid ; KCN, potassium cyanid ; Zn(OH) 2 , zinc hydroxid. In all cases the positive element is named first ; then the negative ele- ment or radical, with the suffix. When more than one compound is formed between the same elements, the name of the positive element ends in ous or if, according as there is less or more of the negative element or radical. Examples Hgl, mercur0/w iodid; HgI 2 , mercur/r iodid. An acid takes its name from the characteristic element in it ; and when there is more than one acid from the same element, the name ends in ous or ic according as there is less or more oxygen. Examples, H 2 SO 3 , sulfuiw/j acid ; H 2 SO 4 , sulfur/V acid. If there are more than two, the prefixes hypo, meta, per, and pyro, are used also. Examples H 2 SO 2 , hypo- sulfurous acid ; HC1O 4 , perchloric acid. The names of acids that contain no oxygen have the prefix hydro. Example, HC1, hydrochloric acid. Salts take their names from the name of the positive ele- ment and that of the acid. The positive element is named first, and then the acid ending in ite or ate. The names of salts from "ous" acids end in ite\ those from "if" acids end in ate. Examples, Na 2 SO 3 sodium sulf//!?/ N^SO^ sodium sultafe. When there is more than one salt from the same element and acid, the name of the positive element ends in ous or if, according as there is less or more of the acid radi- cal. Examples, HgNO 3 mercur0j nitrate ; Hg(NO 3 ) 2 , mer- curif nitrate. Writing of Reactions. In order to be able to write reac- tions by means of chemical symbols, it will be necessary for the student to learn the valence of each of the ele- ments. LAWS AND THEORIES OF CHEMISTRY 139 Suppose we wish to express qualitatively the fact that the action of silver nitrate upon sodium chlorid gives silver chlorid and sodium nitrate, we may do so with a formula like that used in Part I. Silver Sodium Nitrogen + " = Nitrogen + Oxygen Oxygen Chlorin In order to express it quantitatively, we must make use of the molecular formulae of the molecules that act upon each other, keeping in mind the valence of the atoms of which these molecules are composed. Thus, Ag N 3 + Na Cl = Na N 3 + Ag Cl Sometimes it will be necessary to use more than one mole- cule of each substance, as in the case of the action of steam upon red-hot iron. 3 Fe + 4 H 2 = Fe 3 04 + 4 H 2 These formulae not only show what action has taken place between the molecules, but also how much matter is involved. Thus they indicate the fact that there is just as much matter after the action as there was before. For in- stance, the formula, Ag N 3 + Na Cl = Na N Cs + Ag Cl stands for the following sentence. One molecule of silver nitrate, molecular wt. 168.68, acts with one molecule of sodium chlorid, molecular wt. 58.06, producing'one molecule of sodium nitrate, molecular wt. 84.45, an d one molecule of silver chlorid, molecular wt. 142.29. The student will be aided considerably in writing reac- tions by the following illustration. Suppose we wish to write the reaction for calcium nitrate and sodium phos- phate. Write the formulae for the two compounds thus : ^N Oz Na ^ Ca + Na P Na ^ I4O AN ELEMENTARY EXPERIMENTAL CHEMISTRY Knowing that calcium is dyad in valence, and that phos- phoric acid is a tri-basic acid, we see that we must take enough molecules of calcium nitrate to make the total va- lence of the calcium atoms equal to 6, and enough sodium phosphate molecules to make the total valence of the sodium atoms equal to 6, i.e., in both cases the least common multiple of 2 and 3, the valence of calcium and the basicity of phosphoric acid respectively. We then have Na Ca Ca Ca N0 3 N0 3 N0 3 N0 3 N0 3 N0 3 NaP0 4 Na Na NaP0 4 Na The three Ca atoms will now replace the six sodium atoms, unite with the two PO 4 groups, and give one mole- cule of calcium phosphate, indicated thus : N 3 Na Ca N 3 Na P0 4 Na N O 3 Ca N 3 Na N0 8 Na P0 4 Ca N 3 Na LAWS AND THEORIES OF CHEMISTRY The remaining Na's and NO 3 's will unite, forming six molecules of sodium nitrate, thus : We then have the complete equation 3 Ca (N 3 ) 2 + 2 Na 3 P 4 = Ca 3 (P 4 ) 2 + 6 Na N O a EXPERIMENT 20 Write the reactions for all the experiments in Part I., using chemical symbols. Stoichiometry. Since chemical action takes place between molecules of elements or compounds, in order to find the mass weights of elements or compounds formed, knowing the mass weights of elements or compounds used, we have only to make use of the following : First write the reaction representing the chemical change. Then make the proportion, molecular weight of the given substance is to the molecular weight of the required substance as the mass weight of the given substance is to the mass weight of the required substance. 142 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Let us take an example. How many grams of hydrogen will be evolved, when 10 grams of zinc are dissolved in sulfuric acid ? H 2 S 4 + Zn = Zn S 4 + H 2 98 + 65 = 161 + 2 In this case we have 65 : 2 : : 10 : X. .'. X = .30+ grms. Suppose we wish to know the number of grams of zinc sulfate formed. We then have 65 : 161 : : 10 : X. .'. X = 24.7+ grms. Suppose we wish to know the number of grams of sulfuric acid required to dissolve 10 grams of zinc. We then have 6s:g8::io:x, .'. x = 15.0+ grms. In case we wish to know what the volume of the hydrogen evolved would be at o and 760 mm., all we have to do is to divide the weight of the hydrogen by .0896, the weight of a liter of hydrogen at o and 760 mm. (See Exp. 10, Part II.) If we wish to know what this volume would be at any required temperature and pressure, we can easily find it by using the formula VP 273+ (See Exp. 6 and 7, Part II.) Examples, i. How many grams of potassium chlorate must be used to obtain 100 grams of oxygen ? Ans. 2 64+ grms. 2. In order to fill a balloon, 150 kilograms of hydrogen are necessary. How much zinc and sulfuric acid will be required to produce the gas ? Zinc, 4871; kilos. Ans. .,.., , ., Sulfuric acid, 7350 kilos. LAWS AND THEORIES OF CHEMISTRY 143 3. What would the volume of the gas be at o and 760 mm. pressure ? What would it be on a day when the temperature was 20 and the pressure 755 mm. ? 1666.6+ liters. Ans. 1800.5+ liters. 4. How many grams of silver nitrate would be required to make 20 grams of silver chlorid ? Ans. 23.7 + grms. 5. How many grams of iron sulfid must be used to pro- duce, at o and 760 mm., 10 liters of hydrogen sulfid, i liter of hydrogen sulfid weighing 1.52 grams? Ans. 39.3+ grms. 6. Calculate the amount of manganese dioxid that must be used to produce, at o and 760 mm., 10 liters of chlorin from hydrochloric acid, i liter of chlorin weighing 3.17 grms. Ans. 38.9 grms. 7. What weight of copper would be used in making 20 grams of copper nitrate by dissolving the copper in nitric acid ? Ans. 6.9 grms. The reactions also indicate the relative volumes involved in the case of gaseous factors and products. By the equation H 2 + C1 2 = 2 H Cl we indicate that i volume of hydrogen combines with i vol- ume of chlorin, forming 2 volumes of hydrochloric acid. In the same way aH 2 + 2 = 2 H 2 indicates that 2 volumes of hydrogen unite with i volume of oxygen, forming 2 volumes of water vapor. Examples, i. 18 cc. of hydrogen are mixed with 10 cc. of chlorin and exploded. What gases are formed, and what are their volumes ? 2. If 2 volumes of nitric oxid and 5 volumes of hydrogen are united, what volume of ammonia is produced ? 3. To a certain volume of hydrogen sulfid gas, was added 144 AN ELEMENTARY EXPERIMENTAL CHEMISTRY 125 cc. of chlorin gas, which was entirely consumed. What volume of hydrochloric acid gas was evolved ? 4. If 150 liters of marsh gas be exploded with 300 liters of oxygen, what volume of carbon dioxid would result ? Calculation of Percentage Composition, having given the Molecular Formula. It is often required that the student should be able to calculate the percentage composition of a compound when the molecular formula is known. Suppose we have given the formula for alcohol, which is C 2 H 6 O. Carbon ax 12 = 24 Hydrogen 6 x i = 6 Oxygen i x 16 = 16 Molecular weight = 46 In alcohol, therefore, there are 12 parts by weight of carbon, 6 parts of hydrogen, and 16 of oxygen. This reduced to the decimal system becomes Carbon ^ = 52-17% Hydrogen = 13-05 % Oxygen = 34.78 % 100.00 % In case there is water of crystallization in the compound, the water molecules combined with each molecule of the compound must also be taken into account. Examples, i. Calculate the percentage composition of AgCl; NaN0 3 ; Ca 3 (PO 4 ) 2 ; CHC1 3 ; K 4 Fe(CN) 6 ; (NH 4 ) 2 SO 4 . 2. Calculate the percentage composition of MgSO 4 , 7H 2 O; HNaNH 4 P0 4 ,4 H 2 O. Thermochemistry. It is necessary even for the student of elementary chemistry to understand at least the fundamental LAWS AND THEORIES OF CHEMISTRY 145 relations between heat and chemical change. All chemical changes are accompanied either by the using up or the giving out of heat.* For instance, when 2 grams of hydrogen unite with 18 grams of oxygen, 68924 heat units are liberated. Such actions are called exothermic. There are also chemi- cal changes which absorb heat. For instance, to unite i gram of hydrogen with 127 grams of iodin to form hydrio- dic acid, requires 6000 heat units. Such actions are called endothermic. * It is well to note here the relation between chemical energy and other forms of energy. Energy is the power that matter has of doing mechanical work, i.e., of overcoming resistance. Energy may be of two kinds, either kinetic or potential. Kinetic energy is the energy matter has by virtue of its motion, while potential energy is the energy it has by virtue of its position or condition. For instance, a flying cannon ball can overcome resistance by reason of its motion. On the other hand, the same cannon ball, supported at a height from the ground, has in it, by virtue of its position, the power of acquiring motion when the support is removed, and of thus overcoming resistance. The chemical energy possessed by matter is potential energy. Coal, on account of the affinity of carbon for oxygen, will burn, and so give out heat. This heat in turn may be utilized to boil water, the steam from which in ex- panding will move the piston of an engine, and thus do mechanical work. Also the chemical energy of matter may be transformed into electric energy as in the galvanic cell. This electrical energy may be utilized to drive a motor, and thus do mechanical work. It has been shown that carbof? dioxid neither burns nor supports combustion (see Exp. 8 b ). The reason for this is that the carbon has already united with all the oxygen that it can hold; therefore no further combustion can take place. Just as the cannon ball would have to be lifted again to possess potential energy, so the carbon dioxid would have to be decomposed in order to possess chemical energy again. Carbon dioxid does not support combustion, because the affinity between carbon and oxygen is stronger than that between oxygen and most other elements. However, if an element having a stronger affinity for oxygen, as for instance potassium at a high temperature, is brought in contact with carbon dioxid, then the gas supports combustion. 146 AN ELEMENTARY EXPERIMENTAL CHEMISTRY We shall use as our unit of heat the calorie, i.e., the amount of heat necessary to raise one gram of water one degree. The number of calories given out by the burning of one gram of various substances has been determined by a number of investigators. The following table gives a few examples. Hydrogen 34180 Sulfur 2220 Carbon 8080 Zinc 1300 Phosphorus 5747 Iron 1181 We have learned that the equation 2 H 2 + 2 = 2 H 2 is in itself a statement of the law of the indestructibility of matter. But we know that, in the burning of hydrogen, heat is evolved. This equation, as it stands, tells us nothing in regard to this fact. It is very easy, however, to supplement the equation so that it shall convey the full meaning. Let the chemical symbols in an equation stand for the number of grams corresponding to its molecular weight. Thus O 2 stands for 32 grams of oxygen, HC1 stands for 35.2 grains of hydrochloric acid. By adding to the chemical equation the number of calories of heat evolved as the result of the action, we express the meaning in full. In the example of burning hydrogen we have 2 H 2 + 2 = 2 H 2 + 136800 which signifies that 4 grams of hydrogen on uniting with 32 grams of oxygen form 36 grams of water, and at the same time give out 136800 calories of heat. In the case of en- dothermic reactions, the notation is the same, except that the sign instead of the + sign is used. In addition to this, heavy type is used to represent solids, ordinary type liquids, and italics, gases. Thus, c + a, = c a, + 97000 LAWS AND THEORIES OF CHEMISTRY signifies that 12 grams of solid carbon burned in 32 grams of gaseous oxygen gives 44 grams of gaseous carbon dioxid together with 97000 calories of heat. In a great many cases where the heat liberated cannot be found by experiment, it can be found by calculation. Example. Required to find the heat liberated when car- bon is oxidized to carbon monoxid. It is found that when 12 grams of carbon is oxidized to carbon dioxid, 97000* calories are liberated, and that when 28 grams of carbon monoxid is oxidized to carbon dioxid, 136000 | calories are liberated. We then have C + 6> 2 = C O 2 + 97000 2 C O + O- 2 2 C O. 2 + 136000 By dividing the second equation by 2, we obtain 68000 cal. as the amount of heat liberated when 44 grams of carbon dioxid are formed. We divide by 2 in order to obtain the same weight of carbon dioxid as in the first equation. Sub- tracting 68000 from 97000, we evidently obtain the number of calories given out when carbon is oxidized to the monoxid, i.e., 29000 cal. We therefore have the equation C + O = CO + 29000 EXPERIMENT 21 Heat of Chemical Action Dilute 50 cc. of concentrated sulfuric acid with 250 cc. of water, and allow it to cool. Weigh a beaker of thin glass large enough to hold the acid, and place it in a larger beaker. Pack wool around and under the inner beaker in such a way that the inner beaker can be easily removed. Let this inner beaker be used as a calorimeter. Pour the acid into the * More correctly 96960. t More correctly 135920. 148 AN ELEMENTARY EXPERIMENTAL CHEMISTRY calorimeter. Clean a piece of sheet zinc about 2\ inches wide by 5 long. Take the temperature of the acid, then plunge the zinc into it. When the temperature has risen four or five degrees, remove the zinc, stir, and take the tem- perature carefully. Wash the zinc clean, dry it, and weigh again. It will be necessary to know the specific heat of the solu- tion after the zinc has been acted upon. To find this, per- form an experiment exactly like Experiment 15, Part II. with the exception of using this calorimeter and liquid just as they are, instead of the lemonade shaker and water. Calcu- late the specific heat of the liquid, knowing the specific heat of lead to be .031. Arrange your calculations as follows : To find the specific heat of the liquid. Wt. of liquid = 1 Wt. of calorimeter = c = Wt. of shot = s = Specific heat of shot = .031 Specific heat of glass = .2 Temperature liquid = t = Temperature liquid and shot = t' = Temperature of shot = t" = Let x = specific heat of the liquid. .031 s (t" - t') = .2 c (f - t) + xl (f - 1) Solve for x. To find the amount of heat liberated when one gram of zinc is dissolved in sulfuric acid. Wt. of zinc dissolved z = Wt. of water and acid = 1 = Wt. of calorimeter = c = Specific heat of glass = .2 Specific heat of liquid = s (found in foregoing calculation) Temperature before action = t = Temperature after action ='? = LAWS AND THEORIES OF CHEMISTRY 149 Let x = the number of calories liberated when one grm. of zinc is dissolved. The amount of heat evolved by the action of the acid on the zinc equals the amount of heat gained by the solution and the calorimeter. X Z = Sl (f - t) + .2 C (f - t) Solve for x and obtain the number of calories liberated when one gram of zinc is dissolved. Find 65^ in con- formity with principle laid down on page 146. Write the reaction for the chemical change together with the heat evolved. EXPERIMENT 22 Heat of Neutralization Weigh out about 25 grams of c.p. concentrated sulfuric acid. Dilute this with 200 cc. of water, and allow it to cool. Calculate (see page 142) the number of grams of sodium hydroxid necessary to neutralize the acid, take about one- sixth more than this amount, dissolve it in 200 cc. of water, and allow the solution to cool. Let the two vessels con- taining respectively the acid and alkali stand side by side until they are of the same temperature. Then pour them together in the glass calorimeter used in Exp. 21, and note the rise in temperature. Arrange your calculations as follows : Wt. of sulfuric acid = s = Wt. of sodium hydroxid = h = Wt. of calorimeter = c = Specific heat of glass = .2 Temperature before mixing = t = Temperature after mixing = t' = Wt. of water = w = Specific heat of solution = a = I5O AN ELEMENTARY EXPERIMENTAL CHEMISTRY Find a, the specific heat of the sodium sulfate solution, in the same way as you did that of zinc sulfate in Experiment 2 1 . The heat gained by the solution and calorimeter equals the heat evolved by neutralization. Heat gained by solution = (s + w + h) (f t) a Heat gained by calorimeter = .2 c (f t) Heat of neutralization = sx Therefore (s + w + h) (t '- t) a + .2 c (f - t) = sx x = 98 x = We find 98^ in conformity with the principle laid down on page 146. Write the reaction for the chemical change together with the heat evolved. EXPERIMENT 23 Heat of Solution and of Hydration a. Place in a weighed calorimeter about 350 grams of water weighed accurately to one gram. Powder exactly 40 grams of anhydrous magnesium sulfate, place it in a beaker, and cover it with a watch glass. Allow the calorimeter, containing the water, and the beaker to stand side by side until they are of the same temperature. Then pour the powdered salt into the water, and stir the solution with a thermometer, taking the temperature when all is dissc'ved. Arrange your calculations as follows : Wt. of magnesium sulfate m = Wt. of calorimeter = c = Wt. of water = w = Temperature before mixing = t = Temperature after mixing = t' = Specific heat of solution = a = Specific heat of glass = .2 LAWS AND THEORIES OF CHEMISTRY 151 Find a in the same way as in Experiments 2 1 and 2 2 . The heat here given out is made up of two different heats ; first, the heat given out by the anhydrous salt taking on water of crystallization ; second, the heat used up by the salt dissolving in water. We then have the equation : Heat gained by the solution and calorimeter = heat of hydration and of solution. Heat gained by the solution = (w -f m) (f t) a Heat gained by the calorimeter = .2 c (t' t) Heat of hydration and of solution = mx Therefore (w + m) (f - t) a + .2 c (f - 1) = mx x = I2O X = The number 120 is taken for the same reason that 98 was taken in the last experiment. (See page 146.) You should obtain as a result somewhere near the number 20280. b. Now repeat the operation, using exactly 82 grams of the salt with its water of crystallization (MgSO 4 ,7H 2 O). In this case,, calculate the heat lost in dissolving 246 grams of magnesium sulfate. You should obtain a number some- where near 3800. We see then that the heat of hydration (a positive heat) must be 3800 greater than the heat of solution (a negative heat), i.e., 24080. Dissociation. We have found (Exp. 42, Part I.) that when ammonium chlorid is heated it passes directly from the solid to the gaseous state. If the density of ammonium chlorid is found in the state of vapor, the number obtained is 13.34. Now, according to the hypothesis of Avogadro, the number I$2 AN ELEMENTARY EXPERIMENTAL CHEMISTRY should be 26.69, smce the molecular weight is 53.38. The densities of hydrochloric acid gas and ammonia are, how- ever, 18.18 and 8.5 respectively; which, if added, give 26.68, showing that undoubtedly, when ammonium chlorid is vapor- ized, it is not made up of molecules of ammonium chlorid, but of a mixture of molecules of hydrogen chlorid and ammonia. When allowed to cool, these molecules reunite and form molecules of the original compound. This pheno- menon is called dissociation. EXPERIMENT 24 Dissociation Procure a glass tube about 25 cm. long and 2 cm. bore, and close the ends by means of two closely fitting corks, through both of which passes the stem of a clay tobacco pipe. Place in the center of the tube a piece of crystal- lized ammonium chlorid, and at the ends, next to the corks, place pieces of moist blue litmus paper, flat against the glass. Connect one end of the pipe with a pair of bellows. Heat the am- monium chlorid with the flame of a Bunsen burner, at the same time gently forcing air through the pipe stem. The two gases into which the ammonium chlorid has been dissociated will pass through the porous pipe stem in different quantities; and, LAWS AND THEORIES OF CHEMISTRY 153 by holding a piece of moist red litmus paper at the opening of the pipe, the presence of ammonia gas will be shown, while the paper inside the tube will show the presence of free hydrochloric acid gas. Dissociation in Solutions. Dissociation takes place not only in gaseous compounds but also in solutions. In these cases it is probable, and it is believed, that the molecules are con- tinually breaking up into atoms or groups of atoms, and then reuniting again. These atoms or groups of atoms are called ions. In the case of a solution of hydrochloric acid, the ions would be H and Cl, and in the case of a solution of copper sulfate they would be Cu and SO 4 . The ions are believed to be present in greater numbers in dilute solutions than in strong ones. The theory is that the difference be- tween ions and simple atoms is, that the ions are charged electrically ; thus, in a solution of copper sulfate, the Cu ions are charged positively, and the SO 4 ions are charged negatively. EXPERIMENT 25 Dissociation in Liquids Prepare four strong aqueous solutions (10 cc. each) as fol- lows : one of copper nitrate, one of copper chlorid, one of sodium chlorid, and one of sodium nitrate. In a test tube, mix the copper nitrate and the sodium chlorid solutions. In another, mix the sodium nitrate and copper chlorid. Note that the two mixtures are the same. Evidently each mixture contains copper nitrate, copper chlorid, sodium nitrate, and sodium chlorid. Besides these compounds there must be present the ions of Cu, Na, NO 3 and Cl. To destroy the equilibrium between these, add to one of the mixtures a little i$4 AN ELEMENTARY EXPERIMENTAL CMEMISTRV powdered sodium chlorid and shake. The additional green color shows the presence of more copper chlorid, the forma- tion of which has of course necessitated a rearrangement of the compounds and ions. Organic Chemistry. It was formerly thought that those compounds formed by the chemical elements in living bodies were not bound by the same laws as those of the inorganic world. The principle of life was supposed in some mys- terious way to govern them, and it was thought that they could not be prepared artificially. On this account, that branch of chemistry which deals with such compounds was named organic chemistry. When Wohler, a German chemist, succeeded in making artificially the organic compound urea, this theory was overthrown. After many other compounds were thus made, it became no longer tenable, and we now class all chemical compounds as dependent upon the same laws. It happens that the element carbon is the most fre- quently present in the so-called organic compounds ; so it were better to name this branch of the science the chemistry of the compounds of carbon. However, the name organic chemistry has clung to it, and probably always will. EXPERIMENT 26 Burning of Organic Matter; Dry Distillation a. Burn a number of organic substances such as wood, alcohol, a candle, kerosene, etc., under a cold bell-jar, and notice the formation of water. On the end of a glass rod hold a drop of lime-water in the jar, and prove the presence of carbon dioxid. b. To a hard-glass tube about 8 mm. in diameter con- taining a piece of wood, attach by means of a one-holed cork LAWS AND THEORIES OF CHEMISTRY 155 a glass exit tube. Heat the wood, and ignite the gas that escapes. What remains in the tube ? Is the combustion of the wood partial or complete ? Alcohol. Sugar is a compound made up of carbon, hydro- gen, and oxygen (CgH^Oe). Whenever a juice containing sugar is left in the open air, it decomposes, giving off carbon dioxid gas, and forming a new compound called alcohol C 2 H 6 O. C 6 Hi2 6 = 2 C 2 He + 2 C 2 This action, called fermentation, is caused by a small organized body (in this case vegetable) called a ferment. EXPERIMENT 27 Alcohol In a 500 cc. flask, dissolve 40 grms. of grape sugar in 250 cc. of water. Add to this a little brewer's yeast, after con- necting the flask with a wash bottle containing lime water. Notice the evolution of carbon dioxid, as proved by the milky color of the lime water. After the apparatus has stood long enough for the action to cease, re- move the flask and place it in a water bath. Connect it with condenser, and allow the alcohol to distill over. The condenser may be made as follows. Fit each end of a glass tube (about 60 cm. long and 2.5 or 3 cm. bore) with a two-holed rubber stopper. Let a glass tube about 75 cm. long and 5 mm. bore extend through the tube, and through both stoppers. In the other holes of the 156 AN ELEMENTARY EXPERIMENTAL CHEMISTRY stoppers, fit pieces of glass tubing about 10 cm. long and bent at right angles. Fasten the apparatus to a stand, and incline it at a slight angle. Allow cold water by means of rubber tubing to enter the lower end and escape from the higher. EXPERIMENT 28 Saponification In a porcelain dish, boil for an hour about one-eighth of a pound of lard together with a solution of sodium hydroxid (20 grms. to 125 cc. of water). While the mixture is cooling, add a strong solution of salt. The substance that solidifies on the surface of the liquid is soap. Remark. Soaps are the alkali salts of certain fatty organic acids such as stearic, C 18 H 36 O 2 , and palmitic C 16 H 32 O 2 . The calcium and magnesium salts of these acids are insoluble in water; hence, when soap is used with hard water (see Exp. 35, Part I.), these compounds appear on the surface of the water as a scum. PART III. HISTORY, OCCURRENCE AND INDUSTRIAL APPLICATIONS OF THE PRINCIPAL ELEMENTS AND COMPOUNDS PART III. HISTORY, OCCURRENCE AND INDUSTRIAL APPLICATIONS OF THE PRINCIPAL ELEMENTS AND COMPOUNDS OXYGEN History. Until 1774, the air was believed to be a simple substance. In that year, the investigations of Priestley, Rutherford, and Scheele proved that it was a mixture of two different gases. By heating mercuric oxid (see Exp. 5, Part I.), Priestley proved that this substance was composed of a gas and metallic mercury. The gas thus obtained was shown to be the same as one of the constituents of the air. In 1805, Gay Lussac proved that water was composed of two volumes of hydrogen and one of oxygen. (See Exp. 6 e, Part I.) The name oxygen (6v's, sour, yewaw, I produce) was given it by Lavoisier. Occurrence. Oxygen is the most abundant element in nature. It constitutes 23 per cent of the atmosphere, 88.88 per cent of water, and from 44 to 48 per cent of the crust of the earth. The ores of almost all of the metals occur in the earth as oxids, or as other compounds containing oxygen. Industrial Applications. Oxygen is used extensively for medical purposes, and in connection with hydrogen in the oxy-hydrogen blowpipe. It is usually prepared in large quantities by heating potassium chlorate (KC1O 8 ), or by obtaining it from the atmosphere. In the latter method, I6O AN ELEMENTARY EXPERIMENTAL CHEMISTRY advantage is taken of the fact that at a dull red heat barium oxid (BaO) takes on oxygen, becoming barium peroxid (BaO 2 ), and that at a still higher temperature the oxygen thus absorbed is given off, leaving the original oxid. Theoret- ically this process could be continued indefinitely. In prac- tice, however, the barium oxid is not used indefinitely, since it becomes gradually less efficient in its action. The changes are 2 Ba + 2 = 2 Ba 2 2 Ba 2 = 2 Ba + 2 It has been found that this reaction can be accomplished at a constant temperature by changing the pressure. HYDROGEN History. In the sixteenth century. Paracelsus obtained an inflammable gas by treating metals with certain acids. To this gas, Cavendish gave the name " Inflammable Air " in 1766. Later he proved that, when this gas was united with oxygen, water was formed. Lavoisier confirmed this, and gave to the gas the name hydrogen ({SStop, water, yevraw, I produce). Occurrence. Hydrogen occurs in the free state in the atmosphere of the sun, in small quantities mixed with other gases in volcanic eruptions, and sometimes in oil wells. It is also found occluded in meteoric iron and certain iron ores. In chemical combination, it is most widely distributed as water. It also occurs in combination with a number of the non-metals, and as a part of almost all organic compounds. Industrial Applications. From its great lightness, hydrogen is valuable for filling balloons. It has often been made for military balloons by the steam and red-hot iron process. (See Exp. 15 b, Part I.) HISTORY, OCCURRENCE AND APPLICATIONS l6l The great amount of heat given out by the oxidation of hydrogen has made it valuable in melting refractory metals such as platinum. This is done by means of a very simple piece of apparatus called the oxy-hydrogen blowpipe. It is simply a tube within a tube, the tip being made of platinum. The inner tube is connected with a gas holder containing hydrogen, while the other is connected with a similar one containing oxygen. The hydrogen is first turned on and ignited. The oxygen is then turned on until the flame burns quietly. WATER Water is one of the most abundant and most widely distributed compounds. From the fact that it is a solvent for so many substances, it is never found pure. When pure, it is tasteless and colorless. In large quantities, it often has a greenish or a bluish color. The great reservoir of course is the sea, from which the water that is pre- cipitated upon the land as rain, hail, or snow, originally evaporated. Sea water contains in solution about 3^ per cent of solid matter, most of which is sodium chlorid. The water of lakes and rivers contains various substances depend- ing upon the locality. The most common substances are cal- cium carbonate and calcium sulfate. (See Exp. 31, Part I.) Water is so widely used that it is hardly necessary to enumerate many of the ways in which it is of value. The principal uses of water, however, are for drinking, 1 62 AN ELEMENTARY EXPERIMENTAL CHEMISTRY washing, and culinary purposes in the household, and for boiling and solvent purposes in the arts. Good drinking water should be as free as possible from sewage contamina- tion and decaying organic matter. These contaminations are not only injurious in themselves, but they render the water more habitable for the germs of typhoid fever, cholera, and other malignant diseases. The disease germs are almost always carried into water by means of sewage, hence water containing the slightest trace of sewage should be especially avoided. Water that contains organic impurities is usually yellowish in color, and has a disagreeable odor. The water supply of large cities is freed from impurities by a process of filtration through several feet of sand and broken stone extending over a large area. At least two such filters are used, and one is kept empty while the other is in use. The disease germs that are removed by one filter are destroyed by oxidation, when the water is diverted into the other. For boiling and solvent purposes in the arts, soft water is preferable. For washing purposes, hard water may be made soft by the addition of sodium carbonate. HYDROGEN DIOXID, H 2 O 2 This compound was first discovered by Thenard in 1818. He obtained it by treating barium peroxid with dilute hydro- chloric acid. Ba 2 + 2H Cl = H 2 2 + Ba C1 2 It is formed in minute traces in the atmosphere, and is sometimes produced simultaneously with the preparation of ozone. It is an oily, colorless liquid, having a bitter taste. It is usually used in dilute solutions. Industrial Applications. Because of the fact that hydro- HISTORY, OCCURRENCE AND APPLICATIONS 163 gen peroxid gives up part of its oxygen readily, it is very valuable as an oxidizing agent. It has been extensively used for bleaching hair, giving to dark hair the well-known light flaxen tint. It is also used for restoring the colors to old paintings. White paint is composed largely of lead car- bonate, which gradually darkens in time. This occurs on account of the action of sulfur, which forms a black lead sul- fid. If hydrogen peroxid is applied to such a discolored painting, the lead sulfid is converted into lead sulfate, thus restoring in a great measure the original color. Hydrogen peroxid is also used extensively in medicine and as a disinfectant. THE HALOGENS The elements chlorin, bromin, iodin, and fluorin may be grouped together, since their properties are in many ways similar. This group is called the halogen group (oAos, salt, yewda), I produce). CHLORIN History. Chlorin gas was first obtained and studied by Scheele in 1774. He obtained it from hydrochloric acid and manganese dioxid. Its elementary character was proved by Davy in 1810, and he named it chlorin from ^Xwpos, meaning greenish yellow. Occurrence. On account of its strong affinity for other elements, chlorin is not found in the free state. It occurs in combination chiefly with the alkali metals in sea water, and as rock salt in various localities (chiefly at Stassfurt, Ger- many). At Syracuse, New York, salt occurs as a brine at a depth of from 200 to 400 feet below the surface of the earth. Industrial Applications. From its strong affinity for hydro- gen, chlorin is used in the arts in the process of bleaching 164 AN ELEMENTARY EXPERIMENTAL CHEMISTRY cloth. In bleaching cloth, it is necessary for the cloth to be moist. The change is explained by the fact that the chlorin unites with the hydrogen of the water, leaving nascent oxygen, which in turn oxidizes the vegetable coloring matter into colorless compounds. Black colors from carbon, as for instance printer's ink, cannot be bleached. Ordinary writing ink, which is a compound of organic acids and iron, is readily decolorized. Large quantities of chlorin are also used for disinfectant purposes. Manufacture. It will be remembered that chlorin is made in the laboratory (see Exp. 27 b) by treating manganese dioxid with hydrochloric acid, thus : Mn 2 + 4 H Cl = Mn Cls + 2H 2 + C1 2 Since manganese dioxid is the costly material used, it is necessary in the manufacture of chlorin on a large scale to save the manganese chlorid formed, and in some way change it back to the oxid. In the Weldon process this is done. The chlorid is treated with calcium hydrate. The mixture is then heated, and a current of air is blown through. A complicated reaction that need not be entered into here ensues, and the manganese is oxidized to peroxid. Thus the original manganese can be used over again, making chlorin a cheap commercial product. Other processes for the manufacture of chlorin are also in use. HYDROCHLORIC ACID History and Occurrence. Hydrochloric acid was known to Arabian alchemists in aqua regia. Basil Valentine in the fifteenth century wrote of an acid which he called " spirit of salt," obtained from oil of vitriol and common salt. Priestley HISTORY, OCCURRENCE AND APPLICATIONS 165 obtained the gas by collecting it over mercury instead of water, and called it "marine add air" In 1810, Davy showed that this gas was composed of hydrogen and chlorin, and that it was not a compound of oxygen as had formerly been supposed. Hydrochloric acid occurs in the gases which issue from some volcanoes, especially Vesuvius. It is also found in some South American rivers whose sources are in the vol- canic districts of the Andes. Industrial Applications and Manufacture. Hydrochloric acid is used in the manufacture of chlorin, ammonium chlorid, and tin chlorid, the last being extensively used bydyers. In the manufacture of sodium sulfate, large quantities of hydrochloric acid are obtained as a by-product (see Exp. 23, Part I.). The acid fumes pass through a flue to brick chambers filled with coke or broken brick through which water is passing. The water absorbs the gas, and this solu- tion, when collected, is the hydrochloric acid of commerce. OXIDS AND OXY-ACIDS OF CHLORIN There are three oxids of chlorin. Chlorin monoxid, C1 2 Chlorin trioxid, C1 2 3 Chlorin peroxid, Cl 02 They are all unstable compounds. The oxy-acids of chlorin are Hypochlorous acid, H Cl Chlorous acid, H Cl 2 Chloric acid, H Cl 3 Perchloric acid, H Cl 2 These only exist in aqueous solutions. Hypochlorous and chloric acids form salts that are important commercial products. 1 66 AN ELEMENTARY EXPERIMENTAL CHEMISTRY BLEACHING POWDER If chlorin is allowed to pass into chambers containing quantities of freshly slaked lime, it is absorbed by the lime, and a compound commonly called " chlorid of lime " * is formed. This compound has strong bleaching properties because it slowly gives off chlorin. Considerable time has been given to the study of the composition of this substance, but at present no entirely satisfactory explanation has been given. It was first thought to be a salt of hypochlorous acid, but it does not show the proper percentage of chlorin. The best authority on the subject affirms that its composi- tion is CaOCl 2 , and calls it chloro-hypochlorite. It is also called calcium oxychlorid. " Chlorid of lime " is used in enormous quantities as a disinfectant, and for bleaching purposes. In bleaching calico and paper pulp, a two per cent solution of bleaching powder is used. The soaked product is then placed in a dilute solution of sulfuric acid, which liberates the chlorin more freely. When thoroughly bleached, the product is treated with sodium sulfite (called anti-chlor), which removes all traces of- unused chlorin. This is done because otherwise the chlorin would slowly attack the fiber of the cloth or paper. POTASSIUM CHLORATE The potassium salt of chloric acid is a very important product. It can be made by treating warm potassium hy- droxid with chlorin gas. The reaction is as follows : 3 C1 2 +6KOH = KC10 3 + 5KC1 + 3H 2 By this method, five of the six molecules of potassium * If bleaching powder is treated with an acid, the chlorin is given off again. HISTORY, OCCURRENCE AND APPLICATIONS l6/ hydroxid are converted into potassium chlorid ; so it is better first to form the corresponding calcium salts, and then add to the solution enough potassium chlorid to change the cal- cium chlorate to potassium chlorate. The use of potassium- hydroxid, which is the expensive compound in the first method, is avoided in the second. The reactions are 6 Ca (0 H) 2 + 6 C1 2 = Ca (Cl 3 ) 2 + 5 Ca C1 2 + 6 Ho and Ca (Cl s, light, 1 82 AN ELEMENTARY EXPERIMENTAL CHEMISTRY and (j>epa>, I bear. The element never occurs free in nature. It is widely distributed, however, in combination. Many rocks contain phosphorus compounds. When these rocks disintegrate, the soils that are formed become very produc- tive on account of the presence of phosphorus. Vegetation, in turn, is transformed into animal matter. Here the phos- phorus appears mostly in the bones as calcium phosphate. A fossil substance called " caprolites " also contains calcium phosphate, and it is from this and from bones that the phosphorus of commerce is derived. Manufacture. In the manufacture of phosphorus from bones, the first step is to remove the non-phosphorus-bear- ing organic matter. This is done by burning. The ash is then treated with sulfuric acid, which changes the insoluble calcium phosphate to soluble acid calcium phosphate, accord- ing to the following reaction : Ca 3 (P 4 ) 2 + 2 H 2 S 4 = Ca H 4 (P 4 ) 2 + 2 Ca S 4 After the calcium sulfate has been removed, the acid calcium phosphate solution is evaporated to a syrup and mixed into a paste with charcoal powder. This is heated to redness in an earthenware retort, whose mouth dips under water. Water is at first driven off thus : Ca H 4 (P 4 ) 2 = Ca (P 3 ) 2 + 2 H 2 The metaphosphate thus formed then reacts with the char- coal as follows : 3 Ca (P 3 ) 2 + 10 C = P 4 + Ca 3 (P 4 ) 2 + 10 C The phosphorus distills over into the water, is collected, and is purified by redistillation. Sometimes sand is added to- gether with charcoal, in which case all the phosphorus is obtained. The phosphorus obtained by the above method is the yel- low variety. The red amorphous phosphorus is made by HISTORY, OCCURRENCE AND APPLICATIONS 183 heating, for a number of days, the yellow variety at 240 C. in an iron vessel having only a small opening. Not all the yellow phosphorus changes over. The unchanged phos- phorus is dissolved out of the cooled mixture by carbon bisulfid, or is removed by means of sodium hydroxid which acts upon the yellow phosphorus and forms phosphoretted hydrogen. Industrial Applications. Phosphorus is used in the arts principally in the manufacture of matches. The old friction matches were made by first dipping the wood into melted sulfur, cooling it, and then tipping it with a mixture of phosphorus, glue, and some oxidizing agent. The so-called safety matches do not contain phosphorus. Instead of this they are tipped with antimony sulfid. The box is painted with a preparation of red amorphous phosphorus, antimony sulfid, and glue. Sometimes manganese dioxid or some other oxidizing agent is used in this mixture. Phosphin. There are three compounds of phosphorus and hydrogen, only one of which, phosphin (PH 3 ), will be noted here. If phosphorus is heated in a solution of sodium or potassium hydroxid, a gas is liberated that ignites spon- taneously. If the gas is led under water, the bubbles on reaching the surface burst into flame and form white rings of phosphorus pentoxid. The gas is colorless, has a very disagreeable odor, and is poisonous. It forms a class of substances, called phosphonium compounds, analogous to those formed by ammonium. Phosphoric Acid. The phosphoric acid of commerce is made from bones. The main value of phosphoric acid lies in its salts, which are the principal constituents of fertilizers. For the purpose of making fertilizer, bone ash is treated with sulfuric acid to obtain the soluble acid calcium phos- phate, CaH 4 (P0 4 ) 2 . . 1 84 AN ELEMENTARY EXPERIMENTAL CHEMISTRY ARSENIC History and Occurrence. Arsenic compounds were known to the alchemists, who thought they could use them in trans- muting the metals. Brandt, in 1773, first showed that white arsenic was the calx of a metal. Later, its position as an oxid was established. Arsenic occurs in the free state in nature, but usually it is found in combination with iron, cobalt, nickel, or sulfur. Applications. Arsenic is used in the manufacture of cer- tain pigments. In the manufacture of glass, arsenic oxid is used to remove the green tint given by ferrous hydroxid, which by its action becomes ferric oxid. In the manufacture of shot, arsenic has the effect of hardening the lead. In medicine, it is sometimes, used as a tonic. For preserving skins, it is valuable to the taxidermist. In the form of Paris green (CuHAsO 3 ), it is valuable as a poison for the de- struction of insects. ANTIMONY History. Mention is made in the Scriptures of the metal we call antimony. In early times the sulfid was used by the women of the East to paint their eyebrows. Pliny called it stibium, although in Latin it was also known as antimonium. Occurrence and Applications. While antimony occurs some- times free in nature, it is usually found as the sulfid (Stib- nite, Sb 2 S 3 ). When the ore is roasted in the presence of air, the oxid (Sb 2 O 3 ) is formed. This, mixed with carbon and ignited strongly, gives the metal. Antimony is used chiefly in making alloys. It hardens the alloy, and also, by its property of expanding when cooled, makes it invaluable in the manufacture of type metal. Type HISTORY, OCCURRENCE AND APPLICATIONS 185 metal is composed of about two parts lead, and one part each of antimony and tin. Antimony is used also in making white metal, pewter, and Britannia metal. Finely divided anti- mony, prepared by zinc from the chlorid, is called antimony black, and is used for giving a metallic appearance to plaster casts, statues, etc. In medicine its compound, potassium antimony tartarate (tartar emetic), is used as an emetic. BORON History and Occurrence. Gay Lussac, Thenard, and Sir Humphry Davy obtained boron in the elementary state in 1808. There are two varieties, the amorphous and the crys- talline boron. It never occurs free, but is found as boric acid (H 3 BO 3 ) in Tuscany, and as sodium salts in California and Thibet. Boric Acid, ff z BO z . In Tuscany, jets of steam contain- ing boric acid escape to the surface of the earth from sub- terranean sources. Brick basins are built around these steam jets. The heat causes the solution obtained to evaporate. When sufficiently strong, the solution is removed and allowed to crystallize. Uses. Boric acid is a valuable antiseptic, and for this 1 86 AN ELEMENTARY EXPERIMENTAL CHEMISTRY reason is used extensively in surgery. It is also sometimes used in preserving perishable foods. Borax, Na^B^O 1 is a valuable compound, and is much used as a flux by tin-smiths and copper-workers. It gives a clean, bright, metallic surface on account of its property of dissolved metallic oxids. CARBON History. Carbon has of course been known from ancient times, but its allotropic forms were not understood until the end of the i8th century, and its relation to organic chemistry not until somewhat later. In the form of graphite, it was known to the alchemists. Graphite pencils were first made in 1565. In 1772, Lavoisier showed that diamond and charcoal were chemically identical. He burned a diamond and obtained carbon dioxid. Tennant, in 1776, proved that like weights of charcoal, graphite, and diamond give like weights of carbon dioxid. Occurrence. Carbon occurs as diamond principally in India, South Africa, and Brazil. It is usually transparent or slightly tinged with yellow, although it is sometimes found red, green, blue, or even black. As graphite, carbon occurs widely distributed, being found mostly in England, Siberia, Ceylon, Canada, New York, and California. Graphite is used in the manufacture of lead pencils, cruci- bles, and various lubricants. It is also used in foundries for facings, and in electrotyping. Amorphous Carbon. Coal. As mineral coal, carbon is found in almost all countries of the world. The most, how- ever, is found in England and the United States. Coal is all that is left of the great primeval forest that covered the. earth long before the advent of man. When vegetable HISTORY, OCCURRENCE AND APPLICATIONS l8/ matter decays in the absence of air, in the earth, or under water, it gives off gases, and a substance similar to coal remains. In some parts of the earth, the trees of the ancient forests were buried by earthy material, and underwent the process of decay for millions of years. Upon the complete- ness of this process depend the different kinds of coal. Where the decomposition has been most thorough, anthra- cite coal results. Bituminous coal is rich in hydrocarbons, which it gives off when heated. Cannel coal is a variety of bituminous coal, and is especially rich in hydrocarbons. Brown coal and peat belong to a later geological period than the others mentioned. Charcoal. Charcoal is obtained by the partial combus- tion of wood. The wood is arranged in a pile, and covered with earth. It is then ignited, and allowed to burn. On account of being protected from the air, the combustion is imperfect, and only the volatile part of the wood is driven off. Animal Charcoal. By heating bones and other kinds of animal refuse in iron retorts, a variety of charcoal tha.t con- tains calcium phosphate is obtained. It is largely used as a filter for removing vegetable coloring matters from liquids. Coke and Gas Carbon. In the retorts of the gas works, after the volatile products have been driven off from the coal, there remains behind a gray, porous solid that is called coke. It is also made by burning coal in ovens so as to burn out only the volatile products. In the upper parts of the gas retorts, there remains a dull black mass known as gas retort carbon. It is used in making carbon plates for electric batteries, and formerly for making pencils for arc- light lamps. Gas retort carbon is the purest form of amor- phous carbon. Lamp black is made by burning turpentine and other oils 1 88 AN ELEMENTARY EXPERIMENTAL CHEMISTRY rich in hydrocarbons, and collecting the smoke. It is used chiefly in printer's ink and black paint. Hydrocarbons. In Pennsylvania, in Ohio, and in the Cau- casus, an oily liquid called petroleum is found in the earth. This liquid is composed mostly of a mixture of hydrocar- bons. These are partly gaseous, partly liquid, and partly solid. The petroleum is distilled, and the resulting mixtures are afterwards washed with water and alkali. Gasolene, naphtha, benzine, kerosene, and paraffin, are among the principal substances obtained. These are all mixtures of various hydrocarbons. Natural gas occurs in the earth in large quantities, usually near coal beds. Its formation is supposed to be due to the dry distillation of coal in the interior of the earth. The hydrocarbons form a definite series of compounds, each hydrocarbon of the series differing from the next by an atom of carbon and two atoms of hydrogen, i.e., by CH 2 . The first hydrocarbons of the simplest groups are Methane, C H 4 (See Exp. 33 e. Part I.) Ethylene, C 2 H 4 (See Exp. 33 /, Part I.) Actylene, C 2 H 2 METHANE OR MARSH GAS History and Occurrence. This gas is mentioned by Pliny, and Basil Valentine speaks of its presence in mines. In 1785, Berthollet proved that methane contained both hydro- gen and carbon, and in 1805 Henry showed the difference between methane and ethylene. Methane occurs in nature in mines, and wherever vege- table matter is decaying under water ; hence its common name, marsh gas. Miners call it fire damp. Natural gas is rich in methane. HISTORY, OCCURRENCE AND APPLICATIONS 189 ETHYLENE History and Occurrence. In the lyth century, Becher dis- covered ethylene by heating alcohol with sulfuric acid, but it was not until Henry took up the study of methane and ethylene that its composition was understood. It occurs in the gases that emanate from oil wells, and is the principal constituent of coal gas. ACETYLENE History and Application. Acetylene was discovered by Edmund Davy in 1836. It is formed by the incomplete combustion of other hydrocarbons, such as ethylene, coal gas, etc. For instance, when the flame " backs down " in a Bunsen burner, acetylene is formed. Acetylene is rich in carbon, burns with a bright, white flame, and, if it can be produced cheaply enough and regu- lated with perfect safety, is likely to become one of the chief sources of illumination. It is now made by the action of calcium carbid \^CaC 2 ) upon water. Ca C 2 + 2 H 2 = C 2 H 2 + Ca (0 H) 2 ILLUMINATING GAS William Murdock, a Scotchman, first saw the practicabil- ity of making illuminating gas from coal. This was in 1792, but London was not lighted by gas until 1812, nor Paris until 1815. When coal is heated in the absence of air, three classes of products are formed, 1. Illuminating gas. 2. Coal tar, a thick, oily, strong-smelling liquid. AN ELEMENTARY EXPERIMENTAL CHEMISTRY 3. Ammoniacal liquor, containing ammonia and various other compounds in solution. When the coal is heated in the gas retorts, the three classes of products above mentioned pass through a pipe that dips below water in another large pipe called the hydraulic main. The water absorbs some of the soluble gases, and takes up some of the ammoniacal liquor and tarry products. From the hydraulic main, the gas passes through a series of upright iron pipes beneath which is water. As the gas passes through these pipes, it cools, and another portion of the tarry products and ammoniacal liq- uor is deposited, which runs down into HISTORY, OCCURRENCE AND APPLICATIONS IQI the water. From this pipe, the gas passes to a series of towers, called scrubbers, filled with coke. It passes into the bottom of one tower, through the coke, and then at the top meets a spray of ammonia water. In like manner, it passes through the next tower, and so on. The ammonia water removes most of the hydrogen sulfid. There is still left in the gas some hydrogen sulfid, carbon bisulfid, and carbon dioxid. To remove these, the gas passes through a series of chambers containing lime or hydrated ferric oxid. The gas is then ready for use, and passes into the gas holder. CARBON MONOXID History and Occurrence. This compound (see Exp. 33, Part I.) was studied by various investigators during the last part of the i8th century, but it was some time before its true character was explained. It does not occur free in nature, but it is" formed wherever carbon burns in an in- sufficient supply of air. \ Applications. The principal applications of carbon mon- oxid are in the manufacture of water gas and in the reduc- tion of metals from their ores. WATER GAS In the manufacture of water gas superheated steam is passed over red hot anthracite coal, giving the gases carbon monoxid and hydrogen. Both of these gases burn with a blue flame, hence they must be enriched with some gas that burns brightly. 192 AN ELEMENTARY EXPERIMENTAL CHEMISTRY CARBON DIOXID History and Occurrence. Van Helmont, one of the early investigators, held, as early as the beginning of the iyth century, that this gas was different from air. He called it "gas sylvestre." Its chemical nature was first determined by Lavoisier. Carbon dioxid occurs free in the air and in combination with metals as carbonates. It is also found in some mineral waters. Manufacture and Applications. For commercial purposes, carbon dioxid is made by treating sodium carbonate or cal- cium carbonate with acids. Under 38.5 atmospheres pres- sure and at o C., this gas condenses to a liquid. It is then stored in steel cylinders, from which it can be taken at will. Its principal use is in the manufacture of aerated beverages. CARBON BISULFII) This compound (CS 2 ) does not occur in nature, but is made by leading vapors of sulfur over red-hot charcoal. It is extensively used in the arts as a solvent for rubber, phos- phorus, sulfur, iodin, and many oils and gums. CYANOGEN By heating mercuric cyan id, a colorless gas is obtained that burns with a beautiful purple flame, and has the odor of peach kernels. Hg (C N), = (C N) 2 + Hg It combines with hydrogen, forming hydrocyanic acid (HCN), commonly called prussic acid. This is such a deadly poison that it is used only in dilute solutions. HISTORY, OCCURRENCE AND APPLICATIONS 193 SILICON History and Preparation. Silica, or sand, was believed to be a compound body long before Berzelius first obtained impure silicon by fusing together iron, carbon, and silica. It is best prepared by heating potassium silico-fluorid and metallic potassium in an iron tube. K 2 Si F 6 + K 4 = 6 K F + Si Silicon is a brown powder that, when heated in the air, burns to silicon dioxid. If this amorphous silicon is fused with zinc, it forms dark glittering crystals of silicon that may- be obtained by dissolving away the zinc with an acid. Occurrence. Next to oxygen, silicon is the most abundant element in nature. It occurs in combination with oxygen in quartz, and in the form of silicates of the metals. SILICON DIOXID Quartz is the purest form of silicon dioxid. Sand and sandstone are other forms. Silicon dioxid is used in the arts in the manufacture of glass, porcelain, and in potter)-. GLASS The Egyptians are believed to be the first people that manufactured glass. On Egyptian tombs are found pictures of glass blowers carrying on their vocation. During the Middle Ages, Venice was famous for its glass manufactures ; but after their decay the art passed to the workmen of Bohemia. The art of glass making has steadily improved, until at present some of our most beautiful and marvelous works of art are due to the glass-worker's skill. Glass is a IQ4 AN ELEMENTARY EXPERIMENTAL CHEMISTRY mixture of various silicates, especially of the alkalis and alkaline earths. The various kinds of glass are divided into classes accord- ing to their ingredients. 1. Bohemian glass. 2. Crown or window glass. , 3. Common green or bottle glass. 4. Flint glass or crystal. Bohe?nian glass is a silicate of potash and lime, is fusible with difficulty, and withstands chemical reagents better than any other kind. Crown glass is a silicate of soda and lime. It is more readily fusible, and is more easily acted upon by chemicals than is Bohemian glass. Bottle glass is a silicate of soda and lime, mixed with the oxids of aluminum and iron. The green color is due to the iron. This color varies from green to brown, depending upon impurities. Fhnt glass is potash-lead silicate. It is the softest kind of glass, and has a bright luster and high refractive power. On this account, it is used in making lenses for optical in- struments. Glass can be colored by means of various metallic oxids. Gold compounds give it a beautiful ruby tint. Cuprous oxid colors glass an intense red, while cupric oxid colors it green. Cobalt gives blue, manganese gives violet. Black is obtained by the addition of sesqui-oxid of iridium. PORCELAIN Porcelain differs from glass in that it is made of kaolin, a silicate of aluminum, Al 2 (SiO 3 ) 3 . Porcelain is glazed after it has been " fired " for the first time. By one method, HISTORY, OCCURRENCE AND APPLICATIONS 1 95 the ware is dipped into a glazing material, usually finely pow- dered quartz and feldspar, is then dried, placed in earthen- ware pots, and heated in a furnace. The glaze fuses and spreads over the surface of the ware. The ware is then allowed to cool in the furnace to anneal it; otherwise it would be brittle on account of the unequal tension of its dif- ferent parts. Ordinary pottery is simply baked, and not glazed. Bricks are simply baked clay. The red color is due to the presence of iron silicate. POTASSIUM History and Occurrence. Until 1807, the alkalis were believed to be simple substances. The discovery of potas- sium by Davy dispelled this idea, and the true character of these metals was shown. Davy isolated potassium by de- composing potash by means of a strong electric current. Potassium is found widely distributed in na- ture. It forms from 2 to 3 per cent of our granite rocks. In combination with chlorin, as potassium chlorid, it is found in the earth in considerable de- posits. No vegetable growth is possible with- out potassium ; hence all fruitful soils contain it. Manufacture. Potassium is obtained by reducing potas- sium carbonate with carbon. A mixture of potassium tarta- rate and potassium carbonate is first heated. The tartarate decomposes into carbonate and carbon. This mixture is then placed in a wrought iron mercury bottle, and is heated. A, Grate. B, Retort. C, Receiver. 196 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Potassium and carbon monoxid results : K 2 CO 3 + 2 C = K 2 + 3 CO. These two substances readily form an explosive compound at high temperatures, but this is now prevented by cooling the potassium as fast as it is produced. The retort is connected, by means of a short tube, with a very shallow receiver consisting of two thin iron plates clamped together. The potassium here condenses to a liquid, and is transferred to a vessel containing naphtha, as soon as it solidifies. (See ill. p. 195.) POTASSIUM CARBONATE Potassium carbonate was formerly manufactured almost entirely from wood ashes. The ashes were treated with water, filtered, and the solution evaporated. Nowadays, in addition to the above source, most of the " potash " of com- merce is obtained from three sources : (a) from beet-root, () from the sweat of sheep, and (c] from potassium sulfate. (a) The molasses from beet-root sugar is allowed to ferment, and is then evaporated. A black mass containing the potash is then obtained. (//) One-third of the weight of the sweat of sheep is potassium compounds. The washings of sheep wool are evaporated to dryness, and then heated in retorts. What is left is carbon and various potassium salts, which are then separated, (c) Potassium sulfate is obtained as a by- product in many processes. It is converted to the carbonate by a process noted later under the manufacture of sodium carbonate. POTASSIUM HYDROXID Potassium hydroxid is obtained from the carbonate by treating it with slaked lime. K 2 C 3 + Ca (0 H) 2 = 2 K H + Ca C 3 To a hot solution of potassium carbonate, lime is added HISTORY, OCCURRENCE AND APPLICATIONS until, after the calcium carbonate formed has settled, the addition of hydrochloric acid causes no effervescence. The liquid is then drawn off, evaporated, and finally heated to redness in silver crucibles. From the crucibles, it is run into cylindrical molds, and thus cast into sticks. Potassium hydroxid, together with sodium hydroxid, are used in the manufacture of soap. Soap is a salt of an alkali and an organic acid. POTASSIUM NITRATE This substance is found in nature exuding from the soil in warm climates. (See page 179.) It is also obtained artifi- cially on what are called " niter plantations." A mound of chalky soil is built upon a foundation of clay. This is kept moist with the nitrogenous refuse from stables and sewers. In time the organic matter oxidizes, and, uniting with the alkalis, forms nitrates. The soil is washed from time to time to obtain these salts. From the solution thus obtained, crude potassium nitrate is separated by crystallization. Application. The- chief uses of potassium nitrate are in the manufacture of nitric acid and gunpowder. Gunpou>der. Gunpowder is a mixture of potassium nitrate, charcoal, and sulfur. The percentages are approximately, potassium nitrate 75, charcoal 15, and sulfur 10. The oxygen necessary for the burning of the charcoal and sulfur is furnished by the potassium nitrate. This combus- tion takes place very rapidly, forming large quantities of carbon dioxid and nitrogen ; hence the explosion. OTHER POTASSIUM SALTS Among other potassium salts that need not be described here may be mentioned potassium chlorid (KC1, see Exp. 198 AN ELEMENTARY EXPERIMENTAL CHEMISTRY 25 c), potassium bromid (KBr, see Exp. 28 c), potassium iodid (KI, see Exp. 30 ), potassium chlorate (KC1O 3 , see Exp. 6 c), potassium sulfid (K 2 S), and potassium cyanid (KCN). SODIUM History and Occurrence. Sodium was first obtained in 1807 by Davy in the same manner as he obtained potassium. (See page 195.) It never occurs free in nature, but in com- bination it is very plentiful, occurring in the sea in sodium chlorid, and on the land chiefly as sodium chlorid, nitrate, carbonate, and sulfate. Manufacture and Uses. It is prepared in a manner analo- gous to that of potassium. In the case of sodium, however, there is no liability to explosions, since sodium does not form a compound with carbon monoxid. Its chief use is in the preparation of the metals manganese and aluminum. It is comparatively cheap. SODIUM CHLORID Occurrence and Extraction. Sodium chlorid, as has been stated before, occurs plentifully in the sea and in various salt beds. It can be obtained easily from sea water ; but by far the greater amount of the salt of commerce is either mined in the solid state, as rock salt, or else is extracted from the earth in the form of brine. The brine is then evaporated. Uses. Besides its uses for seasoning food and preserving meats, it is the basis of the great soda industry. First the sulfate is made from the chlorid, and then the other sodium compounds are obtained from the sulfate. HISTORY, OCCURRENCE AND APPLICATIONS 1 99 SODIUM SULFATE Occurrence. Sodium sulfate occurs in nature as the min- eral thenardite. The sulfate of commerce, however, is made from common salt. Manufacture and Uses. Sodium chlorid and concentrated sulfuric acid are mixed in a large iron pan or retort, and this mixture is heated in a reverberatory furnace.* The mixture is first gently heated, acid sodium sulfate and hydro- chloric acid being formed. Na Cl + H 2 S 4 = Na H S 4 + H Cl The acid gas passes through a flue into a series of towers containing coke, through which water is trickling. The water collects the acid fumes, and is drawn off at the bottom, forming the hydrochloric acid of commerce. The mass left in the pan is then transferred to another part of the furnace, and is subjected to a higher temperature. This forms the normal sodium sulfate (see Defs. under Exp. 1 8), which is the " salt cake " used in the manufacture of sodium carbonate. * A reverberatory furnace is one that has two or more compartments, one in which the fuel is burned, and others in which the substances to be treated are placed. The compartments are so arranged that the flames from the burning fuel are drawn by the draft over and deflected down upon the substances in the other compartments, thus producing the heat to act upon them. Fuel that A, Firebox. B, Hearth upon which substance gives long flames is generally to be heated is placed . used. 2OO AN ELEMENTARY EXPERIMENTAL CHEMISTRY SODIUM CARBONATE There are two main processes for the manufacture of sodium carbonate. The older process is called the Le Blanc, while the more recent is called the Solvay or the Ammonia process. The Le Blanc Process. In the Le Blanc method, sodium sulfate is treated with charcoal (coal) and calcium carbonate. When heated, the sodium sulfate is reduced by the charcoal to sodium sulfid (Na^S), which in turn acts with calcium carbonate forming sodium carbonate. (1) Na 2 S 4 + 4 C = Na 2 S + 4 C (2) Na 2 S + Ca C 3 = Na 2 C 3 + Ca S The mixture is first heated in the coolest part of a reverbera- tory furnace, to produce the first reaction. After a time it is placed in the hottest part of the furnace, when the second reaction takes place. The " black ash " (a mixture of sodium sulfid, N^S, sodium carbonate, NasCOg, coal, and lime) thus obtained is treated with water. The soluble sodium carbonate is then removed from the solution. The Solvay Process. For many years, the Le Blanc was the best process known for the manufacture of sodium carbonate. It has been demonstrated, however, that the Solvay process is more economical. A solution of salt is impregnated with ammonia gas in the proper proportions (a molecule of ammonia for every mole- cule of salt). Then into this there is led carbon dioxid until the solution is saturated. Ammonium chlorid and acid sodium carbonate are formed according to the following reaction : Na Cl + N H 4 H C 3 = N H 4 Cl + Na H C 3 HISTORY, OCCURRENCE AND APPLICATIONS 2OI On account of its slight solubility in the liquid used, the acid sodium carbonate separates out of the solution, is col- lected, and dried. The liquid remaining is treated with lime, and the ammonia thus obtained is used again for im- pregnating a new solution of salt. The acid carbonate may be heated, and the carbon dioxid from this may also be used again, the residue left after heating being the normal car- bonate. Applications. Immense quantities of sodium carbonate are used in the manufacture of glass and soap. It is also used in the preparation of other sodium compounds. The housewife uses the normal carbonate for softening water, and the acid carbonate * for cooking purposes. SODIUM HYDROXID After sodium carbonate has been separated from the " black ash " obtained in the Le Blanc process, there remains in the liquid sodium hydrate that has been formed by the action of the lime present. The solution is heated* air is blo\vn through it, and a quantity of sodium nitrate is added. This oxidizes the sulfid present to sulfate. The solution is then evaporated to dryness, and raised almost to a red heat. The cooled product is commercial sodium hydroxid. Sodium hydroxid is also manufactured by treating a weak solution of sodium carbonate with lime. * Sodium bicarbonate, the acid salt, is made by passing carbon dioxid over sodium carbonate dissolved in its water of crystallization. The principal use of sodium bicarbonate is in the manufacture of bak- ing powder. The constituents of baking powder are sodium bicarbonate and some acid, or acid salt. When mixed with dough, these consti- tuents, by the aid of water in the dough and the heat of the oven, react upon each other. One of the products of the reaction is carbon dioxid ; it is to this that the leavening process is due. 2O2 AN ELEMENTARY EXPERIMENTAL CHEMISTRY Na 2 C 3 + Ca (0 H) 2 = Ca C 3 + 2 Na H The calcium carbonate, being insoluble, is removed by nitration ; and the filtrate, when evaporated, yields the hydroxid. Sodium hydroxid is purified by treating the crude product with alcohol, which dissolves the hydroxid and leaves the impurities. Such sodium hydroxid is called " soda by alcohol." Applications. Sodium hydroxid finds its main uses in the soap factory. It is also used in purifying petroleum, car- bolic acid, in the manufacture of numerous chemicals, in dye works, and in the manufacture of paper. CALCIUM History and Occurrence. Davy, the discoverer of sodium and potassium, first prepared calcium as a powder by electrolysis ; but Matthiessen obtained the first " piece of calcium " in 1856. The metal itself does not occur free in nature, nor is it of any importance. Its various compounds, however, are very plentiful, forming, as dolomite, CaMg(CO 3 ) 2 , whole mountain ranges. It is also found in large quantities as the sulfate and phosphate. The bones of animals and the shells of eggs and of mollusks are mainly composed of cal- cium salts. In fine, it is one of the most plentiful of the elements. CALCIUM CARBONATE Occurrence and Application. This compound occurs as Iceland spar, in transparent crystals. The Carrara marble is a very pure form of calcium carbonate, while ordinary limestone contains various impurities. The artificial sub- stance is made by treating the chlorid with ammonium carbonate. HISTORY, OCCURRENCE AND APPLICATIONS 2O3 The mineral is used for building purposes, and for the manufacture of lime and cements. The beautiful marbles of Italy are used mostly by the sculptor. CALCIUM OXID Manufacture. The natural carbonate is heated in special furnaces, called kilns. Fuel is allowed to burn under the limestone, thus driving off the carbon dioxid and leaving the calcium oxid (see Exp. 21, Part I.). Applications. The chief use of lime is in the making of mortar and cements. Common mortar is a mixture of one part of slaked lime and three or four parts of sand made into a pasty mass. This hardens or " sets up " in a few days, but it takes years for it to harden completely. This is first due to the evapora- tion of the water ; then, as time goes on, the lime takes on carbon dioxid from the air, and becomes calcium carbonate. Hydraulic mortar is made from lime containing more than 10 per cent of silica, and has the ability to harden under water. Portland cement is a hydraulic mortar made from chalk and clay. The two are ground together in water, dried, and burnt in kilns. CALCIUM HYDROXID Calcium hydroxid is simply slaked lime (see Exp. 20 , Part I.). CALCIUM SULFATE Occurrence and Application. Calcium sulfate occurs as gypsum with water of crystallization, and as anhydrite without it. The artificial salt is formed by the action of sulfuric acid on the carbonate. 2O4 AN ELEMENTARY EXPERIMENTAL CHEMISTRY After gypsum has been heated so as to lose its water of crystallization, it is called " plaster of Paris," and possesses the property of hardening when moistened with water. It is therefore extensively used for making plaster casts, and as a cement. Gypsum is found also in some fertilizers. CALCIUM CHLORID Calcium chlorid is obtained as a by-product in many manufacturing processes. On account of its strong attrac- tion for water, it is used in laboratories for the drying of gases. BARIUM AND STRONTIUM History and Occurrence. The elements barium and strontium were isolated by Davy in 1806, in connection with his experiments on calcium. They occur chiefly as barite (BaSO 4 ), witherite (BaCO 3 ), strontianite (SrCO 3 ), and celestite (SrSO 4 ). Compounds. See Exps. 58 and 59, Part I. MAGNESIUM History and Occurrence. Magnesium also was discovered by Davy. It occurs as dolomite, MgCa(CO 3 ) 2 , on the earth ; as the sulfate, MgSO 4 , in mineral springs ; and as chlorid, MgCl 2 , in the sea. The metal is obtained by reducing the chlorid by means of sodium. Compounds. See Exps. 13 and 17 b, Part I. Applications. Magnesium is used in many chemical oper- ations, and as a means of artificial light in photography. The sulfate (Epsom salts) and the carbonate are exten- sively used in medicine. The chlorid is used in the manu- facture of cotton goods. HISTORY, OCCURRENCE AND APPLICATIONS 20$ ZINC History and Occurrence. The alloy brass was known to the ancients, but they did not know that it contained any metal besides copper. The discoverer of the second metal, zinc, is uncertain, although it is mentioned in writings of the 1 6th century. It is found plenteously as the sulfid (ZnS, zinc blende), as the oxid (ZnO, red zinc ore), as the carbonate (ZnCO 3 , calamine), and as the silicate ( (ZnO) 2 SiO 2 , H 2 O, electric calamine). Extraction. The ore is first roasted to convert it into oxid. The oxid is then mixed with carbon. It is then placed in cylindrical fire-clay retorts about three feet long and eight inches in diameter. These retorts are arranged in tiers and set slantingly. To the open end of each retort is joined a conical receiver about 10 inches long, extending downwards. The retorts are heated, and soon burning carbon monoxid appears at the opening of each receiver. The characteristic greenish-blue flame soon appears, showing that the metal is volatilizing. The reduced metal is removed from time to time, the whole operation requiring about eleven hours. The zinc is afterwards redistilled to purify it. Industrial Applications. Zinc is largely used in " galvaniz- ing " iron, which is done by dipping clean iron into melting zinc. It is also used in the manufacture of brass and other alloys. Zinc dust finds an extensive use in organic chem- istry. It is used on a large scale in the manufacture of indigo blue. It is also used as a paint for iron articles. Compounds. See Exps. 12 and 15 , Part I. Uses. The oxid (ZnO) is largely used as a white paint. The chlorid (ZnCL) is a strong antiseptic, and is used also in soldering. The sulfate (ZnSO 4 ) is used in dyeing and calico printing. 2O6 AN ELEMENTARY EXPERIMENTAL CHEMISTRY CADMIUM History and Occurrence. Cadmium was discovered by Stromeyer in 1817. It occurs in zinc ores. The first batch of zinc obtained from the zinc smelter mentioned above con- tains practically all the cadmium ; and, since it boils at a lower temperature than zinc, the cadmium distills over first. Compounds. See Exp. 51, Part I. ALUMINUM History and Occurrence. Aluminum was first obtained by Wohler in 1827, who separated the metal by means of metallic sodium. It is now obtained by an electrolytic method invented by Hall in 1886. Aluminum occurs almost entirely as a silicate. Clay, mica, slate, etc., are all silicates of aluminum. Its oxid (A1 2 O 3 ) is a common ore, and cryolite (AlNa 3 F 6 ) is a very useful mineral. The precious stones, ruby and sapphire, are oxids of aluminum, colored respectively by chromium and cobalt compounds. Extraction. Hall's method of reducing aluminum com- pounds is as follows : A large iron receptacle is lined with carbon. Into this extend a number of large carbon elec- trodes. A mixture consisting of cryolite and fluorite is placed in the receptacle. A strong electric current is then made to pass through the apparatus. The current passes in through the carbon rods, which act as the positive elec- trode, while the carbon lining acts as the negative electrode. After the flux is melted by the current, aluminum oxid (A1 2 O 8 ) is added at intervals. The melted aluminum collects at the negative electrode, while the oxygen unites with HISTORY, OCCURRENCE AND APPLICATIONS 2O/ carbon at the positive. The aluminum, when it has formed in sufficient quantities, is ladled out from the flux. Applications. On account of its lightness, aluminum is used for the manufacture of utensils where great strength is not required. It forms numerous alloys, one of which, alumi- num bronze, is very beautiful. Compounds of Aluminum. See Exp. 55, Part I. Manufacture of Aluminum Sulfate. Finely powdered China clay, aluminum silicate, is roasted and then heated with sulfuric acid. In case other clays are used, the iron is precipitated by means of potassium ferrocyanid. The product is sold on the market under the name of " alum cake." Alum. An alum is a double sulfate of a triad metal and an alkali, and has twelve molecules of water of crys- tallization. Aluminum sulfate is not readily obtained in crystalline form ; hence it is usually crystallized together with potassium or ammonium sulfate, with which it forms crystals of potassium or ammonium aluminum sulfate, KA1 (SO 4 ) 2 ,i2H 2 O, in the form of regular octahedrons. The best commercial method for obtaining alum is to roast, together with coal, the bituminous shale of coal beds. This shale is generally rich in aluminum silicate and iron pyrites. The ferric sulfid gives off sulfur, which becomes oxidized to sulfuric acid. Both aluminum and ferric sul- fates are formed. This mass is dissolved in water, and evaporated down until the ferric sulfate separates out. Then crude potassium chlorid and potassium sulfate are added, and the mixture is agitated until it is cold. The small crystals then formed are washed with cold water, and recrystallized. Uses of Alum. Alum is used mostly in dyeing establish- ments as a mordant. The cloth is placed in an alum solution 2O8 AN ELEMENTARY EXPERIMENTAL CHEMISTRY to which sodium carbonate has been added until the pre- cipitate first formed has been redissolved. After the cloth is dry the color is applied. The mordant causes the color to stick to the fibers of the material. Alum is also used in the manufacture of paper. CHROMIUM History and Occurrence. Vanquelin and Klaproth dis- covered, independently of each other, the element chromium in the mineral crocoisite. It was named chromium (from Xpw/oa, color), because its compounds are all colored. It is found in crocoisite (PbCrO 4 ), and in chromite (Fe O,Cr 2 O 3 ). The green color of emeralds is due to the presence of chromium. It is a hard, gray, almost infusible metal. Compounds. Its principal compounds are chrome alum, CrK(SO 4 ) 2 ,i2H 2 O, used in dyeing and tanning; lead chro- mate, PbCrO 4 , a valuable pigment; potassium chromate, K 2 CrO 4 ; and potassium bi-chromate, K 2 Cr 2 O 7 . The last is the most important commercial product, and is used ex- tensively in bi-chromate batteries and in dyeing. IRON History. It is supposed that iron, the most important and most useful of all the metals, was first extracted from its ores in India. Moses speaks of iron as used by the Hebrews. The Greeks obtained their iron from the Chalybes, a nation living on the shores of the Black Sea. The Romans obtained theirs from Spain and Elba, and from their own mines. Our own iron is mostly obtained from Pennsylvania, although great iron industries have sprung up in Alabama, Ten- nessee, and Illinois. HISTORY, OCCURRENCE AND APPLICATIONS 2OQ Occurrence. Iron is found widely distributed in the earth. It occurs free in certain rocks and in meteorites. The common ores of iron are: magnetite (Fe 3 O 4 ), haematite (Fe 2 O 3 ), siderite (FeCO 3 ), and limonite A, Tuyere. B, Hearth. C, Molten iron. D, Limestone, coke, and ore. E, Sand. Extraction. It would seem to be a very simple process to obtain iron from its ore. Theoretically, all that is necessary with most ores would be to remove the oxygen. This is done in practice by means of carbon, but difficulties come up that have to be conquered before the final product is ob- tained. Theoretically, carbon would be the only thing necessary to use to remove the oxygen, thus : Fe 3 04 + 2 C = 3 Fe + 2 C 2 2IO AN ELEMENTARY EXPERIMENTAL CHEMISTRY It is found that carbon monoxid is also formed thus : 2 Fe 3 4 + 6 C = 6 Fe + 4 C + 2 C 2 But carbon monoxid is also a reducing agent when the tem- perature is high enough, and acts upon the ore, thus : Fe 3 4 + 4 C = 3 Fe + 4 C 2 Iron ore, however, is rarely pure, the impurities being mostly silicates in some form or other. These must be fused or the carbon cannot get a chance to act upon the ore. To effect this, limestone (CaCO 8 ) is added to the carbon, which is used in the form of coke. The limestone, when heated together with the silicates, produces a flux or molten mass. The process is carried on in a furnace from 75 to 100 ft. high. The ore, coke, and limestone are shoveled into the furnace in layers and ignited. A blast of air is driven through the furnace from the bottom by means of a number of pipes called tuyeres. The reduction takes place ; the iron, from its greater specific gravity, sinks to the bottom of the furnace, and is drawn off into molds of sand three or four feet long and three or four inches wide and deep. This is called "pig iron," and contains as impurities sulfur, sili- con, carbon, manganese, and a small amount of phosphorus. Wrought Iron. The pig iron obtained from the blast fur- nace is made into wrought iron by removing from it the carbon, of which there is from 2 to 6 per cent. The process is called " puddling." The pig iron is placed in a reverbera- tory furnace. The sides of the furnace are lined with a sub- stance containing haematite. At first the iron is melted slowly. Finally it boils, and some of the carbon is removed by the oxygen of the haematite. The heat is then in- creased, and the molten mass is stirred by workmen by means of iron rods. When the carbon is burned away, the iron becomes pasty, and is removed in large masses on the HISTORY, OCCURRENCE AND APPLICATIONS 21 I ends of the rods. The iron thus obtained is tough. Most of the phosphorus, sulfur, and silicon of the pig iron remains in the furnace in the slag, and is afterwards removed. Steel. Steel can be made from either pig iron or wrought iron. To make it from wrought iron, the bars are packed in powdered charcoal, and heated for a number of days at a red heat. The bars are then allowed to cool slowly. The best quality of steel is made in this way. It is called " blister steel," since, when taken from the furnace, the bars are covered with blis- ters. The Bessemer process is the ordinary way for making steel. Good pig iron is melted in huge vessels, called converters, hung on trunnions. A blast of air is blown through the molten iron from the bottom. The flame is watched by an expert until just the right moment, and then the blast is stopped. Carbon is then added to the purified iron by means of molten spiegeleisen, a kind of cast iron that contains carbon and manganese. The converter is then emptied, and the steel is cast into ingots. PROPERTIES OF THE VARIOUS KINDS OF IRON Pig iron is brittle and cannot be welded. It is the kind of iron used in making castings. Wrought iron is tough, malleable, ductile, and can be welded. Steel is usually brittle. Its hardness can be changed by tempering; that is, by 212 AN ELEMENTARY EXPERIMENTAL CHEMISTRY heating at various temperatures and suddenly cooling. By allowing hot steel to cool slowly, it can be made almost as soft as wrought iron. Compounds of Iron. See Exps. 14 a, b, 15 / and Exp. 56, Part I. FERROUS SULFATE One source of ferrous sulfate is the bituminous shale from coal beds. The ferrous sulfate is crystallized out from the solution of the roasted product. (See manufacture of alum.) It is used in the manufacture of black ink, in certain photo- graphic processes, and as a reagent in the laboratory. POTASSIUM FERROCYANID AND FERRICYANID When nitrogenous organic matter is fused with potassium hydroxid and iron filings, and the product is treated with water, a solution is obtained which, when purified and evapo- rated, gives yellow crystals of potassium ferrocyanid, K 4 Fe (CN) 6 . This is commonly called yellow prussiate of potash. If treated with ferric chlorid a beautiful blue precipitate, soluble in excess of the chlorid, is formed and is called Prussian blue. If the potassium ferrocyanid is treated with chlorin, the ferricyanid is formed according to the formula : 2 K 4 Fe (C N) 6 + C1 2 = 2 K Cl + 2 K 3 Fe (C N) 6 When evaporated and crystallized, red crystals of the ferri- cyanid are obtained. No precipitate with ferric compounds is formed by the ferricyanid, but with ferrous salts a dark blue precipitate called TurnbulVs blue is formed. MANGANESE History and Occurrence. In 1807, Gahn first isolated manganese. Its most abundant and best known ore is pyro- HISTORY, OCCURRENCE AND APPLICATIONS 213 lusite (MnO 2 ). The metal itself is reddish-white, oxidizes easily, and must be kept away from contact with air. Applications. It finds its main use in its compound man- ganese dioxid (MnO 2 ), which is a \^ery valuable oxidizing agent. Spiegeleisen is an alloy with iron and contains car- bon. As has been stated before, spiegeleisen is used in the manufacture of steel from pig iron. Potassium permanganate is also a valuable oxidizing agent. Manganese sulfate is often used as a mordant. NICKEL AND COBALT History. Nickel was discovered by Cronstedt in 1751, and cobalt probably by Brandt somewhere about 1735. The two metals both occur in combination with arsenic as arsenids. They are almost always associated, and are sepa- rated with difficulty. Of the two metals, nickel is the more important. Extraction of Nickel. The ore, nickel arsenid, is heated in a reverberatory furnace, thus driving off most of the arsenic. The residue is dissolved in hydrochloric acid ; and ftie impurities, with the exception of cobalt, are removed by various processes. From the remaining solution, the cobalt is precipitated as Co(OH) 2 by bleaching powder. When heated, the cobalt hydroxid becomes cobalt oxid, and is removed by filtering. The nickel solution is then treated with calcium hydrate, and is finally reduced, at a white heat, to nickel by means of carbon. It appears on the market in small cubes. Uses. Nickel is valuable as an alloy with steel, and with copper and zinc it forms the well-known alloy German silver. It is also used for plating iron, for coins, and for laboratory utensils. Principal Compounds. See Exp. 57, Part I. 214 AN ELEMENTARY EXPERIMENTAL CHEMISTRY TIN History and Occurrence. Tin was known to the Romans, who obtained it from England after Caesar's conquest. It occurs chiefly as cassiterite, or tin stone, from which the tin of commerce is obtained. The oldest tin mines are those of Cornwall. Tin is also found in considerable quantities in Australia and in the Black Hills of South Dakota. Extraction. The tin ore is first crushed. It is then washed, and the cleansed ore is roasted, to free it from sulfur and arsenic. Mixed with coal and a small quantity of lime, it is then heated. Metallic tin separates out. The metal obtained is somewhat impure; and in order to remove the other metals, it is heated gently. The tin that melts first is ladled out, and this product is further purified by being melted and then stirred with wet wooden sticks. The impurities separate out and the metal is cast into ingots. Uses. The most important application of tin is in the manufacture of tin plate. Its important alloys are bronze (tin and copper), and plumber's solder (tin and lead). It is also rolled into thin sheets known as tin-foil. Compounds. See Exp. 54, Part I. Reducing Action. Stannous chlorid is a very valuable reducing agent in the laboratory. BISMUTH History and Occurrence. Bismuth was spoken of by Basil Valentine in the i5th century. Pott first made a careful study of its properties in 1739. It occurs in nature and often nearly pure. When impure, it is mixed with iron, carbon, and slag, and is then heated in pots. The purer bismuth settles to the bottom. It may HISTORY, OCCURRENCE AND APPLICATIONS 215 be further purified by melting it on an inclined iron plate, when the pure bismuth melts first and runs off. Applications. Bismuth is added to certain alloys, to lower their melting point. It is valuable for making fuses that melt at a moderate elevation of temperature, as in automatic fire extinguishers ; also for safety fuses and lightning ar- resters in connection with the use of the electric current. Compounds. See Exp. 50, Part I. Basic Bismuth Nitrate is extensively used in medicine in the treatment of cholera. It is also used as a cosmetic, and in the manufacture of porcelain. LEAD History and Occurrence. Lead is first mentioned in the Book of Job, and Pliny pointed out the distinction between lead and tin. Lead is rarely found free in nature, but occurs in large quantities as galena (PbS) in England, Spain, and the United States. Reduction from the Ore. The ore galena is mixed wfth lime, and is at first heated gently, a current of air being drawn through the furnace. It is then heated to a higher temperature, and metallic lead is drawn off. The reactions that occur are : Some lead sulfid is changed to lead oxid, 2 Pb S + 3 2 = 2 Pb +- 2 S 2 and some to lead sulfate. Pb S + 2 2 = Pb S 0* Lead oxid and lead sulfate each react upon lead sulfid, forming lead and sulfur dioxid. Pb S + 2 Pb ^ 3 Pb + S 2 and Pb S + Pb S 4 = 2 Pb + 2 S 2 2l6 AN ELEMENTARY EXPERIMENTAL CHEMISTRY There is probably always some silver in galena. This, if sufficient in amount, is separated by a process described later (see page 219). Applications. Lead is used principally for making water pipes, roofing, shot, and various alloys, such as type-metal, solder, and pewter. In making shot, some arsenic is added to the lead. Lead compounds are used in making white lead, and in glass making. Compounds. Besides the compounds studied in Exp. 53, Part I., there are a few others that should be noted. These are red lead (Pb 3 O 4 ) ; white lead, which is a basic lead car- bonate ; and lead acetate, commonly called " sugar of lead." Red Lead is made by heating litharge (PbO) on trays in a reverberatory furnace. White Lead is made by placing rolls of perforated sheet lead in earthenware pots partly filled with diluted vinegar (acetic acid). The lead is not in contact with the vinegar. The pots are covered with tan bark or manure. In time the acid fumes change the lead to lead acetate, and the carbon dioxid from the tan bark or manure changes this to the carbonate. Lead Acetate is made by dissolving litharge in acetic acid. COPPER History and Occurrence. Copper was probably the first metal used by mankind, having been used by pre-historic man in making his weapons. It is very abundant, which is fortunate, since it is one of the most useful metals we have. It occurs native in the Lake Superior regions, the mines there furnishing almost chemically pure copper. In Mon- tana, Idaho, and Arizona, vast quantities of copper occur, associated with the precious metals. It is also found as the oxid in Russia and Australia, and as the sulfid in England. HISTORY, OCCURRENCE AND APPLICATIONS 21 j Extraction. When the oxid ore is heated with carbon, the metal separates out. The sulfid ore, however, must first be roasted in order that part of the sulfid may be changed to oxid. Then the mixture of sulfid and oxid is heated further in a reverberatory furnace, where the following re- action occurs. Cu S + 2 Cu = 3 Cu + S 2 . Industrial Applications. Since the development of elec- tricity, copper has become one of the staple products, on account of its use as an electric conductor. Its use in the preparation of the alloys, brass and bronze, has already been mentioned. As sheathing for ships, etc., large quantities are used. Copper coins are familiar to us all. Compounds. See Exps. 4; 9 g ; 39 ; Part I. Copper sulfate is an important commercial product The sulfid is gently roasted, and the copper sulfate thus formed is dissolved out with water. It is also formed as a by-pro- duct in the purifying of gold and silver. It is used in the preparation of copper arsenite (Paris green), in the manu- facture of other pigments, and also in calico printing. MERCURY History. Mercury is mentioned by the later Greek writers under the name v&pdyvpos (from vS 5 p/fl msolubl CJ C O rt _~ E .^. i KCN JH" , 1 Na 2 HPO 4 * . rt a) ^ *>. 1 g T3 c (NH 4 ) 2 S S c 8 cu-gS ^1" A n3 -0 c' H 2 S S JD K __ ; .. X iT H 2 S0 4 ^ iS'l ^ JS c rt 5, rt O o" HC1 ^ a J "S a; CD o rt 1 feb t/) . -S ;55 ^ C , 1C _s U (NH 4 ),CO 3 -sis ^| 2 ^ 3 '"D ^||1 0) 1 o 1 rt 1 3 .2 !-S. _s "o 'o .84, 221. and aqua regia, 84. chlorid, 84. Graphic formulae, 132. Graphite, 12, 186. Gun cotton, 180. Gunpowder, 197. Gypsum, 203. H Haemoglobin, 58. Halogens, 163. Heat unit, 1 13. specific, 113. of chemical action, 147. of neutralization, 149. of solution and" of hydration, 1 50. Helium, 181. Hydraulic main, 190. Hydriodic acid, 44, 169. Hydrobromic acid, 42, 167. Hydrochloric acid, 35, 36, 37, 38, 164. Hydrocyanic acid, 192. Hydrofluoric acid, 45, 169. Hydrogen, 160. antimonid, 75. arsenid, 73. burning of, in chlorin, 35. from iron and sulfuric acid, 25. from potassium and water, 23. from sodium and water, 22. from zinc and sulphuric acid, 23. sulfid, 47, 48, 171. action of bromin upon, 49. union with chlorin by light, 35. weight of liter of, 106. Hydrosulfuric acid (see hydrogen sulfid). I Ice, 3. Illuminating gas, 189. Ink, 212. Indestructibility of matter, 90. lodin, 43, 1 68. and phosphorus, 43. extraction from seaweed, 168. replacement by bromin, 45. replacement by chlorin, 45. Ions, 153. Iron, 20. chlorids, 80. galvanized, 205. oxid, 20, 21. pig, 210, 211. specific heat of, 116. sulfate, 212. sulfid, 48. wrought, 210. Isomorphic equivalents, 119. Kerosene, 188. Kindling temperature, 13. Lamp black, 187. Laughing gas, 71."* Law, definition of, 90. of Boyle, 99, 100. of Charles, 101. of definite proportion by w r eight, 92. of definite proportions by vol- ume, 1 20. of Dulong and Petit, 117, 118. INDEX 249 Law of Gay Lussac, 120. of indestructibility of matter, 90. Law of multiple proportions, 93. Lead, 77, 215. acetate, 216. action of acids upon, 77. chlorid, 78. oxids, 77. replacement of, by zinc, 78. specific heat of, 114. sulfate, 78. sulfid, 78. white, 216. Le Blanc process, 200. Lime, 31, 203. Limewater, 31. Limewater and carbonic acid, 31. Litharge, 216. Magnesium, 20, 204. oxid, 20. oxid and water, 20. oxid and sulfuric acid, 28. Manganese, 212. test for, 40. chlorid, 41. dioxid, 7. sulfate, 213. Marble, analysis of, 32. Marsh gas, 52, 53, 188. Matches, 183. Matter, three conditions of, 3. indestructibility of, 90. Mechanical mixture, 4. Mendele Jeff's table, 126. Mercurous chlorid, 77. Mercury, 6, 77, 217. and oxygen, 6, 7. Microcrith, 118. Micro-organisms, 58. Molecular weights, 120. formulae, determination of, 129. Molecules, 97. Monad, 132. Mordant, 207, 208. Mortar, 203. Muriatic acid (see hydrochloric acid). Naming of compounds, 137. Naphtha, 188. Nascent state of elements, 65. Natural gas, 188. Nature of flame, 54. Negative, electro, 137 photographs, 220. Neutralization, 28. Nickel, Si, 213. action of acids upon, 81. electrical equivalent of, 112. Niter plantations, 197. Nitric acid, 59, 62, 179. instability of, 62. neutralization of, 62. Nitric oxid, 63. a reducing agent, 71. peroxid, an oxidizing agent, 71- Nitrous acid, 178. Nitrous oxid, 70. analysis of, 72. Nitrogen, 57, 177. acids of, 178. oxids of, 178. Nitroglycerin, 180. Normal salt, 29. 250 INDEX Olefiant gas, 53, 54. Organic chemistry, 154. Organic matter, burning of, 154. Oxalic acid, ana-ysis of, 50. Oxidation, 27. Oxygen, 7, 159. atomic weight of, 122. burning of, in hydrogen, 26. density of, 122. molecular weight of, 123. number of atoms to the mole- cule, 121. preparation from manganese dioxid, 7. from potassium chlorate, 8. by electrolysis, 8. Oxy-hydrogen blowpipe, 161. Ozone, 10. number of atoms in molecule, 122. P Paraffin, 188. Paris green, 217. Peat, 187. Percentage composition, calcula- tion of, 144. Periodicity of the elements, 125. Petroleum, 188. Pewter, 216. Phosphin, 183. Phosphoric acid, II, 183. Phosphorous acid, n. Phosphorus, 10, 181. and iodin, 43, 44. and oxygen, 10, 11. Photography, 220. Physical change, 3. Physiological chemistry, 59. Pig iron, 210, 211. Plants, chemistry of, 58. Plaster of Paris, 204. Platinum, 85, 222. action of acids on, 85. action of hot metals on, 85. Porcelain, 194. Positive, electro, 137. Positive photographs, 220. Potassium, 18, 195. bromid, 43. and sulfuric acid, 43, carbonate, 196. chlorate, 8, 198. molecular weight of, 123. chlorid, 39. molecular weight of, 124. hydro xid, 19, 196. ferrocyanid, 212. iodid, 44. and sulfuric acid, 44. and phosphoric acid, 44. oxid, 1 8. Pottery, 195. Precipitates, tables of, 228, 229, 230, 239. Products, 90. Prout's hypothesis, 127. Prussian blue, 212. Prussic acid, 192. Puddling, 210. Putreficadon, 58. Qualitative analysis, 227. Quantitative analysis, 94, 227. Quartz, 193. Quick lime, 31, 203. ^ Quicksilver (see mercury). INDEX 251 Radical, 69. Reactions, writing of, 138. Reduction, 27. Reverberatory furnace, 199. S Safety lamp, 55. Salt, definition of, 29. acid, basic, and normal, 29. cake, 199. Saponification. 156. Seaweed, 168. Selenium, 176. Shot, 216. Silicon, 193. dioxid, 193. fluorid, 46. Silver, 83, 219. action of acids upon, 83. bromid, 83. chlorid, 83. reduction of, by hydrogen, 84. analysis of, 94. iodid, 83. oxid, 83. replacement of, by copper, 84. Soap, 156, 197. Sodium, 17, 198. amalgam, 18. carbonate, 200. chlorid, 36, 198. analysis of, 94. hydro xid, 17, 201. and sulfuric acid, 28. oxid, 17. sulfate, 28, 29, 199. Solder, 205, 214, 216. Solvay process, 200. Specific heat, 113. Spiegeleisen, 211, 213. Steel, 211. Stibin, 75. Stoic biometry, 141. Strontium, 82, 204. chlorid, 82. flame test, 82. hydroxid, 82. nitrate, 82. oxid, 82. Specific heat, 113. apparatus for, 114. Subtimation, 70. Sugar of lead, 216. Sulfates, 28, 29, 138. Sulfids, 47, 138. Sulfites, 138. Sulfur, 13, 170. allotropic forms of, 13. dioxid, 14, 15, 1 6, 172. trioxid, 14. Sulfuric acid, 16, 173. Sulfurous acid, 14. Supporter of combustion, 13. Synthesis, 6. Tellurium, 176. Temperature, effect of changes upon gases, 101, 102. absolute zero of, 103. Ternary, 137. Theory, 90. Thermal equivalent, 117. Thermochemistry, 144, 145. Tin, 79, 214. action of acids upon, 79. chlorid, 214. INDEX Tin, replacement by zinc, 79. oxid, 79. Toning pictures, 221. Triad, 132. TurnbulPs blue, 212. Tuyeres, 210. Type metal, 216. Typhoid, 162. U Unstable equilibrium, 71. Valence, 131. Vapor density, 107. Ventilation, 32. W Water, 161. synthesis of, from burning hy- drogen, 25. from oxidation of organic sub- stances, 26. gas, 191. Water, hard and soft, 56. of crystallization, 24. permanently hard, 56. synthesis of, 25. temporarily hard, 56. Weights, atomic, 108, 109, 124, 128. Weldon's process, 164. White lead, 216. Writing reactions, 138. Wrought iron, 210. Yeast, 155. Zero, absolute, 103. Zinc, 19, 205. chemical equivalence of, no. dust, 19. oxid, 19, 205. oxid and sulfuric acid, 27. sulfid, 48. sulfate, 205. UNIVERSITY OF CALIFORNIA LIB LAST DATE THIS BOOK IS STAMPED BELOW NOV261917 ,v 8 1918 f I 1919 ft 1923 DEC I 1931 30m-6,'14 6862 235384