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A more convenient form of eudiometer. A form of eudiometer (Fig. 21) different from that shown on page 43 is sometimes used to avoid the calculations necessary in reducing the volumes of the gases to the same conditions of temperature and pressure in order to make comparisons. With this apparatus it is possible to take the readings of the volumes under the same conditions of temperature and pressure, and thus compare them directly. The apparatus (Fig. 21) is filled with mercury and the gases introduced into the tube A. The experiment is carried out as in the preceding one, except that before taking the reading of the gas volumes, mercury is either added to the tube B or withdrawn from it by means of the stopcock C, until it stands at exactly the same height in both tubes. The gas inclosed in tube A is then under atmospheric pressure; and since but a few minutes are required for performing the experiment, the conditions of temperature and pressure may be regarded as constant. Hence the volumes of the hydrogen and oxygen and of the residual gas may be read off from the tube and directly compared.
Method used by Berzelius and Dumas. The method used by these investigators enables us to determine directly the proportion by weight in which the hydrogen and oxygen combine. Fig. 22 illustrates the apparatus used in making this determination. B is a glass tube containing copper oxide. C and D are glass tubes filled with calcium chloride, a substance which has great affinity for water. The tubes B and C, including their contents, are carefully weighed, and the apparatus connected as shown in the figure. A slow current of pure hydrogen is then passed through A, and that part of the tube B which contains copper oxide is carefully heated. The hydrogen combines with the oxygen present in the copper oxide to form water, which is absorbed by the calcium chloride in tube C. The calcium chloride in tube D prevents any moisture entering tube C from the air. The operation is continued until an appreciable amount of water has been formed. The tubes B and C are then weighed once more. The loss of weight in the tube B will exactly equal the weight of oxygen taken up from the copper oxide in the formation of the water. The gain in weight in the tube C will exactly equal the weight of the water formed. The difference in these weights will of course equal the weight of the hydrogen present in the water formed.
Dumas' results. The above method for the determination of the composition of water was first used by Berzelius in 1820. The work was repeated in 1843 by Dumas, the average of whose results is as follows:
Weight of water formed 236.36 g. Oxygen given up by the copper oxide 210.04 ——— Weight of hydrogen present in water 26.32
According to this experiment the ratio of hydrogen to oxygen in water is therefore 26.32 to 210.04, or as l to 7.98
Morley's results. The American chemist Morley has recently determined the composition of water, extreme precautions being taken to use pure materials and to eliminate all sources of error. The hydrogen and oxygen which combined, as well as the water formed, were all accurately weighed. According to Morley's results, 1 part of hydrogen by weight combines with 7.94 parts of oxygen to form water.
Comparison of results obtained. From the above discussions it is easy to see that it is by experiment alone that the composition of a compound can be determined. Different methods may lead to slightly different results. The more accurate the method chosen and the greater the skill with which the experiment is carried out, the more accurate will be the results. It is generally conceded by chemists that the results obtained by Morley in reference to the composition of water are the most accurate ones. In accordance with these results, then, water must be regarded as a compound containing hydrogen and oxygen in the proportion of 1 part by weight of hydrogen to 7.94 parts by weight of oxygen.
Relation between the volume of aqueous vapor and the volumes of the hydrogen and oxygen which combine to form it. When the quantitative synthesis of water is carried out in the eudiometer as described above, the water vapor formed by the union of the hydrogen and oxygen at once condenses. The volume of the resulting liquid is so small that it may be disregarded in making the calculations. If, however, the experiment is carried out at a temperature of 100 deg. or above, the water-vapor formed is not condensed and it thus becomes possible to compare the volume of the vapor with the volumes of hydrogen and oxygen which combined to form it. This can be accomplished by surrounding the arm A of the eudiometer (Fig. 23) with the tube B through which is passed the vapor obtained by boiling some liquid which has a boiling point above 100 deg.. In this way it has been proved that 2 volumes of hydrogen and 1 volume of oxygen combine to form exactly 2 volumes of water vapor, the volumes all being measured under the same conditions of temperature and pressure. It will be noted that the relation between these volumes may be expressed by whole numbers. The significance of this very important fact will be discussed in a subsequent chapter.
Occurrence of water. Water not only covers about three fourths of the surface of the earth, and is present in the atmosphere in the form of moisture, but it is also a common constituent of the soil and rocks and of almost every form of animal and vegetable organism. The human body is nearly 70% water. This is derived not only from the water which we drink but also from the food which we eat, most of which contains a large percentage of water. Thus potatoes contain about 78% of water, milk 85%, beef over 50%, apples 84%, tomatoes 94%.
Impurities in water. Chemically pure water contains only hydrogen and oxygen. Such a water never occurs in nature, however, for being a good solvent, it takes up certain substances from the rocks and soil with which it comes in contact. When such waters are evaporated these substances are deposited in the form of a residue. Even rain water, which is the purest form occurring in nature, contains dust particles and gases dissolved from the atmosphere. The foreign matter in water is of two kinds, namely, mineral, such as common salt and limestone, and organic, that is the products of animal and vegetable life.
Mineral matter in water. The amount and nature of the mineral matter present in different waters vary greatly, depending on the character of the rocks and soil with which the waters come in contact. The more common of the substances present are common salt and compounds of calcium, magnesium, and iron. One liter of the average river water contains about 175 mg. of mineral matter. Water from deep wells naturally contains more mineral matter than river water, generally two or three times as much, while sea water contains as much as 35,000 mg. to the liter.
Effect of impurities on health. The mineral matter in water does not, save in very exceptional cases, render the water injurious to the human system. In fact the presence of a certain amount of such matter is advantageous, supplying the mineral constituents necessary for the formation of the solid tissues of the body. The presence of organic matter, on the other hand, must always be regarded with suspicion. This organic matter may consist not only of the products of animal and vegetable life but also of certain microscopic forms of living organisms which are likely to accompany such products. Contagious diseases are known to be due to the presence in the body of minute living organisms or germs. Each disease is caused by its own particular kind of germ. Through sewage these germs may find their way from persons afflicted with disease into the water supply, and it is principally through the drinking water that certain of these diseases, especially typhoid fever, are spread. It becomes of great importance, therefore, to be able to detect such matter when present in drinking water as well as to devise methods whereby it can be removed or at least rendered harmless.
Analysis of water. The mineral analysis of a water is, as the name suggests, simply the determination of the mineral matter present. Sanitary analysis, on the other hand, is the determination of the organic matter present. The physical properties of a water give no conclusive evidence as to its purity, since a water may be unfit for drinking purposes and yet be perfectly clear and odorless. Neither can any reliance be placed on the simple methods often given for testing the purity of water. Only the trained chemist can carry out such methods of analysis as can be relied upon.
Purification of water. Three general methods are used for the purification of water, namely, distillation, filtration, and boiling.
1. Distillation. The most effective way of purifying natural waters is by the process of distillation. This consists in boiling the water and condensing the steam. Fig. 24 illustrates the process of distillation, as commonly conducted in the laboratory. Ordinary water is poured into the flask A and boiled. The steam is conducted through the condenser B, which consists essentially of a narrow glass tube sealed within a larger one, the space between the two being filled with cold water, which is admitted at C and escapes at D. The inner tube is thus kept cool and the steam in passing through it is condensed. The water formed by the condensation of the steam collects in the receiver E and is known as distilled water. Such water is practically pure, since the impurities are nonvolatile and remain in the flask A.
Commercial distillation. In preparing distilled water on a large scale, the steam is generated in a boiler or other metal container and condensed by passing it through a pipe made of metal, generally tin. This pipe is wound into a spiral and is surrounded by a current of cold water. Distilled water is used by the chemist in almost all of his work. It is also used in the manufacture of artificial ice and for drinking water.
Fractional distillation. In preparing distilled water, it is evident that if the natural water contains some substance which is volatile its vapor will pass over and be condensed with the steam, so that the distillate will not be pure water. Even such mixtures, however, may generally be separated by repeated distillation. Thus, if a mixture of water (boiling point 100 deg.) and alcohol (boiling point 78 deg.) is distilled, the alcohol, having the lower boiling point, tends to distill first, followed by the water. The separation of the two is not perfect, however, but may be made nearly so by repeated distillations. The process of separating a mixture of volatile substances by distillation is known as fractional distillation.
2. Filtration. The process of distillation practically removes all nonvolatile foreign matter, mineral as well as organic. In purifying water for drinking purposes, however, it is only necessary to eliminate the latter or to render it harmless. This is ordinarily done either by filtration or boiling. In filtration the water is passed through some medium which will retain the organic matter. Ordinary charcoal is a porous substance and will condense within its pores the organic matter in water if brought in contact with it. It is therefore well adapted to the construction of filters. Such filters to be effective must be kept clean, since it is evident that the charcoal is useless after its pores are filled. A more effective type of filter is the Chamberlain-Pasteur filter. In this the water is forced through a porous cylindrical cup, the pores being so minute as to strain out the organic matter.
City filtration beds. For purifying the water supply of cities, large filtration beds are prepared from sand and gravel, and the water is allowed to filter through these. Some of the impurities are strained out by the filter, while others are decomposed by the action of certain kinds of bacteria present in the sand. Fig. 25 shows a cross section of a portion of the filter used in purifying the water supply of Philadelphia. The water filters through the sand and gravel and passes into the porous pipe A, from which it is pumped into the city mains. The filters are covered to prevent the water from freezing in cold weather.
3. Boiling. A simpler and equally efficient method for purifying water for drinking purposes consists in boiling the water. It is the germs in water that render it dangerous to health. These germs are living forms of matter. If the water is boiled, the germs are killed and the water rendered safe. While these germs are destroyed by heat, cold has little effect upon them. Thus Dewar, in working with liquid hydrogen, exposed some of these minute forms of life to the temperature of boiling hydrogen (-252 deg.) without killing them.
Self-purification of water. It has long been known that water contaminated with organic matter tends to purify itself when exposed to the air. This is due to the fact that the water takes up a small amount of oxygen from the air, which gradually oxidizes the organic matter present in the water. While water is undoubtedly purified in this way, the method cannot be relied upon to purify a contaminated water so as to render it safe for drinking purposes.
Physical properties. Pure water is an odorless and tasteless liquid, colorless in thin layers, but having a bluish tinge when observed through a considerable thickness. It solidifies at 0 deg. and boils at 100 deg. under the normal pressure of one atmosphere. If the pressure is increased, the boiling point is raised. When water is cooled it steadily contracts until the temperature of 4 deg. is reached: it then expands. Water is remarkable for its ability to dissolve other substances, and is the best solvent known. Solutions of solids in water are more frequently employed in chemical work than are the solid substances, for chemical action between substances goes on more readily when they are in solution than it does when they are in the solid state.
Chemical properties. Water is a very stable substance, or, in other words, it does not undergo decomposition readily. To decompose it into its elements by heat alone requires a very high temperature; at 2500 deg., for example, only about 5% of the entire amount is decomposed. Though very stable towards heat, water can be decomposed in other ways, as by the action of the electrical current or by certain metals.
Heat of formation and heat of decomposition are equal. The fact that a very high temperature is necessary to decompose water into hydrogen and oxygen is in accord with the fact that a great deal of heat is evolved by the union of hydrogen and oxygen; for it has been proved that the heat necessary to decompose a compound into its elements (heat of decomposition) is equal to the heat evolved in the formation of a compound from its elements (heat of formation).
Water of crystallization. When a solid is dissolved in water and the resulting solution is allowed to evaporate, the solid separates out, often in the form of crystals. It has been found that the crystals of many compounds, although perfectly dry, give up a definite amount of water when heated, the substance at the same time losing its crystalline form. Such water is called water of crystallization. This varies in amount with different compounds, but is perfectly definite for the same compound. Thus, if a perfectly dry crystal of copper sulphate is strongly heated in a tube, water is evolved and condenses on the sides of the tube, the crystal crumbling to a light powder. The weight of the water evolved is always equal to exactly 36.07% of the weight of copper sulphate crystals heated. The water must therefore be in chemical combination with the substance composing the crystal; for if simply mixed with it or adhering to it, not only would the substance appear moist but the amount present would undoubtedly vary. The combination, however, must be a very weak one, since the water is often expelled by even a gentle heat. Indeed, in some cases the water is given up on simple exposure to air. Such compounds are said to be efflorescent. Thus a crystal of sodium sulphate (Glauber's salt) on exposure to air crumbles to a fine powder, owing to the escape of its water of crystallization. Other substances have just the opposite property: they absorb moisture when exposed to the air. For example, if a bit of dry calcium chloride is placed in moist air, in the course of a few hours it will have absorbed sufficient moisture to dissolve it. Such substances are said to be deliquescent. A deliquescent body serves as a good drying or desiccating agent. We have already employed calcium chloride as an agent for absorbing the moisture from hydrogen. Many substances, as for example quartz, form crystals which contain no water of crystallization.
Mechanically inclosed water. Water of crystallization must be carefully distinguished from water which is mechanically inclosed in a crystal and which can be removed by powdering the crystal and drying. Thus, when crystals of common salt are heated, the water inclosed in the crystal is changed into steam and bursts the crystal with a crackling sound. Such crystals are said to decrepitate. That this water is not combined is proved by the fact that the amount present varies and that it has all the properties of water.
Uses of water. The importance of water in its relation to life and commerce is too well known to require comment. Its importance to the chemist has also been pointed out. It remains to call attention to the fact that it is used as a standard in many physical measurements. Thus 0 deg. and 100 deg. on the centigrade scale are respectively the freezing and the boiling points of water under normal pressure. The weight of 1 cc. of water at its point of greatest density is the unit of weight in the metric system, namely, the gram. It is also taken as the unit for the determination of the density of liquids and solids as well as for the measurement of amounts of heat.
HYDROGEN DIOXIDE
Composition. As has been shown, 1 part by weight of hydrogen combines with 7.94 parts by weight of oxygen to form water. It is possible, however, to obtain a second compound of hydrogen and oxygen differing from water in composition in that 1 part by weight of hydrogen is combined with 2 x 7.94, or 15.88 parts, of oxygen. This compound is called hydrogen dioxide or hydrogen peroxide, the prefixes di- and per- signifying that it contains more oxygen than hydrogen oxide, which is the chemical name for water.
Preparation. Hydrogen dioxide cannot be prepared cheaply by the direct union of hydrogen and oxygen, and indirect methods must therefore be used. It is commonly prepared by the action of a solution of sulphuric acid on barium dioxide. The change which takes place may be indicated as follows:
sulphuric acid + barium dioxide = barium sulphate + hydrogen dioxide ——————— ——————— ———————- ———————— hydrogen barium barium hydrogen sulphur oxygen sulphur oxygen oxygen oxygen
In other words, the barium and hydrogen in the two compounds exchange places. By this method a dilute solution of the dioxide in water is obtained. It is possible to separate the dioxide from the water by fractional distillation. This is attended with great difficulties, however, since the pure dioxide is explosive. The distillation is carried on under diminished pressure so as to lower the boiling points as much as possible; otherwise the high temperature would decompose the dioxide.
Properties. Pure hydrogen dioxide is a colorless sirupy liquid having a density of 1.49. Its most characteristic property is the ease with which it decomposes into water and oxygen. One part by weight of hydrogen is capable of holding firmly only 7.94 parts of oxygen. The additional 7.94 parts of oxygen present in hydrogen dioxide are therefore easily evolved, the compound breaking down into water and oxygen. This decomposition is attended by the generation of considerable heat. In dilute solution hydrogen dioxide is fairly stable, although such a solution should be kept in a dark, cool place, since both heat and light aid in the decomposition of the dioxide.
Uses. Solutions of hydrogen dioxide are used largely as oxidizing agents. The solution sold by druggists contains 3% of the dioxide and is used in medicine as an antiseptic. Its use as an antiseptic depends upon its oxidizing properties.
EXERCISES
1. Why does the chemist use distilled water in making solutions, rather than filtered water?
2. How could you determine the total amount of solid matter dissolved in a sample of water?
3. How could you determine whether a given sample of water is distilled water?
4. How could the presence of air dissolved in water be detected?
5. How could the amount of water in a food such as bread or potato be determined?
6. Would ice frozen from impure water necessarily be free from disease germs?
7. Suppose that the maximum density of water were at 0 deg. in place of 4 deg.; what effect would this have on the formation of ice on bodies of water?
8. Is it possible for a substance to contain both mechanically inclosed water and water of crystallization?
9. If steam is heated to 2000 deg. and again cooled, has any chemical change taken place in the steam?
10. Why is cold water passed into C instead of D (Fig. 24)?
11. Mention at least two advantages that a metal condenser has over a glass condenser.
12. Draw a diagram of the apparatus used in your laboratory for supplying distilled water.
13. 20 cc. of hydrogen and 7 cc. of oxygen are placed in a eudiometer and the mixture exploded. (a) How many cubic centimeters of aqueous vapor are formed? (b) What gas and how much of it remains in excess?
14. (a) What weight of water can be formed by the combustion of 100 L of hydrogen, measured under standard conditions? (b)What volume of oxygen would be required in (a)? (c)What weight of potassium chlorate is necessary to prepare this amount of oxygen?
15. What weight of oxygen is present in 1 kg. of the ordinary hydrogen dioxide solution? In the decomposition of this weight of the dioxide into water and oxygen, what volume of oxygen (measured under standard conditions) is evolved?
CHAPTER V
THE ATOMIC THEORY
Three fundamental laws of matter. Before we can gain any very definite idea in regard to the structure of matter, and the way in which different kinds of substances act chemically upon each other, it is necessary to have clearly in view three fundamental laws of matter. These laws have been established by experiment, and any conception which may be formed concerning matter must therefore be in harmony with them. The laws are as follows:
Law of conservation of matter. This law has already been touched upon in the introductory chapter, and needs no further discussion. It will be recalled that it may be stated thus: Matter can neither be created nor destroyed, though it can be changed from one form into another.
Law of definite composition. In the earlier days of chemistry there was much discussion as to whether the composition of a given compound is always precisely the same or whether it is subject to some variation. Two Frenchmen, Berthollet and Proust, were the leaders in this discussion, and a great deal of most useful experimenting was done to decide the question. Their experiments, as well as all succeeding ones, have shown that the composition of a pure chemical compound is always exactly the same. Water obtained by melting pure ice, condensing steam, burning hydrogen in oxygen, has always 11.18% hydrogen and 88.82% oxygen in it. Red oxide of mercury, from whatever source it is obtained, contains 92.6% mercury and 7.4% oxygen. This truth is known as the law of definite composition, and may be stated thus: The composition of a chemical compound never varies.
Law of multiple proportion. It has already been noted, however, that hydrogen and oxygen combine in two different ratios to form water and hydrogen dioxide respectively. It will be observed that this fact does not contradict the law of definite composition, for entirely different substances are formed. These compounds differ from each other in composition, but the composition of each one is always constant. This ability of two elements to unite in more than one ratio is very frequently observed. Carbon and oxygen combine in two different ratios; nitrogen and oxygen combine to form as many as five distinct compounds, each with its own precise composition.
In the first decade of the last century John Dalton, an English school-teacher and philosopher, endeavored to find some rule which holds between the ratios in which two given substances combine. His studies brought to light a very simple relation, which the following examples will make clear. In water the hydrogen and oxygen are combined in the ratio of 1 part by weight of hydrogen to 7.94 parts by weight of oxygen. In hydrogen dioxide the 1 part by weight of hydrogen is combined with 15.88 parts by weight of oxygen. The ratio between the amounts of oxygen which combine with the same amount of hydrogen to form water and hydrogen dioxide respectively is therefore 7.94: 15.88, or 1: 2.
Similarly, the element iron combines with oxygen to form two oxides, one of which is black and the other red. By analysis it has been shown that the former contains 1 part by weight of iron combined with 0.286 parts by weight of oxygen, while the latter contains 1 part by weight of iron combined with 0.429 parts by weight of oxygen. Here again we find that the amounts of oxygen which combine with the same fixed amount of iron to form the two compounds are in the ratio of small whole numbers, viz., 2:3.
Many other examples of this simple relation might be given, since it has been found to hold true in all cases where more than one compound is, formed from the same elements. Dalton's law of multiple proportion states these facts as follows: When any two elements, A and B, combine to form more than one compound, the amounts of B which unite with any fixed amount of A bear the ratio of small whole numbers to each other.
Hypothesis necessary to explain the laws of matter. These three generalizations are called laws, because they express in concise language truths which are found by careful experiment to hold good in all cases. They do not offer any explanation of the facts, but merely state them. The human mind, however, does not rest content with the mere bare facts, but seeks ever to learn the explanation of the facts. A suggestion which is offered to explain such a set of facts is called an hypothesis. The suggestion which Dalton offered to explain the three laws of matter, called the atomic hypothesis, was prompted by his view of the constitution of matter, and it involves three distinct assumptions in regard to the nature of matter and chemical action. Dalton could not prove these assumptions to be true, but he saw that if they were true the laws of matter become very easy to understand.
Dalton's atomic hypothesis. The three assumptions which Dalton made in regard to the nature of matter, and which together constitute the atomic hypothesis, are these:
1. All elements are made up of minute, independent particles which Dalton designated as atoms.
2. All atoms of the same element have equal masses; those of different elements have different masses; in any change to which an atom is subjected its mass does not change.
3. When two or more elements unite to form a compound, the action consists in the union of a definite small number of atoms of each element to form a small particle of the compound. The smallest particles of a given compound are therefore exactly alike in the number and kinds of atoms which they contain, and larger masses of the substances are simply aggregations of these least particles.
Molecules and atoms. Dalton applied the name atom not only to the minute particles of the elements but also to the least particles of compounds. Later Avogadro, an Italian scientist, pointed out the fact that the two are different, since the smallest particle of an element is a unit, while that of a compound must have at least two units in it. He suggested the name molecule for the least particle of a compound which can exist, retaining the name atom for the smallest particle of an element. In accordance with this distinction, we may define the atom and the molecule as follows: An atom is the smallest particle of an element which can exist. A molecule is the smallest particle of a compound which can exist. It will be shown in a subsequent chapter that sometimes two or more atoms of the same element unite with each other to form molecules of the element. While the term atom, therefore, is applicable only to elements, the term molecule is applicable both to elements and compounds.
The atomic hypothesis and the laws of matter. Supposing the atomic hypothesis to be true, let us now see if it is in harmony with the laws of matter.
1. The atomic hypothesis and the law of conservation of matter. It is evident that if the atoms never change their masses in any change which they undergo, the total quantity of matter can never change and the law of conservation of matter must follow.
2. The atomic hypothesis and the law of definite composition. According to the third supposition, when iron combines with sulphur the union is between definite numbers of the two kinds of atoms. In the simplest case one atom of the one element combines with one atom of the other. If the sulphur and the iron atoms never change their respective masses when they unite to form a molecule of iron sulphide, all iron sulphide molecules will have equal amounts of iron in them and also of sulphur. Consequently any mass made up of iron sulphide molecules will have the same fraction of iron by weight as do the individual iron sulphide molecules. Iron sulphide, from whatever source, will have the same composition, which is in accordance with the law of definite composition.
3. The atomic hypothesis and the law of multiple proportion. But this simplest case may not always be the only one. Under other conditions one atom of iron might combine with two of sulphur to form a molecule of a second compound. In such a case the one atom of iron would be in combination with twice the mass of sulphur that is in the first compound, since the sulphur atoms all have equal masses. What is true for one molecule will be true for any number of them; consequently when such quantities of these two compounds are selected as are found to contain the same amount of iron, the one will contain twice as much sulphur as the other.
The combination between the atoms may of course take place in other simple ratios. For example, two atoms of one element might combine with three or with five of the other. In all such cases it is clear that the law of multiple proportion must hold true. For on selecting such numbers of the two kinds of molecules as have the same number of the one kind of atoms, the numbers of the other kind of atoms will stand in some simple ratio to each other, and their weights will therefore stand in the same simple ratio.
Testing the hypothesis. Efforts have been made to find compounds which do not conform to these laws, but all such attempts have resulted in failure. If such compounds should be found, the laws would be no longer true, and the hypothesis of Dalton would cease to possess value. When an hypothesis has been tested in every way in which experiment can test it, and is still found to be in harmony with the facts in the case, it is termed a theory. We now speak of the atomic theory rather than of the atomic hypothesis.
Value of a theory. The value of a theory is twofold. It aids in the clear understanding of the laws of nature because it gives an intelligent idea as to why these laws should be in operation.
A theory also leads to discoveries. It usually happens that in testing a theory much valuable work is done, and many new facts are discovered. Almost any theory in explaining given laws will involve a number of consequences apart from the laws it seeks to explain. Experiment will soon show whether these facts are as the theory predicts they will be. Thus Dalton's atomic theory predicted many properties of gases which experiment has since verified.
Atomic weights. It would be of great advantage in the study of chemistry if we could determine the weights of the different kinds of atoms. It is evident that this cannot be done directly. They are so small that they cannot be seen even with a most powerful microscope. It is calculated that it would take 200,000,000 hydrogen atoms placed side by side to make a row one centimeter long. No balance can weigh such minute objects. It is possible, however, to determine their relative weights,—that is, how much heavier one is than another. These relative weights of the atoms are spoken of as the atomic weights of the elements.
If elements were able to combine in only one way,—one atom of one with one atom of another,—the problem of determining the atomic weights would be very simple. We should merely have to take some one convenient element as a standard, and find by experiment how much of each other element would combine with a fixed weight of it. The ratios thus found would be the same ratios as those between the atoms of the elements, and thus we should have their relative atomic weights. The law of multiple proportion calls attention to the fact that the atoms combine in other ratios than 1: 1, and there is no direct way of telling which one, if any, of the several compounds in a given case is the one consisting of a single atom of each element.
If some way were to be found of telling how much heavier the entire molecule of a compound is than the atom chosen as a standard,—that is, of determining the molecular weights of compounds,—the problem could be solved, though its solution would not be an entirely simple matter. There are ways of determining the molecular weights of compounds, and there are other experiments which throw light directly upon the relative weights of the atoms. These methods cannot be described until the facts upon which they rest have been studied. It will be sufficient for the present to assume that these methods are trustworthy.
Standard for atomic weights. Since the atomic weights are merely relative to some one element chosen as a standard, it is evident that any one of the elements may serve as this standard and that any convenient value may be assigned to its atom. At one time oxygen was taken as this standard, with the value 100, and the atomic weights of the other elements were expressed in terms of this standard. It would seem more rational to take the element of smallest atomic weight as the standard and give it unit value; accordingly hydrogen was taken as the standard with an atomic weight of 1. Very recently, however, this unit has been replaced by oxygen, with an atomic weight of 16.
Why oxygen is chosen as the standard for atomic weights. In the determination of the atomic weight of an element it is necessary to find the weight of the element which combines with a definite weight of another element, preferably the element chosen as the standard. Since oxygen combines with the elements far more readily than does hydrogen to form definite compounds, it is far better adapted for the standard element, and has accordingly replaced hydrogen as the standard. Any definite value might be given to the weight of the oxygen atom. In assigning a value to it, however, it is convenient to choose a whole number, and as small a number as possible without making the atomic weight of any other element less than unity. For these reasons the number 16 has been chosen as the atomic weight of oxygen. This makes the atomic weight of hydrogen equal to 1.008, so that there is but little difference between taking oxygen as 16 and hydrogen as 1 for the unit.
The atomic weights of the elements are given in the Appendix.
EXERCISES
1. Two compounds were found to have the following compositions: (a) oxygen = 69.53%, nitrogen = 30.47%; (b) oxygen = 53.27%, nitrogen = 46.73%. Show that the law of multiple proportion holds in this case.
2. Two compounds were found to have the following compositions: (a) oxygen = 43.64%, phosphorus = 56.36%; (b) oxygen = 56.35%, phosphorus = 43.65%. Show that the law of multiple proportion holds in this case.
3. Why did Dalton assume that all the atoms of a given element have the same weight?
CHAPTER VI
CHEMICAL EQUATIONS AND CALCULATIONS
Formulas. Since the molecule of any chemical compound consists of a definite number of atoms, and this number never changes without destroying the identity of the compound, it is very convenient to represent the composition of a compound by indicating the composition of its molecules. This can be done very easily by using the symbols of the atoms to indicate the number and the kind of the atoms which constitute the molecule. HgO will in this way represent mercuric oxide, a molecule of which has been found to contain 1 atom each of mercury and oxygen. H_{2}O will represent water, the molecules of which consist of 1 atom of oxygen and 2 of hydrogen, the subscript figure indicating the number of the atoms of the element whose symbol precedes it. H_{2}SO_{4} will stand for sulphuric acid, the molecules of which contain 2 atoms of hydrogen, 1 of sulphur, and 4 of oxygen. The combination of symbols which represents the molecule of a substance is called its _formula_.
Equations. When a given substance undergoes a chemical change it is possible to represent this change by the use of such symbols and formulas. In a former chapter it was shown that mercuric oxide decomposes when heated to form mercury and oxygen. This may be expressed very briefly in the form of the equation
(1) HgO = Hg + O.
When water is electrolyzed two new substances, hydrogen and oxygen, are formed from it. This statement in the form of an equation is
(2) H_{2}O = 2H + O.
The coefficient before the symbol for hydrogen indicates that a single molecule of water yields two atoms of hydrogen on decomposition.
In like manner the combination of sulphur with iron is expressed by the equation
(3) Fe + S = FeS.
The decomposition of potassium chlorate by heat takes place as represented by the equation
(4) KClO_{3} = KCl + 3O.
Reading of equations. Since equations are simply a kind of shorthand way of indicating chemical changes which occur under certain conditions, in reading an equation the full statement for which it stands should be given. Equation (1) should be read, "Mercuric oxide when heated gives mercury and oxygen"; equation (2) is equivalent to the statement, "When electrolyzed, water produces hydrogen and oxygen"; equation (3), "When heated together iron and sulphur unite to form iron sulphide"; equation (4), "Potassium chlorate when heated yields potassium chloride and oxygen."
Knowledge required for writing equations. In order to write such equations correctly, a considerable amount of exact knowledge is required. Thus, in equation (1) the fact that red oxide of mercury has the composition represented by the formula HgO, that it is decomposed by heat, that in this decomposition mercury and oxygen are formed and no other products,—all these facts must be ascertained by exact experiment before the equation can be written. An equation expressing these facts will then have much value.
Having obtained an equation describing the conduct of mercuric oxide on being heated, it will not do to assume that other oxides will behave in like manner. Iron oxide (FeO) resembles mercuric oxide in many respects, but it undergoes no change at all when heated. Manganese dioxide, the black substance used in the preparation of oxygen, has the formula MnO_{2}. When this substance is heated oxygen is set free, but the metal manganese is not liberated; instead, a different oxide of manganese containing less oxygen is produced. The equation representing the reaction is
3MnO_{2} = Mn_{3}O_{4} + 2O.
Classes of reactions. When a chemical change takes place in a substance the substance is said to undergo a reaction. Although a great many different reactions will be met in the study of chemistry, they may all be grouped under the following heads.
1. Addition. This is the simplest kind of chemical action. It consists in the union of two or more substances to produce a new substance. The combination of iron with sulphur is an example:
Fe + S = FeS.
2. Decomposition. This is the reverse of addition, the substance undergoing reaction being parted into its constituents. The decomposition of mercuric oxide is an example: HgO = Hg + O.
3. _Substitution._ It is sometimes possible for an element in the free state to act upon a compound in such a way that it takes the place of one of the elements of the compound, liberating it in turn. In the study of the element hydrogen it was pointed out that hydrogen is most conveniently prepared by the action of sulphuric or hydrochloric acid upon zinc. When sulphuric acid is used a substance called zinc sulphate, having the composition represented by the formula ZnSO_{4}, is formed together with hydrogen. The equation is
Zn + H_{2}SO_{4} = ZnSO_{4} + 2H.
When hydrochloric acid is used zinc chloride and hydrogen are the products of reaction:
Zn + 2HCl = ZnCl_{2} + 2H.
When iron is used in place of zinc the equation is
Fe + H_{2}SO_{4} = FeSO_{4} + 2H.
These reactions are quite similar, as is apparent from an examination of the equations. In each case 1 atom of the metal replaces 2 atoms of hydrogen in the acid, and the hydrogen escapes as a gas. When an element in the free state, such as the zinc in the equations just given, takes the place of some one element in a compound, setting it free from chemical combination, the act is called substitution.
Other reactions illustrating substitution are the action of sodium on water,
Na + H_{2}O = NaOH + H;
and the action of heated iron upon water,
3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H.
4. Double decomposition. When barium dioxide (BaO{2}) is treated with sulphuric acid two compounds are formed, namely, hydrogen dioxide (H{2}O{2}) and barium sulphate (BaSO{4}). The equation is
BaO{2} + H{2}SO{4} = BaSO{4} + H{2}O{2}.
In this reaction it will be seen that the two elements barium and hydrogen simply exchange places. Such a reaction is called a double decomposition. We shall meet with many examples of this kind of chemical reactions.
Chemical equations are quantitative. The use of symbols and formulas in expressing chemical changes has another great advantage. Thus, according to the equation
H_{2}O = 2H + O,
1 molecule of water is decomposed into 2 atoms of hydrogen and 1 atom of oxygen. But, as we have seen, the relative weights of the atoms are known, that of hydrogen being 1.008, while that of oxygen is 16. The molecule of water, being composed of 2 atoms of hydrogen and 1 atom of oxygen, must therefore weigh relatively 2.016 + 16, or 18.016. The amount of hydrogen in this molecule must be 2.016/18.016, or 11.18% of the whole, while the amount of oxygen must be 16/18.018, or 88.82% of the whole. Now, since any definite quantity of water is simply the sum of a great many molecules of water, it is plain that the fractions representing the relative amounts of hydrogen and oxygen present in a molecule must likewise express the relative amounts of hydrogen and oxygen present in any quantity of water. Thus, for example, in 20 g. of water there are 2.016/18.016 x 20, or 2.238 g. of hydrogen, and 16/18.016 x 20, or 17.762 g. of oxygen. These results in reference to the composition of water of course agree exactly with the facts obtained by the experiments described in the chapter on water, for it is because of those experiments that the values 1.008 and 16 are given to hydrogen and oxygen respectively.
It is often easier to make calculations of this kind in the form of a proportion rather than by fractions. Since the molecule of water and the two atoms of hydrogen which it contains have the ratio by weight of 18.016: 2.016, any mass of water has the same ratio between its total weight and the weight of the hydrogen in it. Hence, to find the number of grams (x) of hydrogen in 20 g. of water, we have the proportion
18.016 : 2.016 :: 20 g. : x (grams of hydrogen).
Solving for x, we get 2.238 for the number of grams of hydrogen. Similarly, to find the amount (x) of oxygen present in the 20 g. of water, we have the proportion
18.016 : 16 :: 20 : x
from which we find that x = 17.762 g.
Again, suppose we wish to find what weight of oxygen can be obtained from 15 g. of mercuric oxide. The equation representing the decomposition of mercuric oxide is
HgO = Hg + O.
The relative weights of the mercury and oxygen atoms are respectively 200 and 16. The relative weight of the mercuric oxide molecule must therefore be the sum of these, or 216. The molecule of mercuric oxide and the atom of oxygen which it contains have the ratio 216: 16. This same ratio must therefore hold between the weight of any given quantity of mercuric oxide and that of the oxygen which it contains. Hence, to find the weight of oxygen in 15 g. of mercuric oxide, we have the proportion
216 : 16 :: 15 : x (grams of oxygen).
On the other hand, suppose we wish to prepare, say, 20 g. of oxygen. The problem is to find out what weight of mercuric oxide will yield 20 g. of oxygen. The following proportion evidently holds
216 : 16 :: x (grams of mercuric oxide) : 20;
from which we get x = 270.
In the preparation of hydrogen by the action of sulphuric acid upon zinc, according to the equation,
Zn + H_{2}SO_{4} = ZnSO_{4} + 2 H,
suppose that 50 g. of zinc are available; let it be required to calculate the weight of hydrogen which can be obtained. It will be seen that 1 atom of zinc will liberate 2 atoms of hydrogen. The ratio by weight of a zinc to an hydrogen atom is 65.4: 1.008; of 1 zinc atom to 2 hydrogen atoms, 65.4: 2.016. Zinc and hydrogen will be related in this reaction in this same ratio, however many atoms of zinc are concerned. Consequently in the proportion
65.4 : 2.016 :: 50 : x,
x will be the weight of hydrogen set free by 50 g. of zinc. The weight of zinc sulphate produced at the same time can be found from the proportion
65.4 : 161.46 :: 50 : x;
where 161.46 is the molecular weight of the zinc sulphate, and x the weight of zinc sulphate formed. In like manner, the weight of sulphuric acid used up can be calculated from the proportion
65.4 : 98.076 :: 50 : x.
These simple calculations are possible because the symbols and formulas in the equations represent the relative weights of the substances concerned in a chemical reaction. When once the relative weights of the atoms have been determined, and it has been agreed to allow the symbols to stand for these relative weights, an equation or formula making use of the symbols becomes a statement of a definite numerical fact, and calculations can be based on it.
Chemical equations not algebraic. Although chemical equations are quantitative, it must be clearly understood that they are not algebraic. A glance at the equations
7 + 4 = 11, 8 + 5 = 9 + 4
will show at once that they are true. The equations
HgO = Hg + O, FeO = Fe + O
are equally true in an algebraic sense, but experiment shows that only the first is true chemically, for iron oxide (FeO) cannot be directly decomposed into iron and oxygen. Only such equations as have been found by careful experiment to express a real chemical transformation, true both for the kinds of substances as well as for the weights, have any value.
Chemical formulas and equations, therefore, are a concise way of representing qualitatively and quantitatively facts which have been found by experiment to be true in reference to the composition of substances and the changes which they undergo.
Formulas representing water of crystallization. An examination of substances containing water of crystallization has shown that in every case the water is present in such proportion by weight as can readily be represented by a formula. For example, copper sulphate (CuSO{4}) and water combine in the ratio of 1 molecule of the sulphate to 5 of water; calcium sulphate (CaSO{4}) and water combine in the ratio 1: 2 to form gypsum. These facts are expressed by writing the formulas for the two substances with a period between them. Thus the formula for crystallized copper sulphate is CuSO{4}.5H{2}O; that of gypsum is CaSO{4}.2H{2}O.
Heat of reaction. Attention has frequently been directed to the fact that chemical changes are usually accompanied by heat changes. In general it has been found that in every chemical action heat is either absorbed or given off. By adopting a suitable unit for the measurement of heat, the heat change during a chemical reaction can be expressed in the equation for the reaction.
Heat cannot be measured by the use of a thermometer alone, since the thermometer measures the intensity of heat, not its quantity. The easiest way to measure a quantity of heat is to note how warm it will make a definite amount of a given substance chosen as a standard. Water has been chosen as the standard, and the unit of heat is called a calorie. A calorie is defined as the amount of heat required to raise the temperature of one gram of water one degree.
By means of this unit it is easy to indicate the heat changes in a given chemical reaction. The equation
2H + O = H_{2}O + 68,300 cal.
means that when 2.016 g. of hydrogen combine with 16 g. of oxygen, 18.016 g. of water are formed and 68,300 cal. are set free.
C + 2S = CS_{2} - 19,000 cal.
means that an expenditure of 19,000 cal. is required to cause 12 g. of carbon to unite with 64.12 g. of sulphur to form 76.12 g. of carbon disulphide. In these equations it will be noted that the symbols stand for as many grams of the substance as there are units in the weights of the atoms represented by the symbols. This is always understood to be the case in equations where the heat of reaction is given.
Conditions of a chemical action are not indicated by equations. Equations do not tell the conditions under which a reaction will take place. The equation
HgO = Hg + O
does not tell us that it is necessary to keep the mercuric oxide at a high temperature in order that the decomposition may go on. The equation
Zn + 2HCl = ZnCl_{2} + 2H
in no way indicates the fact that the hydrochloric acid must be dissolved in water before it will act upon the zinc. From the equation
H + Cl = HCl
it would not be suspected that the two gases hydrogen and chlorine will unite instantly in the sunlight, but will stand mixed in the dark a long time without change. It will therefore be necessary to pay much attention to the details of the conditions under which a given reaction occurs, as well as to the expression of the reaction in the form of an equation.
EXERCISES
1. Calculate the percentage composition of the following substances: (a) mercuric oxide; (b) potassium chlorate; (c) hydrochloric acid; (d) sulphuric acid. Compare the results obtained with the compositions as given in Chapters II and III.
2. Determine the percentage of copper, sulphur, oxygen, and water in copper sulphate crystals. What weight of water can be obtained from 150 g. of this substance?
3. What weight of zinc can be dissolved in 10 g. of sulphuric acid? How much zinc sulphate will be formed?
4. How many liters of hydrogen measured under standard conditions can be obtained from the action of 8 g. of iron on 10 g. of sulphuric acid? How much iron sulphate (FeSO_{4}) will be formed?
5. 10 g. of zinc were used in the preparation of hydrogen; what weight of iron will be required to prepare an equal volume?
6. How many grams of barium dioxide will be required to prepare 1 kg. of common hydrogen dioxide solution? What weight of barium sulphate will be formed at the same time?
7. What weight of the compound Mn{3}O{4} will be formed by strongly heating 25 g. of manganese dioxide? What volume of oxygen will be given off at the same time, measured under standard conditions?
8. (a) What is the weight of 100 l. of hydrogen measured in a laboratory in which the temperature is 20 deg. and pressure 750 mm.? (b) What weight of sulphuric acid is necessary to prepare this amount of hydrogen? (c) The density of sulphuric acid is 1.84. Express the acid required in (b) in cubic centimeters.
9. What weight of potassium chlorate is necessary to furnish sufficient oxygen to fill four 200 cc. bottles in your laboratory (the gas to be collected over water)?
CHAPTER VII
NITROGEN AND THE RARE ELEMENTS: ARGON, HELIUM, NEON, KRYPTON, XENON
Historical. Nitrogen was discovered by the English chemist Rutherford in 1772. A little later Scheele showed it to be a constituent of air, and Lavoisier gave it the name azote, signifying that it would not support life. The name nitrogen was afterwards given it because of its presence in saltpeter or niter. The term azote and symbol Az are still retained by the French chemists.
Occurrence. Air is composed principally of oxygen and nitrogen in the free state, about 78 parts by volume out of every 100 parts being nitrogen. Nitrogen also occurs in nature in the form of potassium nitrate (KNO{3})—commonly called saltpeter or niter—as well as in sodium nitrate (NaNO{3}). Nitrogen is also an essential constituent of all living organisms; for example, the human body contains about 2.4% of nitrogen.
Preparation from air. Nitrogen can be prepared from air by the action of some substance which will combine with the oxygen, leaving the nitrogen free. Such a substance must be chosen, however, as will combine with the oxygen to form a product which is not a gas, and which can be readily separated from the nitrogen. The substances most commonly used for this purpose are phosphorus and copper.
1. By the action of phosphorus. The method used for the preparation of nitrogen by the action of phosphorus is as follows:
The phosphorus is placed in a little porcelain dish, supported on a cork and floated on water (Fig. 26). It is then ignited by contact with a hot wire, and immediately a bell jar or bottle is brought over it so as to confine a portion of the air. The phosphorus combines with the oxygen to form an oxide of phosphorus, known as phosphorus pentoxide. This is a white solid which floats about in the bell jar, but in a short time it is all absorbed by the water, leaving the nitrogen. The withdrawal of the oxygen is indicated by the rising of the water in the bell jar.
2. By the action of copper. The oxygen present in the air may also be removed by passing air slowly through a heated tube containing copper. The copper combines with the oxygen to form copper oxide, which is a solid. The nitrogen passes on and may be collected over water.
Nitrogen obtained from air is not pure. Inasmuch as air, in addition to oxygen and nitrogen, contains small amounts of other gases, and since the phosphorus as well as the copper removes only the oxygen, it is evident that the nitrogen obtained by these methods is never quite pure. About 1% of the product is composed of other gases, from which it is very difficult to separate the nitrogen. The impure nitrogen so obtained may, however, be used for a study of most of the properties of nitrogen, since these are not materially affected by the presence of the other gases.
Preparation from compounds of nitrogen. Pure nitrogen may be obtained from certain compounds of the element. Thus, if heat is applied to the compound ammonium nitrite (NH{4}NO{2}), the change represented in the following equation takes place:
NH_{4}NO_{2} = 2H_{2}O + 2N.
Physical properties. Nitrogen is similar to oxygen and hydrogen in that it is a colorless, odorless, and tasteless gas. One liter of nitrogen weighs 1.2501 g. It is almost insoluble in water. It can be obtained in the form of a colorless liquid having a boiling point of -195 deg. at ordinary pressure. At -214 deg. it solidifies.
Chemical properties. Nitrogen is characterized by its inertness. It is neither combustible nor a supporter of combustion. At ordinary temperatures it will not combine directly with any of the elements except under rare conditions. At higher temperatures it combines with magnesium, lithium, titanium, and a number of other elements. The compounds formed are called nitrides, just as compounds of an element with oxygen are called oxides. When it is mixed with oxygen and subjected to the action of electric sparks, the two gases slowly combine forming oxides of nitrogen. A mixture of nitrogen and hydrogen when treated similarly forms ammonia, a gaseous compound of nitrogen and hydrogen. Since we are constantly inhaling nitrogen, it is evident that it is not poisonous. Nevertheless life would be impossible in an atmosphere of pure nitrogen on account of the exclusion of the necessary oxygen.
Argon, helium, neon, krypton, xenon. These are all rare elements occurring in the air in very small quantities. Argon, discovered in 1894, was the first one obtained. Lord Rayleigh, an English scientist, while engaged in determining the exact weights of various gases, observed that the nitrogen obtained from the air is slightly heavier than pure nitrogen obtained from its compounds. After repeating his experiments many times, always with the same results, Rayleigh finally concluded that the nitrogen which he had obtained from the air was not pure, but was mixed with a small amount of some unknown gas, the density of which is greater than that of nitrogen. Acting on this assumption, Rayleigh, together with the English chemist Ramsay, attempted to separate the nitrogen from the unknown gas. Knowing that nitrogen would combine with magnesium, they passed the nitrogen obtained from the air and freed from all known substances through tubes containing magnesium heated to the necessary temperature. After repeating this operation, they finally succeeded in obtaining from the atmospheric nitrogen a small volume of gas which would not combine with magnesium and hence could not be nitrogen. This proved to be a new element, to which they gave the name argon. As predicted, this new element was found to be heavier than nitrogen, its density as compared with hydrogen as a standard being approximately 20, that of nitrogen being only 14. About 1% of the atmospheric nitrogen proved to be argon. The new element is characterized by having no affinity for other elements. Even under the most favorable conditions it has not been made to combine with any other element. On this account it was given the name argon, signifying lazy or idle. Like nitrogen, it is colorless, odorless, and tasteless. It has been liquefied and solidified. Its boiling point is -187 deg..
Helium was first found in the gases expelled from certain minerals by heating. Through the agency of the spectroscope it had been known to exist in the sun long before its presence on the earth had been demonstrated,—a fact suggested by the name helium, signifying the sun. Its existence in traces in the atmosphere has also been proven. It was first liquefied by Onnes in July, 1908. Its boiling point, namely -269 deg., is the lowest temperature yet reached.
The remaining elements of this group—neon, krypton, and xenon—have been obtained from liquid air. When liquid air is allowed to boil, the constituents which are the most difficult to liquefy, and which therefore have the lowest boiling points, vaporize first, followed by the others in the order of their boiling points. It is possible in this way to make at least a partial separation of the air into its constituents, and Ramsay thus succeeded in obtaining from liquid air not only the known constituents, including argon and helium, but also the new elements, neon, krypton, and xenon. These elements, as well as helium, all proved to be similar to argon in that they are without chemical activity, apparently forming no compounds whatever. The percentages present in the air are very small. The names, neon, krypton, xenon, signify respectively, new, hidden, stranger.
EXERCISES
1. How could you distinguish between oxygen, hydrogen, and nitrogen?
2. Calculate the relative weights of nitrogen and oxygen; of nitrogen and hydrogen.
3. In the preparation of nitrogen from the air, how would hydrogen do as a substance for the removal of the oxygen?
4. What weight of nitrogen can be obtained from 10 l. of air measured under the conditions of temperature and pressure which prevail in your laboratory?
5. How many grams of ammonium nitrite are necessary in the preparation of 20 l. of nitrogen measured over water under the conditions of temperature and pressure which prevail in your laboratory?
6. If 10 l. of air, measured under standard conditions, is passed over 100 g. of hot copper, how much will the copper gain in weight?
CHAPTER VIII
THE ATMOSPHERE
Atmosphere and air. The term atmosphere is applied to the gaseous envelope surrounding the earth. The term air is generally applied to a limited portion of this envelope, although the two words are often used interchangeably. Many references have already been made to the composition and properties of the atmosphere. These statements must now be collected and discussed somewhat more in detail.
Air formerly regarded as an element. Like water, air was at first regarded as elementary in character. Near the close of the eighteenth century Scheele, Priestley, and Lavoisier showed by their experiments that it is a mixture of at least two gases,—those which we now call oxygen and nitrogen. By burning substances in an inclosed volume of air and noting the contraction in volume due to the removal of the oxygen, they were able to determine with some accuracy the relative volumes of oxygen and nitrogen present in the air.
The constituents of the atmosphere. The constituents of the atmosphere may be divided into two general groups: those which are essential to life and those which are not essential.
1. Constituents essential to life. In addition to oxygen and nitrogen at least two other substances, namely, carbon dioxide and water vapor, must be present in the atmosphere in order that life may exist. The former of these is a gaseous compound of carbon and oxygen having the formula CO{2}. Its properties will be discussed in detail in the chapter on the compounds of carbon. Its presence in the air may be shown by causing the air to bubble through a solution of calcium hydroxide (Ca(OH){2}), commonly called lime water. The carbon dioxide combines with the calcium hydroxide in accordance with the following equation:
Ca(OH){2} + CO{2} = CaCO{3} + H{2}O.
The resulting calcium carbonate (CaCO_{3}) is insoluble in water and separates in the form of a white powder, which causes the solution to appear milky.
The presence of water vapor is readily shown by its condensation on cold objects as well as by the fact that a bit of calcium chloride when exposed to the air becomes moist, and may even dissolve in the water absorbed from the air.
2. Constituents not essential to life. In addition to the essential constituents, the air contains small percentages of various other gases, the presence of which so far as is known is not essential to life. This list includes the rare elements, argon, helium, neon, krypton, and xenon; also hydrogen, ammonia, hydrogen dioxide, and probably ozone. Certain minute forms of life (germs) are also present, the decay of organic matter being due to their presence.
Function of each of the essential constituents. (1) The oxygen directly supports life through respiration. (2) The nitrogen, on account of its inactivity, serves to dilute the oxygen, and while contrary to the older views, it is possible that life might continue to exist in the absence of the atmospheric nitrogen, yet the conditions of life would be entirely changed. Moreover, nitrogen is an essential constituent of all animal and plant life. It was formerly supposed that neither animals nor plants could assimilate the free nitrogen, but it has been shown recently that the plants of at least one natural order, the Leguminosae, to which belong the beans, peas, and clover, have the power of directly assimilating the free nitrogen from the atmosphere. This is accomplished through the agency of groups of bacteria, which form colonies in little tubercles on the roots of the plants. These bacteria probably assist in the absorption of nitrogen by changing the free nitrogen into compounds which can be assimilated by the plant. Fig. 27 shows the tubercles on the roots of a variety of bean. (3) The presence of water vapor in the air is necessary to prevent excessive evaporation from both plants and animals. (4) Carbon dioxide is an essential plant food.
The quantitative analysis of air. A number of different methods have been devised for the determination of the percentages of the constituents present in the atmosphere. Among these are the following.
1. Determination of oxygen. (1) The oxygen is withdrawn from a measured volume of air inclosed in a tube, by means of phosphorus.
To make the determination, a graduated tube is filled with water and inverted in a vessel of water. Air is introduced into the tube until it is partially filled with the gas. The volume of the inclosed air is carefully noted and reduced to standard conditions. A small piece of phosphorus is attached to a wire and brought within the tube as shown in Fig. 28. After a few hours the oxygen in the inclosed air will have combined with the phosphorus, the water rising to take its place. The phosphorus is removed and the volume is again noted and reduced to standard conditions. The contraction in the volume of the air is equal to the volume of oxygen absorbed.
(2) The oxygen may also be estimated by passing a measured volume of air through a tube containing copper heated to a high temperature. The oxygen in the air combines with the copper to form copper oxide (CuO). Hence the increase in the weight of the copper equals the weight of the oxygen in the volume of air taken.
(3) A more accurate method is the following. A eudiometer tube is filled with mercury and inverted in a vessel of the same liquid. A convenient amount of air is then introduced into the tube and its volume accurately noted. There is then introduced more than sufficient hydrogen to combine with the oxygen present in the inclosed air, and the volume is again accurately noted. The mixture is then exploded by an electric spark, and the volume is once more taken. By subtracting this volume from the total volume of the air and hydrogen there is obtained the contraction in volume due to the union of the oxygen and hydrogen. The volume occupied by the water formed by the union of the two gases is so small that it may be disregarded in the calculation. Since oxygen and hydrogen combine in the ratio 1: 2 by volume, it is evident that the contraction in volume due to the combination is equal to the volume occupied by the oxygen in the air contained in the tube, plus twice this volume of hydrogen. In other words, one third of the total contraction is equal to the volume occupied by the oxygen in the inclosed air. The following example will make this clear:
Volume of air in tube 50.0 cc. Volume after introducing hydrogen 80.0 Volume after combination of oxygen and hydrogen 48.5 Contraction in volume due to combination (80 cc.-48.5 cc.) 31.5 Volume of oxygen in 50 cc. of air (1/3 of 31.5) 10.5
All these methods agree in showing that 100 volumes of dry air contain approximately 21 volumes of oxygen.
2. Determination of nitrogen. If the gas left after the removal of oxygen from a portion of air is passed over heated magnesium, the nitrogen is withdrawn, argon and the other rare elements being left. It may thus be shown that of the 79 volumes of gas left after the removal of the oxygen from 100 volumes of air, approximately 78 are nitrogen and 0.93 argon. The other elements are present in such small quantities that they may be neglected.
3. Determination of carbon dioxide. The percentage of carbon dioxide in any given volume of air may be determined by passing the air over calcium hydroxide or some other compound which will combine with the carbon dioxide. The increase in the weight of the hydroxide equals the weight of the carbon dioxide absorbed. The amount present in the open normal air is from 3 to 4 parts by volume in 10,000 volumes of air, or about 0.04%.
4. Determination of water vapor. The water vapor present in a given volume of air may be determined by passing the air over calcium chloride (or some other compound which has a strong affinity for water), and noting the increase in the weight of the chloride. The amount present varies not only with the locality, but there is a wide variation from day to day in the same locality because of the winds and changes in temperature.
Processes affecting the composition of the air. The most important of these processes are the following.
1. Respiration. In the process of respiration some of the oxygen in the inhaled air is absorbed by the blood and carried to all parts of the body, where it combines with the carbon of the worn-out tissues. The products of oxidation are carried back to the lungs and exhaled in the form of carbon dioxide. The amount exhaled by an adult averages about 20 l. per hour. Hence in a poorly ventilated room occupied by a number of people the amount of carbon dioxide rapidly increases. While this gas is not poisonous unless present in large amounts, nevertheless air containing more than 15 parts in 10,000 is not fit for respiration.
2. Combustion. All of the ordinary forms of fuel contain large percentages of carbon. On burning, this carbon combines with oxygen in the air, forming carbon dioxide. Combustion and respiration, therefore, tend to diminish the amount of oxygen in the air and to increase the amount of carbon dioxide.
3. Action of plants. Plants have the power, when in the sunlight, of absorbing carbon dioxide from the air, retaining the carbon and returning at least a portion of the oxygen to the air. It will be observed that these changes are just the opposite of those brought about by the processes of respiration and combustion.
Poisonous effect of exhaled air. The differences in the percentages of oxygen, carbon dioxide, and moisture present in inhaled air and exhaled air are shown in the following analyses.
INHALED AIR EXHALED AIR Oxygen 21.00% 16.00% Carbon dioxide 0.04 4.38 Moisture variable saturated
The foul odor of respired air is due to the presence of a certain amount of organic matter. It is possible that this organic matter rather than the carbon dioxide is responsible for the injurious effects which follow the respiration of impure air. The extent of such organic impurities present may be judged, however, by the amount of carbon dioxide present, since the two are exhaled together.
The cycle of carbon in nature. Under the influence of sunlight, the carbon dioxide absorbed from the air by plants reacts with water and small amounts of other substances absorbed from the soil to form complex compounds of carbon which constitute the essential part of the plant tissue. This reaction is attended by the evolution of oxygen, which is restored to the air. The compounds resulting from these changes are much richer in their energy content than are the substances from which they are formed; hence a certain amount of energy must have been absorbed in their formation. The source of this energy is the sun's rays.
If the plant is burned, the changes which took place in the formation of the compounds present are largely reversed. The carbon and hydrogen present combine with oxygen taken from the air to form carbon dioxide and water, while the energy absorbed from the sun's rays is liberated in the form of energy of heat. If, on the other hand, the plant is used as food, the compounds present are used in building up the tissues of the body. When this tissue breaks down, the changes which it undergoes are very similar to those which take place when the plant is burned. The carbon and hydrogen combine with the inhaled oxygen to form carbon dioxide and water, which are exhaled. The energy possessed by the complex substances is liberated partly in the form of energy of heat, which maintains the heat of the body, and partly in the various forms of muscular energy. The carbon originally absorbed from the air by the plant in the form of carbon dioxide is thus restored to the air and is ready to repeat the cycle of changes.
The composition of the air is constant. Notwithstanding the changes constantly taking place which tend to alter the composition of the air, the results of a great many analyses of air collected in the open fields show that the percentages of oxygen and nitrogen as well as of carbon dioxide are very nearly constant. Indeed, so constant are the percentages of oxygen and nitrogen that the question has arisen, whether these two elements are not combined in the air, forming a definite chemical compound. That the two are not combined but are simply mixed together can be shown in a number of ways, among which are the following.
1. When air dissolves in water it has been found that the ratio of oxygen to nitrogen in the dissolved air is no longer 21: 78, but more nearly 35: 65. If it were a chemical compound, the ratio of oxygen to nitrogen would not be changed by solution in water.
2. A chemical compound in the form of a liquid has a definite boiling point. Water, for example, boils at 100 deg.. Moreover the steam which is thus formed has the same composition as the water. The boiling point of liquid air, on the other hand, gradually rises as the liquid boils, the nitrogen escaping first followed by the oxygen. If the two were combined, they would pass off together in the ratio in which they are found in the air.
Why the air has a constant composition. If air is a mixture and changes are constantly taking place which tend to modify its composition, how, then, do we account for the constancy of composition which the analyses reveal? This is explained by several facts. (1) The changes which are caused by the processes of combustion and respiration, on the one hand, and the action of plants, on the other, tend to equalize each other. (2) The winds keep the air in constant motion and so prevent local changes. (3) The volume of the air is so vast and the changes which occur are so small compared with the total amount of air that they cannot be readily detected. (4) Finally it must be noted that only air collected in the open fields shows this constancy in composition. The air in a poorly ventilated room occupied by a number of people rapidly changes in composition.
The properties of the air. Inasmuch as air is composed principally of a mixture of oxygen and nitrogen, which elements have already been discussed, its properties may be inferred largely from those of the two gases. One liter weighs 1.2923 g. It is thus 14.38 times as heavy as hydrogen. At the sea level it exerts an average pressure sufficient to sustain a column of mercury 760 mm. in height. This is taken as the standard pressure in determining the volumes of gases as well as the boiling points of liquids. Water may be made to boil at any temperature between 0 deg. and considerably above 100 deg. by simply varying the pressure. It is only when the pressure upon it is equal to the normal pressure of the atmosphere at the sea level, as indicated by a barometric reading of 760 mm., that it boils at 100 deg..
Preparation of liquid air. Attention has been called to the fact that both oxygen and nitrogen can be obtained in the liquid state by strongly cooling the gases and applying great pressure to them. Since air is largely a mixture of these two gases, it can be liquefied by the same methods.
The methods for liquefying air have been simplified greatly in that the low temperature required is obtained by allowing a portion of the compressed air to expand. The expansion of a gas is always attended by the absorption of heat. In liquefying air the apparatus is so constructed that the heat absorbed is withdrawn from air already under great pressure. This process is continued until the temperature is lowered to the point of liquefaction.
The Dewar bulb. It is not possible to preserve air in the liquid state in a closed vessel, on account of the enormous pressure exerted by it in its tendency to pass into the gaseous state. It may however be preserved for some hours or even days before it will completely evaporate, by simply placing it in an open vessel surrounded by a nonconducting material. The most efficient vessel for this purpose is the Dewar bulb shown in Fig. 29. The air is withdrawn from the space between the two walls, thus making it nonconducting.
Properties and uses of liquid air. When first prepared, liquid air is cloudy because of the presence of particles of solid carbon dioxide. These may be filtered off, leaving a liquid of slightly bluish color. It begins to boil at about -190 deg., the nitrogen passing off first, gradually followed by the oxygen, the last portions being nearly pure oxygen. To a certain extent oxygen is now prepared in this way for commercial purposes.
The extremely low temperature of liquid air may be inferred from the fact that mercury when cooled by it is frozen to a mass so hard that it may be used for driving nails.
Liquid air is used in the preparation of oxygen and as a cooling agent in the study of the properties of matter at low temperatures. It has thus been found that elements at extremely low temperatures largely lose their chemical activity.
EXERCISES
1. When oxygen and nitrogen are mixed in the proportion in which they exist in the atmosphere, heat is neither evolved nor absorbed by the process. What important point does this suggest?
2. What essential constituent of the air is found in larger amount in manufacturing districts than in the open country?
3. Can you suggest any reason why the growth of clover in a field improves the soil?
4. Why are the inner walls of a Dewar bulb sometimes coated with a film of silver?
5. To what is the blue color of liquid air due? Does this color increase in intensity on standing?
6. When ice is placed in a vessel containing liquid air, the latter boils violently. Explain.
7. Taking the volumes of the oxygen and nitrogen in 100 volumes of air as 21 and 78 respectively, calculate the percentages of these elements present by weight.
8. Would combustion be more intense in liquid air than in the gaseous substance?
9. A tube containing calcium chloride was found to weigh 30.1293 g. A volume of air which weighed 15.2134 g. was passed through, after which the weight of the tube was found to be 30.3405 g. What was the percentage amount of moisture present in the air?
10. 10 l. of air measured at 20 deg. and 740 mm. passed through lime water caused the precipitation of 0.0102 g. of CaCO_{3}. Find the number of volumes of carbon dioxide in 10,000 volumes of the air.
CHAPTER IX
SOLUTIONS
Definitions. When a substance disappears in a liquid in such a way as to thoroughly mix with it and to be lost to sight as an individual body, the resulting liquid is called a solution. The liquid in which the substance dissolves is called the solvent, while the dissolved substance is called the solute.
Classes of solutions. Matter in any one of its physical states may dissolve in a liquid, so that we may have solutions of gases, of liquids, and of solids. Solutions of liquids in liquids are not often mentioned in the following pages, but the other two classes will become very familiar in the course of our study, and deserve special attention.
SOLUTION OF GASES IN LIQUIDS
It has already been stated that oxygen, hydrogen, and nitrogen are slightly soluble in water. Accurate study has led to the conclusion that all gases are soluble to some extent not only in water but in many other liquids. The amount of a gas which will dissolve in a liquid depends upon a number of conditions, and these can best be understood by supposing a vessel B (Fig. 30), to be filled with the gas and inverted over the liquid. Under these circumstances the gas cannot escape or become mixed with another gas.
Circumstances affecting the solubility of gases. A number of circumstances affect the solubility of a gas in a liquid.
1. Nature of the gas. Other conditions being equal, each gas has its own peculiar solubility, just as it has its own special taste or odor. The solubility of gases varies between wide limits, as will be seen from the following table, but as a rule a given volume of a liquid will not dissolve more than two or three times its own volume of a gas.
Solubility of Gases in Water
1 l. of water at 760 mm. pressure and at 0 deg. will dissolve:
Ammonia 1148.00 l. Hydrochloric acid 503.00 Sulphur dioxide 79.79 Carbon dioxide 1.80 Oxygen 41.14 cc. Hydrogen 21.15 Nitrogen 20.03
In the case of very soluble gases, such as the first three in the table, it is probable that chemical combination between the liquid and the gas takes place.
2. Nature of the liquid. The character of the liquid has much influence upon the solubility of a gas. Water, alcohol, and ether have each its own peculiar solvent power. From the solubility of a gas in water, no prediction can be made as to its solubility in other liquids.
3. Influence of pressure. It has been found that the weight of gas which dissolves in a given case is proportional to the pressure exerted upon the gas. If the pressure is doubled, the weight of gas going into solution is doubled; if the pressure is diminished to one half of its original value, half of the dissolved gas will escape. Under high pressure, large quantities of gas can be dissolved in a liquid, and when the pressure is removed the gas escapes, causing the liquid to foam or effervesce.
4. Influence of temperature. In general, the lower the temperature of the liquid, the larger the quantity of gas which it can dissolve. 1000 volumes of water at 0 deg. will dissolve 41.14 volumes of oxygen; at 50 deg., 18.37 volumes; at 100 deg. none at all. While most gases can be expelled from a liquid by boiling the solution, some cannot. For example, it is not possible to expel hydrochloric acid gas completely from its solution by boiling.
SOLUTION OF SOLIDS IN LIQUIDS
This is the most familiar class of solutions, since in the laboratory substances are much more frequently used in the form of solutions than in the solid state.
Circumstances affecting the solubility of a solid. The solubility of a solid in a liquid depends upon several factors.
1. Nature of the solid. Other conditions being the same, solids vary greatly in their solubility in liquids. This is illustrated in the following table:
Table of Solubility of Solids at 18 deg.
100 cc. of water will dissolve:
Calcium chloride 71.0 g. Sodium chloride 35.9 Potassium nitrate 29.1 Copper sulphate 21.4 Calcium sulphate 0.207
No solids are absolutely insoluble, but the amount dissolved may be so small as to be of no significance for most purposes. Thus barium sulphate, one of the most insoluble of common substances, dissolves in water to the extent of 1 part in 400,000.
2. Nature of the solvent. Liquids vary much in their power to dissolve solids. Some are said to be good solvents, since they dissolve a great variety of substances and considerable quantities of them. Others have small solvent power, dissolving few substances, and those to a slight extent only. Broadly speaking, water is the most general solvent, and alcohol is perhaps second in solvent power.
3. Temperature. The weight of a solid which a given liquid can dissolve varies with the temperature. Usually it increases rapidly as the temperature rises, so that the boiling liquid dissolves several times the weight which the cold liquid will dissolve. In some instances, as in the case of common salt dissolved in water, the temperature has little influence upon the solubility, and a few solids are more soluble in cold water than in hot. The following examples will serve as illustrations:
Table of Solubility at 0 deg. and at 100 deg.
100 cc. of water will dissolve:
At 0 deg. At 100 deg.
Calcium chloride 49.6 g. 155.0 g. Sodium chloride 35.7 39.8 Potassium nitrate 13.3 247.0 Copper sulphate 15.5 73.5 Calcium sulphate 0.205 0.217 Calcium hydroxide 0.173 0.079
Saturated solutions. A liquid will not dissolve an unlimited quantity of a solid. On adding the solid to the liquid in small portions at a time, it will be found that a point is reached at which the liquid will not dissolve more of the solid at that temperature. The solid and the solution remain in contact with each other unchanged. This condition may be described by saying that they are in equilibrium with each other. A solution is said to be saturated when it remains unchanged in concentration in contact with some of the solid. The weight of the solid which will completely saturate a definite volume of a liquid at a given temperature is called the solubility of the substance at that temperature.
Supersaturated solutions. When a solution, saturated at a given temperature, is allowed to cool it sometimes happens that no solid crystallizes out. This is very likely to occur when the vessel used is perfectly smooth and the solution is not disturbed in any way. Such a solution is said to be supersaturated. That this condition is unstable can be shown by adding a crystal of the solid to the solution. All of the solid in excess of the quantity required to saturate the solution at this temperature will at once crystallize out, leaving the solution saturated. Supersaturation may also be overcome in many cases by vigorously shaking or stirring the solution.
General physical properties of solutions. A few general statements may be made in reference to the physical properties of solutions.
1. Distribution of the solid in the liquid. A solid, when dissolved, tends to distribute itself uniformly through the liquid, so that every part of the solution has the same concentration. The process goes on very slowly unless hastened by stirring or shaking the solution. Thus, if a few crystals of a highly colored substance such as copper sulphate are placed in the bottom of a tall vessel full of water, it will take weeks for the solution to become uniformly colored.
2. Boiling points of solutions. The boiling point of a liquid is raised by the presence of a substance dissolved in it. In general the extent to which the boiling point of a solvent is raised by a given substance is proportional to the concentration of the solution, that is, to the weight of the substance dissolved in a definite weight of the solvent.
3. Freezing points of solutions. A solution freezes at a lower temperature than the pure solvent. The lowering of the freezing point obeys the same law which holds for the raising of the boiling point: the extent of lowering is proportional to the weight of dissolved substance, that is, to the concentration of the solution.
Electrolysis of solutions. Pure water does not appreciably conduct the electric current. If, however, certain substances such as common salt are dissolved in the water, the resulting solutions are found to be conductors of electricity. Such solutions are called electrolytes. When the current passes through an electrolyte some chemical change always takes place. This change is called electrolysis.
The general method used in the electrolysis of a solution is illustrated in Fig. 31. The vessel D contains the electrolyte. Two plates or rods, A and B, made of suitable material, are connected with the wires from a battery (or dynamo) and dipped into the electrolyte, as shown in the figure. These plates or rods are called electrodes. The electrode connected with the zinc plate of the battery is the negative electrode or cathode, while that connected with the carbon plate is the positive electrode or anode.
Theory of electrolytic dissociation. The facts which have just been described in connection with solutions, together with many others, have led chemists to adopt a theory of solutions called the theory of electrolytic dissociation. The main assumptions in this theory are the following.
1. Formation of ions. Many compounds when dissolved in water undergo an important change. A portion of their molecules fall apart, or dissociate, into two or more parts, called ions. Thus sodium nitrate (NaNO{3}) dissociates into the ions Na and NO{3}; sodium chloride, into the ions Na and Cl. These ions are free to move about in the solution independently of each other like independent molecules, and for this reason were given the name ion, which signifies a wanderer.
2. The electrical charge of ions. Each ion carries a heavy electrical charge, and in this respect differs from an atom or molecule. It is evident that the sodium in the form of an ion must differ in some important way from ordinary sodium, for sodium ions, formed from sodium nitrate, give no visible evidence of their presence in water, whereas metallic sodium at once decomposes the water. The electrical charge, therefore, greatly modifies the usual chemical properties of the element.
3. The positive charges equal the negative charges. The ions formed by the dissociation of any molecule are of two kinds. One kind is charged with positive electricity and the other with negative electricity; moreover the sum of all the positive charges is always equal to the sum of all the negative charges. The solution as a whole is therefore electrically neutral. If we represent dissociation by the usual chemical equations, with the electrical charges indicated by + and - signs following the symbols, the dissociation of sodium chloride molecules is represented thus:
NaCl —> Na^{+}, Cl^{-}.
The positive charge on each sodium ion exactly equals the negative charge on each chlorine ion. Sodium sulphate dissociates, as shown in the equation
Na_{2}SO_{4} —> 2Na^{+}, SO_{4}^{—}.
Here the positive charge on the two sodium ions equals the double negative charge on the SO_{4} ion.
4. Not all compounds dissociate. Only those compounds dissociate whose solutions form electrolytes. Thus salt dissociates when dissolved in water, the resulting solution being an electrolyte. Sugar, on the other hand, does not dissociate and its solution is not a conductor of the electric current.
5. Extent of dissociation differs in different liquids. While compounds most readily undergo dissociation in water, yet dissociation often occurs to a limited extent when solution takes place in liquids other than water. In the discussion of solutions it will be understood that the solvent is water unless otherwise noted.
The theory of electrolytic dissociation and the properties of solutions. In order to be of value, this theory must give a reasonable explanation of the properties of solutions. Let us now see if the theory is in harmony with certain of these properties.
The theory of electrolytic dissociation and the boiling and freezing points of solutions. We have seen that the boiling point of a solution of a substance is raised in proportion to the concentration of the dissolved substance. This is but another way of saying that the change in the boiling point of the solution is proportional to the number of molecules of the dissolved substance present in the solution.
It has been found, however, that in the case of electrolytes the boiling point is raised more than it should be to conform to this law. If the solute dissociates into ions, the reason for this becomes clear. Each ion has the same effect on the boiling point as a molecule, and since their number is greater than the number of molecules from which they were formed, the effect on the boiling point is abnormally great.
In a similar way, the theory furnishes an explanation of the abnormal lowering of the freezing point of electrolytes.
The theory of electrolytic dissociation and electrolysis. The changes taking place during electrolysis harmonize very completely with the theory of dissociation. This will become clear from a study of the following examples. |
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