Appearance of flames. The flame caused by the union of hydrogen and oxygen is almost colorless and invisible. Chlorine and hydrogen combine with a pale violet flame, carbon monoxide burns in oxygen with a blue flame, while ammonia burns with a deep yellow flame. The color and appearance of flames are therefore often quite characteristic of the particular combustion which occasions them.

Structure of flames. When the gas undergoing combustion issues from a round opening into an atmosphere of the gas supporting combustion, as is the case with the burning Bunsen burner (Fig. 63), the flame is generally conical in outline. It consists of several distinct cones, one within the other, the boundary between them being marked by differences of color or luminosity. In the simplest flame, of which hydrogen burning in oxygen is a good example, these cones are two in number,—an inner one, formed by unburned gas, and an outer one, usually more or less luminous, consisting of the combining gases. This outer one is in turn surrounded by a third envelope of the products of combustion; this envelope is sometimes invisible, as in the present case, but is sometimes faintly luminous. The lower part of the inner cone of the flame is quite cool and consists of unburned gas. Toward the top of the inner cone the gas has become heated to a high temperature by the burning envelope surrounding it. On reaching the supporter of combustion on the outside it is far above its kindling temperature, and combustion follows with the evolution of much heat. The region of combustion just outside the inner cone is therefore the hottest part of the flame.



Oxidizing and reducing flames. Since the tip of the outside cone consists of very hot products of combustion mixed with oxygen from the air, a substance capable of oxidation placed in this part of the flame becomes very hot and is easily oxidized. The oxygen with which it combines comes, of course, from the atmosphere, and not from the products of combustion. This outer tip of the flame is called the oxidizing flame.

At the tip of the inner cone the conditions are quite different. This region consists of a highly heated combustible gas, which has not yet reached a supply of oxygen.

If a substance rich in oxygen, such as a metallic oxide, is placed in this region of the flame, the heated gases combine with its oxygen and the substance is reduced. This part of the flame is called the reducing flame. These flames are used in testing certain substances, especially minerals. For this purpose they are produced by blowing into a small luminous Bunsen flame from one side through a blowpipe. This is a tube of the shape shown in Fig. 64. The flame is directed in any desired way and has the oxidizing and reducing regions very clearly marked (Fig. 65). It is non-luminous from the same causes which render the open Bunsen burner flame non-luminous, the gases from the lungs serving to furnish oxygen and to dilute the combustible gas.



Luminosity of flames. The luminosity of flames is due to a number of distinct causes, and may therefore be increased or diminished in several ways.

1. Presence of solid matter. The most obvious of these causes is the presence in the flame of incandescent solid matter. Thus chalk dust sifted into a non-luminous flame renders it luminous. When hydrocarbons form a part of the combustible gas, as they do in nearly all illuminating gases and oils, some carbon is usually set free in the process of combustion. This is made very hot by the flame and becomes incandescent, giving out light. In a well-regulated flame it is afterward burned up, but when the supply of oxygen is insufficient it escapes from the flame as lampblack or soot. That it is temporarily present in a well-burning luminous flame may be demonstrated by holding a cold object, such as a small evaporating dish, in the flame for a few seconds. This cold object cools the carbon below its kindling temperature, and it is deposited on the object as soot.

2. Pressure. A second factor in the luminosity of flames is the pressure under which the gases are burning. Under increased pressure there is more matter in a given volume of a gas, and the chemical action is more energetic than when the gases are rarefied. Consequently there is more heat and light. A candle burning on a high mountain gives less light than when it burns at the sea level.

If the gas is diluted with a non-combustible gas, the effect is the same as if it is rarefied, for under these conditions there is less combustible gas in a given volume.

3. Temperature. The luminosity also depends upon the temperature attained in the combustion. In general the hotter the flame the greater the luminosity; hence cooling the gases before combustion diminishes the luminosity of the flame they will make, because it diminishes the temperature attained in the combustion. Thus the luminosity of the Bunsen flame is largely diminished by the air drawn up with the gas. This is due in part to the fact that the burning gas is diluted and cooled by the air drawn in. The oxygen thus introduced into the flame also causes the combustion of the hot particles of carbon which would otherwise tend to make the flame luminous.

Illuminating and fuel gases. A number of mixtures of combustible gases, consisting largely of carbon compounds and hydrogen, find extensive use for the production of light and heat. The three chief varieties are coal gas, water gas, and natural gas. The use of acetylene gas has already been referred to.

Coal gas. Coal gas is made by heating bituminous coal in large retorts out of contact with the air. Soft or bituminous coal contains, in addition to large amounts of carbon, considerable quantities of compounds of hydrogen, oxygen, nitrogen, and sulphur. When distilled the nitrogen is liberated partly in the form of ammonia and cyanides and partly as free nitrogen gas; the sulphur is converted into hydrogen sulphide, carbon disulphide, and oxides of sulphur; the oxygen into water and oxides of carbon. The remaining hydrogen is set free partly as hydrogen and partly in combination with carbon in the form of hydrocarbons. The most important of these is methane, with smaller quantities of many others, some of which are liquids or solids at ordinary temperatures. The great bulk of the carbon remains behind as coke and retort carbon.

The manufacture of coal gas. In the manufacture of coal gas it is necessary to separate from the volatile constituents formed by the heating of the coal all those substances which are either solid or liquid at ordinary temperature, since these would clog the gas pipes. Certain gaseous constituents, such as hydrogen sulphide and ammonia, must also be removed. The method used to accomplish this is shown in Fig. 66. The coal is heated in air-tight retorts illustrated by A. The volatile products escape through the pipe X and bubble into the tarry liquid in the large pipe B, known as the hydraulic main, which runs at right angles to the retorts. Here is deposited the greater portion of the solid and liquid products, forming a tarry mass known as coal tar. Much of the ammonia also remains dissolved in this liquid. The partially purified gas then passes into the pipes C, which serve to cool it and further remove the solid and liquid matter. The gas then passes into D, which is filled with coke over which a jet of water is sprayed. The water still further cools the gas and at the same time partially removes such gaseous products as hydrogen sulphide and ammonia, which are soluble in water. In E the gas passes over some material such as lime, which removes the last portions of the sulphur compounds as well as much of the carbon dioxide present. From E the gas passes into the large gas holder F, from which it is distributed through pipes to the places where it is burned.



One ton of good gas coal yields approximately 10,000 cu. ft. of gas, 1400 lb. of coke, 120 lb. of tar, and 20 gal. of ammoniacal liquor.

Not only is the ammonia obtained in the manufacture of the gas of great importance, but the coal tar also serves as the source of many very useful substances, as will be explained in Chapter XXXII.

Water gas. Water gas is essentially a mixture of carbon monoxide and hydrogen. It is made by passing steam over very hot anthracite coal, when the reaction shown in the following equation takes place:

C + H_{2}O = CO + 2H.

When required merely to produce heat the gas is at once ready for use. When made for illuminating purposes it must be enriched, that is, illuminants must be added, since both carbon monoxide and hydrogen burn with non-luminous flames. This is accomplished by passing it into heaters containing highly heated petroleum oils. The gas takes up hydrocarbon gases formed in the decomposition of the petroleum oils, which make it burn with a luminous flame.

Water gas is very effective as a fuel, since both carbon monoxide and hydrogen burn with very hot flames. It has little odor and is very poisonous. Its use is therefore attended with some risk, since leaks in pipes are very likely to escape notice.

Natural gas. This substance, so abundant in many localities, varies much in composition, but is composed principally of methane. When used for lighting purposes it is usually burned in a burner resembling an open Bunsen, the illumination being furnished by an incandescent mantle. This is the case in the familiar Welsbach burner. Contrary to statements frequently made, natural gas contains no free hydrogen.

TABLE SHOWING COMPOSITION OF GASES

======================================================= PENNSYLVANIA COAL WATER ENRICHED NATURAL GAS GAS WATER GAS GAS - Hydrogen 41.3 52.88 30.00 Methane 90.64 43.6 2.16 24.00 Illuminants 3.9 12.05 Carbon monoxide 6.4 36.80 29.00 Carbon dioxide 0.30 2.0 3.47 0.30 Nitrogen 9.06 1.2 4.69 2.50 Oxygen 0.3 1.50 Hydrocarbon vapors 1.5 1.50 ===================================================

These are analyses of actual samples, and may be taken as about the average for the various kinds of gases. Any one of these may vary considerably. The nitrogen and oxygen in most cases is due to a slight admixture of air which is difficult to exclude entirely in the manufacture and handling of gases.

Fuels. A variety of substances are used as fuels, the most important of them being wood, coal, and the various gases mentioned above. Wood consists mainly of compounds of carbon, hydrogen, and oxygen. The composition of coal and the fuel gases has been given. Since these fuels are composed principally of carbon and hydrogen or their compounds, the chief products of combustion are carbon dioxide and water. The practice of heating rooms with portable gas or oil stoves with no provision for removing the products of combustion is to be condemned, since the carbon dioxide is generated in sufficient quantities to render the air unfit for breathing. Rooms so heated also become very damp from the large amount of water vapor formed in the combustion, and which in cold weather condenses on the window glass, causing the glass to "sweat." Both coal and wood contain a certain amount of mineral substances which constitute the ashes.

The electric furnace. In recent years electric furnaces have come into wide use in operations requiring a very high temperature. Temperatures as high as 3500 deg. can be easily reached, whereas the hottest oxyhydrogen flame is not much above 2000 deg.. These furnaces are constructed on one of two general principles.



1. Arc furnaces. In the one type the source of heat is an electric arc formed between carbon electrodes separated a little from each other, as shown in Fig. 67. The substance to be heated is placed in a vessel, usually a graphite crucible, just below the arc. The electrodes and crucible are surrounded by materials which fuse with great difficulty, such as magnesium oxide, the walls of the furnace being so shaped as to reflect the heat downwards upon the contents of the crucible.



2. Resistance furnaces. In the other type of furnace the heat is generated by the resistance offered to the current in its passage through the furnace. In its simplest form it may be represented by Fig. 68. The furnace is merely a rectangular box built up of loose bricks. The electrodes E, each consisting of a bundle of carbon rods, are introduced through the sides of the furnace. The materials to be heated, C, are filled into the furnace up to the electrodes, and a layer of broken coke is arranged so as to extend from one electrode to the other. More of the charge is then placed on top of the coke. In passing through the broken coke the electrical current encounters great resistance. This generates great heat, and the charge surrounding the coke is brought to a very high temperature. The advantage of this type of furnace is that the temperature can be regulated to any desired intensity.

EXERCISES

1. Why does charcoal usually burn with no flame? How do you account for the flame sometimes observed when it burns?

2. How do you account for the fact that a candle burns with a flame?

3. What two properties must the mantle used in the Welsbach lamp possess?

4. (a) In what respects does the use of the Welsbach mantle resemble that of lime in the calcium light? (b) If the mantle were made of carbon, would it serve the same purpose?

5. Would anthracite coal be suitable for the manufacture of coal gas?

6. How could you prove the formation of carbon dioxide and water in the combustion of illuminating gases?

7. Suggest a probable way in which natural gas has been formed.

8. Coal frequently contains a sulphide of iron. (a) What two sulphur compounds are likely to be formed when gas is made from such coal? (b) Suggest some suitable method for the removal of these compounds.

9. Why does the use of the bellows on the blacksmith's forge cause a more intense heat?

10. What volume of oxygen is necessary to burn 100 l. of marsh gas and what volume of carbon dioxide would be formed, all of the gases being measured under standard conditions?

11. Suppose a cubic meter of Pennsylvania natural gas, measured under standard conditions, were to be burned. How much water by weight would result?



CHAPTER XIX

MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS

Introduction. In the chapter on The Atomic Theory, it was shown that if it were true that two elements uniting to form a compound always combined in the ratio of one atom of one element to one atom of the other element, it would be a very easy matter to decide upon figures which would represent the relative weights of the different atoms. It would only be necessary to select some one element as a standard and determine the weight of every element which combines with a definite weight (say 1 g.) of the standard element. The figures so obtained would evidently represent the relative weights of the atoms.

But the law of multiple proportion at once reminds us that two elements may unite in several proportions; and there is no simple way to determine the number of atoms present in the molecule of any compound. Consequently the problem of deciding upon the relative atomic weights is not an easy one. To the solution of this problem we must now turn.

Dalton's method of determining atomic weights. When Dalton first advanced the atomic theory he attempted to solve this problem by very simple methods. He thought that when only one compound of two elements is known it is reasonable to suppose that it contains one atom of each element. He therefore gave the formula HO to water, and HN to ammonia. When more than two compounds were known he assumed that the most familiar or the most stable one had the simple formula. He then determined the atomic weight as explained above. The results he obtained were contradictory and very far from satisfactory, and it was soon seen that some other method, resting on much more scientific grounds, must be found to decide what compounds, if any, have a single atom of each element present.

Determination of atomic weights. Three distinct steps are involved in the determination of the atomic weight of an element: (1) determination of the equivalent, (2) determination of molecular weights of its compounds, and (3) deduction of the exact atomic weight from the equivalent and molecular weights.

1. Determination of the equivalent. By the equivalent of an element is meant the weight of the element which will combine with a fixed weight of some other element chosen as a standard. It has already been explained that oxygen has been selected as the standard element for atomic weights, with a weight of 16. This same standard will serve very well as a standard for equivalents. The equivalent of an element is the weight of the element which will combine with 16 g. of oxygen. Thus 16 g. of oxygen combines with 16.03 g. of sulphur, 65.4 g. of zinc, 215.86 g. of silver, 70.9 g. of chlorine. These figures, therefore, represent the equivalent weights of these elements.

Relation of atomic weights to equivalents. According to the atomic theory combination always takes place between whole numbers of atoms. Thus one atom unites with one other, or with two or three; or two atoms may unite with three, or three with five, and so on.

When oxygen combines with zinc the combination must be between definite numbers of the two kinds of atoms. Experiment shows that these two elements combine in the ratio of 16 g. of oxygen to 65.4 g. of zinc. If one atom of oxygen combines with one atom of zinc, then this ratio must be the ratio between the weights of the two atoms. If one atom of oxygen combines with two atoms of zinc, then the ratio between the weights of the two atoms will be 16: 32.7. If two atoms of oxygen combine with one atom of zinc, the ratio by weight between the two atoms will be 8: 65.4. It is evident, therefore, that the real atomic weight of an element must be some multiple or submultiple of the equivalent; in other words, the equivalent multiplied by 1/2, 1, 2, or 3 will give the atomic weight.

Combining weights. A very interesting relation holds good between the equivalents of the various elements. We have just seen that the figures 16.03, 65.4, 215.86, and 70.9 are the equivalents respectively of sulphur, zinc, silver, and chlorine. These same figures represent the ratios by weight in which these elements combine among themselves. Thus 215.86 g. of silver combine with 70.9 g. of chlorine and with 2 x 16.03 g. of sulphur. 65.4 g. of zinc combine with 70.9 g. of chlorine and 2 x 16.03 g. of sulphur.

By taking the equivalent or some multiple of it a value can be obtained for each element which will represent its combining value, and for this reason is called its combining weight. It is important to notice that the fact that a combining weight can be obtained for each element is not a part of a theory, but is the direct result of experiment.

Elements with more than one equivalent. It will be remembered that oxygen combines with hydrogen in two ratios. In one case 16 g. of oxygen combine with 2.016 g. of hydrogen to form water; in the other 16 g. of oxygen combine with 1.008 g. of hydrogen to form hydrogen dioxide. The equivalents of hydrogen are therefore 2.016 and 1.008. Barium combines with oxygen in two proportions: in barium oxide the proportion is 16 g. of oxygen to 137.4 g. of barium; in barium dioxide the proportion is 16 g. of oxygen to 68.7 g. of barium.

In each case one equivalent is a simple multiple of the other, so the fact that there may be two equivalents does not add to the uncertainty. All we knew before was that the true atomic weight is some multiple of the equivalent.

2. The determination of molecular weights. To decide the question as to which multiple of the equivalent correctly represents the atomic weight of an element, it has been found necessary to devise a method of determining the molecular weights of compounds containing the element in question. Since the molecular weight of a compound is merely the sum of the weights of all the atoms present in it, it would seem to be impossible to determine the molecular weight of a compound without first knowing the atomic weights of the constituent atoms, and how many atoms of each element are present in the molecule. But certain facts have been discovered which suggest a way in which this can be done.

Avogadro's hypothesis. We have seen that the laws of Boyle, Charles, and Gay-Lussac apply to all gases irrespective of their chemical character. This would lead to the inference that the structure of gases must be quite simple, and that it is much the same in all gases.

In 1811 Avogadro, an Italian physicist, suggested that if we assume all gases under the same conditions of temperature and pressure to have the same number of molecules in a given volume, we shall have a probable explanation of the simplicity of the gas laws. It is difficult to prove the truth of this hypothesis by a simple experiment, but there are so many facts known which are in complete harmony with this suggestion that there is little doubt that it expresses the truth. Avogadro's hypothesis may be stated thus: Equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules.

Avogadro's hypothesis and molecular weights. Assuming that Avogadro's hypothesis is correct, we have a very simple means for deciding upon the relative weights of molecules; for if equal volumes of two gases contain the same number of molecules, the weights of the two volumes must be in the same ratio as the weights of the individual molecules which they contain. If we adopt some one gas as a standard, we can express the weights of all other gases as compared with this one, and the same figures will express the relative weights of the molecules of which the gases are composed.

Oxygen as the standard. It is important that the same standard should be adopted for the determination of molecular weights as has been decided upon for atomic weights and equivalents, so that the three values may be in harmony with each other. Accordingly it is best to adopt oxygen as the standard element with which to compare the molecular weights of other gases, being careful to keep the oxygen atom equal to 16.

The oxygen molecule contains two atoms. One point must not be overlooked, however. We desire to have our unit, the oxygen atom, equal to 16. The method of comparing the weights of gases just suggested compares the molecules of the gases with the molecule of oxygen. Is the molecule and the atom of oxygen the same thing? This question is answered by the following considerations.

We have seen that when steam is formed by the union of oxygen and hydrogen, two volumes of hydrogen combine with one volume of oxygen to form two volumes of steam. Let us suppose that the one volume of oxygen contains 100 molecules; then the two volumes of steam must, according to Avogadro's hypothesis, contain 200 molecules. But each of these 200 molecules must contain at least one atom of oxygen, or 200 in all, and these 200 atoms came from 100 molecules of oxygen. It follows that each molecule of oxygen must contain at least two atoms of oxygen.

Evidently this reasoning merely shows that there are at least two atoms in the oxygen molecule. There may be more than that, but as there is no evidence to this effect, we assume that the molecule contains two atoms only.

It is evident that if we wish to retain the value 16 for the atom of oxygen we must take twice this value, or 32, for the value of the oxygen molecule, when using it as a standard for molecular weights.

Determination of the molecular weights of gases from their weights compared with oxygen. Assuming the molecular weight of oxygen to be 32, Avogadro's hypothesis gives us a ready means for determining the molecular weight of any other gas, for all that is required is to know its weight compared with that of an equal volume of oxygen. For example, 1 l. of chlorine is found by experiment to weigh 2.216 times as much as 1 l. of oxygen. The molecular weight of chlorine must therefore be 2.216 x 32, or 70.91.

If, instead of comparing the relative weights of 1 l. of the two gases, we select such a volume of oxygen as will weigh 32 g., or the weight in grams corresponding to the molecular weight of the gas, the calculation is much simplified. It has been found that 32 g. of oxygen, under standard conditions, measure 22.4 l. This same volume of hydrogen weighs 2.019 g.; of chlorine 70.9 g.; of hydrochloric acid 36.458 g. The weights of these equal volumes must be proportional to their molecular weights, and since the weight of the oxygen is the same as the value of its molecular weight, so too will the weights of the 22.4 l. of the other gases be equal to the value of their molecular weights.

As a summary we can then make the following statement: The molecular weight of any gas may be determined by calculating the weight of 22.4 l. of the gas, measured under standard conditions.

Determination of molecular weights from density of gases. In an actual experiment it is easier to determine the density of a gas than the weight of a definite volume of it. The density of a gas is usually defined as its weight compared with that of an equal volume of air. Having determined the density of a gas, its weight compared with oxygen may be determined by multiplying its density by the ratio between the weights of air and oxygen. This ratio is 0.9046. To compare it with our standard for atomic weights we must further multiply it by 32, since the standard is 1/32 the weight of oxygen molecules. The steps then are these:

1. Determine the density of the gas (its weight compared with air).

2. Multiply by 0.9046 to make the comparison with oxygen molecules.

3. Multiply by 32 to make the comparison with the unit for atomic weights.

We have, then, the formula:

molecular weight = density x 0.9046 x 32;

or, still more briefly,

M. = D. x 28.9.

The value found by this method for the determination of molecular weights will of course agree with those found by calculating the weight of 22.4 l. of the gas, since both methods depend on the same principles.



Determination of densities of gases. The relative weights of equal volumes of two gases can be easily determined. The following is one of the methods used. A small flask, such as is shown in Fig. 69, is filled with one of the gases, and after the temperature and pressure have been noted the flask is sealed up and weighed. The tip of the sealed end is then broken off, the flask filled with the second gas, and its weight determined. If the weight of the empty flask is subtracted from these two weighings, the relative weights of the gases is readily found.

3. Deduction of atomic weights from molecular weights and equivalents. We have now seen how the equivalent of an element and the molecular weight of compounds containing the element can be obtained. Let us see how it is possible to decide which multiple of the equivalent really is the true atomic weight. As an example, let us suppose that the equivalent of nitrogen has been found to be 7.02 and that it is desired to obtain its atomic weight. The next step is to obtain the molecular weights of a large number of compounds containing nitrogen. The following will serve:

========================================================= APPROXIMATE PERCENTAGE OF PART OF DENSITY BY MOLECULAR NITROGEN BY MOLECULAR EXPERIMENT WEIGHT EXPERIMENT WEIGHT DUE (D. x 28.9) TO NITROGEN - Nitrogen gas 0.9671 27.95 100.00 27.95 Nitrous oxide 1.527 44.13 63.70 27.11 Nitric oxide 1.0384 30.00 46.74 14.02 Nitrogen peroxide 1.580 45.66 30.49 13.90 Ammonia 0.591 17.05 82.28 14.03 Nitric acid 2.180 63.06 22.27 14.03 Hydrocyanic acid 0.930 26.87 51.90 13.94 =========================================================

Method of calculation. The densities of the various gases in the first column of this table are determined by experiment, and are fairly accurate but not entirely so. By multiplying these densities by 28.9 the molecular weights of the compounds as given in the second column are obtained. By chemical analysis it is possible to determine the percentage composition of these substances, and the percentages of nitrogen in them as determined by analysis are given in the third column. If each of these molecular weights is multiplied in turn by the percentage of nitrogen in the compound, the product will be the weight of the nitrogen in the molecular weight of the compound. This will be the sum of the weights of the nitrogen atoms in the molecule. These values are given in the fourth column in the table.

If a large number of compounds containing nitrogen are studied in this way, it is probable that there will be included in the list at least one substance whose molecule contains a single nitrogen atom. In this case the number in the fourth column will be the approximate atomic weight of nitrogen. On comparing the values for nitrogen in the table it will be seen that a number which is approximately 14 is the smallest, and that the others are multiples of this. These compounds of higher value, therefore, contain more than one nitrogen atom in the molecule.

Accurate determination of atomic weights. Molecular weights cannot be determined very accurately, and consequently the part in them due to nitrogen is a little uncertain, as will be seen in the table. All we can tell by this method is that the true weight is very near 14. The equivalent can however be determined very accurately, and we have seen that it is some multiple or submultiple of the true atomic weight. Since molecular-weight determinations have shown that in the case of nitrogen the atomic weight is near 14, and we have found the equivalent to be 7.02, it is evident that the true atomic weight is twice the equivalent, or 7.02 x 2 = 14.04.

Summary. These, then, are the steps necessary to establish the atomic weight of an element.

1. Determine the equivalent accurately by analysis.

2. Determine the molecular weight of a large number of compounds of the element, and by analysis the part of the molecular weight due to the element. The smallest number so obtained will be approximately the atomic weight.

3. Multiply the equivalent by the small whole number (usually 1, 2, or 3), which will make a number very close to the approximate atomic weight. The figure so obtained will be the true atomic weight.

Molecular weights of the elements. It will be noticed that the molecular weight of nitrogen obtained by multiplying its density by 28.9 is 28.08. Yet the atomic weight of nitrogen as deduced from a study of its gaseous compounds is 14.04. The simplest explanation that can be given for this is that the gaseous nitrogen is made up of molecules, each of which contains two atoms. In this respect it resembles oxygen; for we have seen that an entirely different line of reasoning leads us to believe that the molecule of oxygen contains two atoms. When we wish to indicate molecules of these gases the symbols N{2} and O{2} should be used. When we desire to merely show the weights taking part in a reaction this is not necessary.

The vapor densities of many of the elements show that, like oxygen and nitrogen, their molecules consist of two atoms. In other cases, particularly among the metals, the molecule and the atom are identical. Still other elements have four atoms in their molecules.

While oxygen contains two atoms in its molecules, a study of ozone has led to the conclusion that it has three. The formation of ozone from oxygen can therefore be represented by the equation

3O{2} = 2O{3}.

Other methods of determining molecular weights. It will be noticed that Avogadro's law gives us a method by which we can determine the relative weights of the molecules of two gases because it enables us to tell when we are dealing with an equal number of the two kinds of molecules. If by any other means we can get this information, we can make use of the knowledge so gained to determine the molecular weights of the two substances.

Raoult's laws. Two laws have been discovered which give us just such information. They are known as Raoult's laws, and can be stated as follows:

1. When weights of substances which are proportional to their molecular weights are dissolved in the same weight of solvent, the rise of the boiling point is the same in each case.

2. When weights of substances which are proportional to their molecular weights are dissolved in the same weight of solvent, the lowering of the freezing point is the same in each case.

By taking advantage of these laws it is possible to determine when two solutions contain the same number of molecules of two dissolved substances, and consequently the relative molecular weights of the two substances.

Law of Dulong and Petit. In 1819 Dulong and Petit discovered a very interesting relation between the atomic weight of an element and its specific heat, which holds true for elements in the solid state. If equal weights of two solids, say, lead and silver, are heated through the same range of temperature, as from 10 deg. to 20 deg., it is found that very different amounts of heat are required. The amount of heat required to change the temperature of a solid or a liquid by a definite amount compared with the amount required to change the temperature of an equal weight of water by the same amount is called its specific heat. Dulong and Petit discovered the following law: The specific heat of an element in the solid form multiplied by its atomic weight is approximately equal to the constant 6.25. That is,

at. wt. x sp. ht. = 6.25.

Consequently,

6.25 at. wt. = ———— sp. ht.

This law is not very accurate, but it is often possible by means of it to decide upon what multiple of the equivalent is the real atomic weight. Thus the specific heat of iron is found by experiment to be 0.112, and its equivalent is 27.95. 6.25 / 0.112 = 55.8. We see, therefore, that the atomic weight is twice the equivalent, or 55.9.

How formulas are determined. It will be well in connection with molecular weights to consider how the formula of a compound is decided upon, for the two subjects are very closely associated. Some examples will make clear the method followed.

The molecular weight of a substance containing hydrogen and chlorine was 36.4. By analysis 36.4 parts of the substance was found to contain 1 part of hydrogen and 35.4 parts of chlorine. As these are the simple atomic weights of the two elements, the formula of the compound must be HCl.

A substance consisting of oxygen and hydrogen was found to have a molecular weight of 34. Analysis showed that in 34 parts of the substance there were 2 parts of hydrogen and 32 parts of oxygen. Dividing these figures by the atomic weights of the two elements, we get 2 / 1 = 2 for H; 32 / 16 = 2 for O. The formula is therefore H{2}O{2}.

A substance containing 2.04% H, 32.6% S, and 65.3% O was found to have a molecular weight of 98. In these 98 parts of the substance there are 98 x 2.04% = 2 parts of H, 98 x 32.6% = 32 parts of S, and 98 x 65.3% = 64 parts of O. If the molecule weighs 98, the hydrogen atoms present must together weigh 2, the sulphur atoms 32, and the oxygen atoms 64. Dividing these figures by the respective atomic weights of the three elements, we have, for H, 2 / 1 = 2 atoms; for S, 32 / 32 = 1 atom; for O, 64 / 16 = 4 atoms. Hence the formula is H{2}SO{4}.

We have, then, this general procedure: Find the percentage composition of the substance and also its molecular weight. Multiply the molecular weight successively by the percentage of each element present, to find the amount of the element in the molecular weight of the compound. The figures so obtained will be the respective parts of the molecular weight due to the several atoms. Divide by the atomic weights of the respective elements, and the quotient will be the number of atoms present.

Avogadro's hypothesis and chemical calculations. This law simplifies many chemical calculations.

1. Application to volume relations in gaseous reactions. Since equal volumes of gases contain an equal number of molecules, it follows that when an equal number of gaseous molecules of two or more gases take part in a reaction, the reaction will involve equal volumes of the gases. In the equation

C_{2}H_{2}O_{4} = H_{2}O + CO_{2} + CO,

since 1 molecule of each of the gases CO_{2} and CO is set free from each molecule of oxalic acid, the two substances must always be set free in equal volumes.

Acetylene burns in accordance with the equation

2C_{2}H_{2} + 5O_{2} = 4CO_{2} + 2H_{2}O.

Hence 2 volumes of acetylene will react with 5 volumes of oxygen to form 4 volumes of carbon dioxide and 2 volumes of steam. That the volume relations may be correct a gaseous element must be given its molecular formula. Thus oxygen must be written O_{2} and not 2O.

2. Application to weights of gases. It will be recalled that the molecular weight of a gas is determined by ascertaining the weight of 22.4 l. of the gas. This weight in grams is called the gram-molecular weight of a gas. If the molecular weight of any gas is known, the weight of a liter of the gas under standard conditions may be determined by dividing its gram-molecular weight by 22.4. Thus the gram-molecular weight of a hydrochloric acid gas is 36.458. A liter of the gas will therefore weigh 36.458 / 22.4 = 1.627 g.

EXERCISES

1. From the following data calculate the atomic weight of sulphur. The equivalent, as obtained by an analysis of sulphur dioxide, is 16.03. The densities and compositions of a number of compounds containing sulphur are as follows:

NAME DENSITY COMPOSITION BY PERCENTAGE Hydrosulphuric acid 1.1791 S = 94.11 H = 5.89 Sulphur dioxide 2.222 S = 50.05 O = 49.95 Sulphur trioxide 2.74 S = 40.05 O = 59.95 Sulphur chloride 4.70 S = 47.48 Cl = 52.52 Sulphuryl chloride 4.64 S = 23.75 Cl = 52.53 O = 23.70 Carbon disulphide 2.68 S = 84.24 C = 15.76

2. Calculate the formulas for compounds of the following compositions:

MOLECULAR WEIGHT (1) S = 39.07% O = 58.49% H = 2.44% 81.0 (2) Ca = 29.40 S = 23.56 O = 47.04 136.2 (3) K = 38.67 N = 13.88 O = 47.45 101.2

3. The molecular weight of ammonia is 17.06; of sulphur dioxide is 64.06; of chlorine is 70.9. From the molecular weight calculate the weight of 1 l. of each of these gases. Compare your results with the table on the back cover of the book.

4. From the molecular weight of the same gases calculate the density of each, referred to air as a standard.

5. A mixture of 50 cc. of carbon monoxide and 50 cc. of oxygen was exploded in a eudiometer, (a) What gases remained in the tube after the explosion? (b) What was the volume of each?

6. In what proportion must acetylene and oxygen be mixed to produce the greatest explosion?

7. Solve Problem 18, Chapter XVII, without using molecular weights. Compare your results.

8. Solve Problem 10, Chapter XVIII, without using molecular weights. Compare your results.

9. The specific heat of aluminium is 0.214; of lead is 0.031. From these specific heats calculate the atomic weights of each of the elements.



CHAPTER XX

THE PHOSPHORUS FAMILY

================================================== ATOMIC MELTING SYMBOL WEIGHT DENSITY POINT - - - - Phosphorus P 31.0 1.8 43.3 deg. Arsenic As 75.0 5.73 - Antimony Sb 120.2 6.7 432 deg. Bismuth Bi 208.5 9.8 270 deg. ==================================================

The family. The elements constituting this family belong in the same group with nitrogen and therefore resemble it in a general way. They exhibit a regular gradation of physical properties, as is shown in the above table. The same general gradation is also found in their chemical properties, phosphorus being an acid-forming element, while bismuth is essentially a metal. The other two elements are intermediate in properties.

Compounds. In general the elements of the family form compounds having similar composition, as is shown in the following table:

PH_{3} PCl_{3} PCl_{5} P_{2}O_{3} P_{2}O_{5} AsH_{3} AsCl_{3} AsCl_{5} As_{2}O_{3} As_{2}O_{5} SbH_{3} SbCl_{3} SbCl_{5} Sb_{2}O_{3} Sb_{2}O_{5} .... BiCl_{3} BiCl_{5} Bi_{2}O_{3} Bi_{2}O_{5}

In the case of phosphorus, arsenic, and antimony the oxides are acid anhydrides. Salts of at least four acids of each of these three elements are known, the free acid in some instances being unstable. The relation of these acids to the corresponding anhydrides may be illustrated as follows, phosphorus being taken as an example:

P_{2}O_{3} + 3H_{2}O = 2H_{3}PO_{3} (phosphorous acid).

P_{2}O_{5} + 3H_{2}O = 2H_{3}PO_{4} (phosphoric acid).

P{2}O{5} + 2H{2}O = H{4}P{2}O{7} (pyrophosphoric acid).

P{2}O{5} + H{2}O = 2HPO{3} (metaphosphoric acid).

PHOSPHORUS

History. The element phosphorus was discovered by the alchemist Brand, of Hamburg, in 1669, while searching for the philosopher's stone. Owing to its peculiar properties and the secrecy which was maintained about its preparation, it remained a very rare and costly substance until the demand for it in the manufacture of matches brought about its production on a large scale.

Occurrence. Owing to its great chemical activity phosphorus never occurs free in nature. In the form of phosphates it is very abundant and widely distributed. Phosphorite and sombrerite are mineral forms of calcium phosphate, while apatite consists of calcium phosphate together with calcium fluoride or chloride. These minerals form very large deposits and are extensively mined for use as fertilizers. Calcium phosphate is a constituent of all fertile soil, having been supplied to the soil by the disintegration of rocks containing it. It is the chief mineral constituent of bones of animals, and bone ash is therefore nearly pure calcium phosphate.

Preparation. Phosphorus is now manufactured from bone ash or a pure mineral phosphate by heating the phosphate with sand and carbon in an electric furnace. The materials are fed in at M (Fig. 70) by the feed screw F. The phosphorus vapor escapes at P and is condensed under water, while the calcium silicate is tapped off as a liquid at S. The phosphorus obtained in this way is quite impure, and is purified by distillation.



Explanation of the reaction. To understand the reaction which occurs, it must be remembered that a volatile acid anhydride is expelled from its salts when heated with an anhydride which is not volatile. Thus, when sodium carbonate and silicon dioxide are heated together the following reaction takes place:

Na{2}CO{3} + SiO{2} = Na{2}SiO{3} + CO{2}.

Silicon dioxide is a less volatile anhydride than phosphoric anhydride (P{2}O{5}), and when strongly heated with a phosphate the phosphoric anhydride is driven out, thus:

Ca_{3}(PO_{4})_{2} + 3SiO_{2} = 3CaSiO_{3} + P_{2}O_{5}.

If carbon is added before the heat is applied, the P{2}O{5} is reduced to phosphorus at the same time, according to the equation

P{2}O{5} + 5C = 2P + 5CO.

Physical properties. The purified phosphorus is a pale yellowish, translucent, waxy solid which melts at 43.3 deg. and boils at 269 deg.. It can therefore be cast into any convenient form under warm water, and is usually sold in the market in the form of sticks. It is quite soft and can be easily cut with a knife, but this must always be done while the element is covered with water, since it is extremely inflammable, and the friction of the knife blade is almost sure to set it on fire if cut in the air. It is not soluble in water, but is freely soluble in some other liquids, notably in carbon disulphide. Its density is 1.8.

Chemical properties. Exposed to the air phosphorus slowly combines with oxygen, and in so doing emits a pale light, or phosphorescence, which can be seen only in a dark place. The heat of the room may easily raise the temperature to the kindling point of phosphorus, when it burns with a sputtering flame, giving off dense fumes of oxide of phosphorus. It burns with dazzling brilliancy in oxygen, and combines directly with many other elements, especially with sulphur and the halogens. On account of its great affinity for oxygen it is always preserved under water.

Phosphorus is very poisonous, from 0.2 to 0.3 gram being a fatal dose. Ground up with flour and water or similar substances, it is often used as a poison for rats and other vermin.

Precaution. The heat of the body is sufficient to raise phosphorus above its kindling temperature, and for this reason it should always be handled with forceps and never with the bare fingers. Burns occasioned by it are very painful and slow in healing.

Red phosphorus. On standing, yellow phosphorus gradually undergoes a remarkable change, being converted into a dark red powder which has a density of 2.1. It no longer takes fire easily, neither does it dissolve in carbon disulphide. It is not poisonous and, in fact, seems to be an entirely different substance. The velocity of this change increases with rise in temperature, and the red phosphorus is therefore prepared by heating the yellow just below the boiling point (250 deg.-300 deg.). When distilled and quickly condensed the red form changes back to the yellow. This is in accordance with the general rule that when a substance capable of existing in several allotropic forms is condensed from a gas or crystallized from the liquid state, the more unstable variety forms first, and this then passes into the more stable forms.

Matches. The chief use of phosphorus is in the manufacture of matches. Common matches are made by first dipping the match sticks into some inflammable substance, such as melted paraffin, and afterward into a paste consisting of (1) phosphorus, (2) some oxidizing substance, such as manganese dioxide or potassium chlorate, and (3) a binding material, usually some kind of glue. On friction the phosphorus is ignited, the combustion being sustained by the oxidizing agent and communicated to the wood by the burning paraffin. In sulphur matches the paraffin is replaced by sulphur.

In safety matches red phosphorus, an oxidizing agent, and some gritty material such as emery is placed on the side of the box, while the match tip is provided as before with an oxidizing agent and an easily oxidized substance, usually antimony sulphide. The match cannot be ignited easily by friction, save on the prepared surface.

Compounds of phosphorus with hydrogen. Phosphorus forms several compounds with hydrogen, the best known of which is phosphine (PH{3}) analogous to ammonia (NH{3}).

Preparation of phosphine. Phosphine is usually made by heating phosphorus with a strong solution of potassium hydroxide, the reaction being a complicated one.



The experiment can be conveniently made in the apparatus shown in Fig. 71. A strong solution of potassium hydroxide together with several small bits of phosphorus are placed in the flask A, and a current of coal gas is passed into the flask through the tube B until all the air has been displaced. The gas is then turned off and the flask is heated. Phosphine is formed in small quantities and escapes through the delivery tube, the exit of which is just covered by the water in the vessel C. Each bubble of the gas as it escapes into the air takes fire, and the product of combustion (P{2}O{5}) forms beautiful small rings, which float unbroken for a considerable time in quiet air. The pure phosphine does not take fire spontaneously. When prepared as directed above, impurities are present which impart this property.

Properties. Phosphine is a gas of unpleasant odor and is exceedingly poisonous. Like ammonia it forms salts with the halogen acids. Thus we have phosphonium chloride (PH{4}Cl) analogous to ammonium chloride (NH{4}Cl). The phosphonium salts are of but little importance.

Oxides of phosphorus. Phosphorus forms two well-known oxides,—the trioxide (P{2}O{3}) and the pentoxide (P{2}O{5}), sometimes called phosphoric anhydride. When phosphorus burns in an insufficient supply of air the product is partially the trioxide; in oxygen or an excess of air the pentoxide is formed. The pentoxide is much the better known of the two. It is a snow-white, voluminous powder whose most marked property is its great attraction for water. It has no chemical action upon most gases, so that they can be very thoroughly dried by allowing them to pass through properly arranged vessels containing phosphorus pentoxide.

Acids of phosphorus. The important acids of phosphorus are the following:

H{3}PO{3} phosphorous acid. H{3}PO{4} phosphoric acid. H{4}P{2}O{7} pyrophosphoric acid. HPO{3} metaphosphoric acid.

These may be regarded as combinations of the oxides of phosphorus with water according to the equations given in the discussion of the characteristics of the family.

1. Phosphorous acid (H{3}PO{3}). Neither the acid nor its salts are at all frequently met with in chemical operations. It can be easily obtained, however, in the form of transparent crystals when phosphorus trichloride is treated with water and the resulting solution is evaporated:

PCl{3} + 3H{2}O = H{3}PO{3} + 3HCl.

Its most interesting property is its tendency to take up oxygen and pass over into phosphoric acid.

2. Orthophosphoric acid (phosphoric acid) (H{3}PO{4}). This acid can be obtained by dissolving phosphorus pentoxide in boiling water, as represented in the equation

P_{2}O_{5} + 3H_{2}O = 2H_{3}PO_{4}.

It is usually made by treating calcium phosphate with concentrated sulphuric acid. The calcium sulphate produced in the reaction is nearly insoluble, and can be filtered off, leaving the phosphoric acid in solution. Very pure acid is made by oxidizing phosphorus with nitric acid. It forms large colorless crystals which are exceedingly soluble in water. Being a tribasic acid, it forms acid as well as normal salts. Thus the following compounds of sodium are known:

NaH{2}PO{4} monosodium hydrogen phosphate. Na{2}HPO{4} disodium hydrogen phosphate. Na{3}PO{4} normal sodium phosphate.

These salts are sometimes called respectively primary, secondary, and tertiary phosphates. They may be prepared by bringing together phosphoric acid and appropriate quantities of sodium hydroxide. Phosphoric acid also forms mixed salts, that is, salts containing two different metals. The most familiar compound of this kind is microcosmic salt, which has the formula Na(NH{4})HPO{4}.

Orthophosphates. The orthophosphates form an important class of salts. The normal salts are nearly all insoluble and many of them occur in nature. The secondary phosphates are as a rule insoluble, while most of the primary salts are soluble.

3. _Pyrophosphoric acid_ (H_{4}P_{2}O_{7}). On heating orthophosphoric acid to about 225 deg. pyrophosphoric acid is formed in accordance with the following equation:

2H{3}PO{4} = H{4}P{2}O{7} + H{2}O.

It is a white crystalline solid. Its salts can be prepared by heating a secondary phosphate:

2Na{2}HPO{4} = Na{4}P{2}O{7} + H{2}O.

4. _Metaphosphoric acid (glacial phosphoric acid)_ (HPO_{3}). This acid is formed when orthophosphoric acid is heated above 400 deg.:

H{3}PO{4} = HPO{3} + H{2}O.

It is also formed when phosphorus pentoxide is treated with cold water:

P{2}O{5} + H{2}O = 2HPO{3}.

It is a white crystalline solid, and is so stable towards heat that it can be fused and even volatilized without decomposition. On cooling from the fused state it forms a glassy solid, and on this account is often called glacial phosphoric acid. It possesses the property of dissolving small quantities of metallic oxides, with the formation of compounds which, in the case of certain metals, have characteristic colors. It is therefore used in the detection of these metals.

While the secondary phosphates, on heating, give salts of pyrophosphoric acid, the primary phosphates yield salts of metaphosphoric acid. The equations representing these reactions are as follows:

2Na{2}HPO{4} = Na{4}P{3}O{7} + H{2}O,

NaH{2}PO{4} = NaPO{3} + H{2}O.

Fertilizers. When crops are produced year after year on the same field certain constituents of the soil essential to plant growth are removed, and the soil becomes impoverished and unproductive. To make the land once more fertile these constituents must be replaced. The calcium phosphate of the mineral deposits or of bone ash serves well as a material for restoring phosphorus to soils exhausted of that essential element; but a more soluble substance, which the plants can more readily assimilate, is desirable. It is better, therefore, to convert the insoluble calcium phosphate into the soluble primary phosphate before it is applied as fertilizer. It will be seen by reference to the formulas for the orthophosphates (see page 244) that in a primary phosphate only one hydrogen atom of phosphoric acid is replaced by a metal. Since the calcium atom always replaces two hydrogen atoms, it might be thought that there could be no primary calcium phosphate; but if the calcium atom replaces one hydrogen atom from each of two molecules of phosphoric acid, the salt Ca(H_{2}PO_{4})_{2} will result, and this is a primary phosphate. It can be made by treatment of the normal phosphate with the necessary amount of sulphuric acid, calcium sulphate being formed at the same time, thus:

Ca_{3}(PO_{4})_{2} + 2H_{2}SO_{4} = Ca(H_{2}PO_{4})_{2} + 2CaSO_{4}.

The resulting mixture is a powder, which is sold as a fertilizer under the name of "superphosphate of lime."

ARSENIC

Occurrence. Arsenic occurs in considerable quantities in nature as the native element, as the sulphides realgar (As{2}S{2}) and orpiment (As{2}S{3}), as oxide (As{2}O{3}), and as a constituent of many metallic sulphides, such as arsenopyrite (FeAsS).

Preparation. The element is prepared by purifying the native arsenic, or by heating the arsenopyrite in iron tubes, out of contact with air, when the reaction expressed by the following equation occurs:

FeAsS = FeS + As.

The arsenic, being volatile, condenses in chambers connected with the heated tubes. It is also made from the oxide by reduction with carbon:

2As_{2}O_{3} + 3C = 4As + 3CO_{2}.

Properties. Arsenic is a steel-gray, metallic-looking substance of density 5.73. Though resembling metals in appearance, it is quite brittle, being easily powdered in a mortar. When strongly heated it sublimes, that is, it passes into a vapor without melting, and condenses again to a crystalline solid when the vapor is cooled. Like phosphorus it can be obtained in several allotropic forms. It alloys readily with some of the metals, and finds its chief use as an alloy with lead, which is used for making shot, the alloy being harder than pure lead. When heated on charcoal with the blowpipe it is converted into an oxide which volatilizes, leaving the charcoal unstained by any oxide coating. It burns readily in chlorine gas, forming arsenic trichloride,—

As + 3Cl = AsCl_{3}.

Unlike most of its compounds, the element itself is not poisonous.

Arsine (AsH{3}). When any compound containing arsenic is brought into the presence of nascent hydrogen, arsine (AsH{3}), corresponding to phosphine and ammonia, is formed. The reaction when oxide of arsenic is so treated is

As{2}O{3} + 12H = 2AsH{3} + 3H{2}O.

Arsine is a gas with a peculiar garlic-like odor, and is intensely poisonous. A single bubble of pure gas has been known to prove fatal. It is an unstable compound, decomposing into its elements when heated to a moderate temperature. It is combustible, burning with a pale bluish-white flame to form arsenic trioxide and water when air is in excess:

2AsH{3} + 6O = As{2}O{3} + 3H{2}O.

When the supply of air is deficient water and metallic arsenic are formed:

2AsH{3} + 3O = 3H{2}O + 2As.

These reactions make the detection of even minute quantities of arsenic a very easy problem.



Marsh's test for arsenic. The method devised by Marsh for detecting arsenic is most frequently used, the apparatus being shown in Fig. 72. Hydrogen is generated in the flask A by the action of dilute sulphuric acid on zinc, is dried by passing over calcium chloride in the tube B, and after passing through the hard-glass tube C is ignited at the jet D. If a substance containing arsenic is now introduced into the generator A, the arsenic is converted into arsine by the action of the nascent hydrogen, and passes to the jet along with the hydrogen. If the tube C is strongly heated at some point near the middle, the arsine is decomposed while passing this point and the arsenic is deposited just beyond the heated point in the form of a shining, brownish-black mirror. If the tube is not heated, the arsine burns along with the hydrogen at the jet. Under these conditions a small porcelain dish crowded down into the flame is blackened by a spot of metallic arsenic, for the arsine is decomposed by the heat of the flame, and the arsenic, cooled below its kindling temperature by the cold porcelain, deposits upon it as a black spot. Antimony conducts itself in the same way as arsenic, but the antimony deposit is more sooty in appearance. The two can also be distinguished by the fact that sodium hypochlorite (NaClO) dissolves the arsenic deposit, but not that formed by antimony.

Oxides of arsenic. Arsenic forms two oxides, As{2}O{3} and As{2}O{5}, corresponding to those of phosphorus. Of these arsenious oxide, or arsenic trioxide (As{2}O{3}), is much better known, and is the substance usually called white arsenic, or merely arsenic. It is found as a mineral, but is usually obtained as a by-product in burning pyrite in the sulphuric-acid industry. The pyrite has a small amount of arsenopyrite in it, and when this is burned arsenious oxide is formed as a vapor together with sulphur dioxide:

2FeAsS + 10O = Fe_{2}O_{3} + As_{2}O_{3} + 2SO_{2}.

The arsenious oxide is condensed in appropriate chambers. It is a rather heavy substance, obtained either as a crystalline powder or as large, vitreous lumps, resembling lumps of porcelain in appearance. It is very poisonous, from 0.2 to 0.3 g. being a fatal dose. It is frequently given as a poison, since it is nearly tasteless and does not act very rapidly. This slow action is due to the fact that it is not very soluble, and hence is absorbed slowly by the system. Arsenious oxide is also used as a chemical reagent in glass making and in the dye industry.

Acids of arsenic. Like the corresponding oxides of phosphorus, the oxides of arsenic are acid anhydrides. In solution they combine with bases to form salts, corresponding to the salts of the acids of phosphorus. Thus we have salts of the following acids:

H{3}AsO{3} arsenious acid.

H{3}AsO{4} orthoarsenic acid.

H_{4}As_{2}O_{3} pyroarsenic acid.

HAsO_{3} metarsenic acid.

Several other acids of arsenic are also known. Not all of these can be obtained as free acids, since they tend to lose water and form the oxides. Thus, instead of obtaining arsenious acid (H{3}AsO{3}), the oxide As{2}O{3} is obtained:

2H_{3}AsO_{3} = As_{2}O_{3} + 3H_{2}O.

Salts of all the acids are known, however, and some of them have commercial value. Most of them are insoluble, and some of the copper salts, which are green, are used as pigments. Paris green, which has a complicated formula, is a well-known insecticide.

Antidote for arsenical poisoning. The most efficient antidote for arsenic poisoning is ferric hydroxide. It is prepared as needed, according to the equation

Fe{2}(SO{4}){3} + 3Mg(OH){2} = 2Fe(OH){3} + 3MgSO{4}.

Sulphides of arsenic. When hydrogen sulphide is passed into an acidified solution containing an arsenic compound the arsenic is precipitated as a bright yellow sulphide, thus:

2H{3}AsO{3} + 3H{2}S = As{2}S{3} + 6H{2}O,

2H{3}AsO{4} + 5H{2}S = As{2}S{5} + 8H{2}O.

In this respect arsenic resembles the metallic elements, many of which produce sulphides under similar conditions. The sulphides of arsenic, both those produced artificially and those found in nature, are used as yellow pigments.

ANTIMONY

Occurrence. Antimony occurs in nature chiefly as the sulphide (Sb{2}S{3}), called stibnite, though it is also found as oxide and as a constituent of many complex minerals.

Preparation. Antimony is prepared from the sulphide in a very simple manner. The sulphide is melted with scrap iron in a furnace, when the iron combines with the sulphur to form a slag, or liquid layer of melted iron sulphide, while the heavier liquid, antimony, settles to the bottom and is drawn off from time to time. The reaction involved is represented by the equation

Sb{2}S{3} + 3Fe = 2Sb + 3FeS.

Physical properties. Antimony is a bluish-white, metallic-looking substance whose density is 6.7. It is highly crystalline, hard, and very brittle. It has a rather low melting point (432 deg.) and expands very noticeably on solidifying.

Chemical properties. In chemical properties antimony resembles arsenic in many particulars. It forms the oxides Sb_{2}O_{3} and Sb_{2}O_{5}, and in addition Sb_{2}O_{4}. It combines with the halogen elements with great energy, burning brilliantly in chlorine to form antimony trichloride (SbCl_{3}). When heated on charcoal with the blowpipe it is oxidized and forms a coating of antimony oxide on the charcoal which has a characteristic bluish-white color.

Stibine (SbH{3}). The gas stibine (SbH{3}) is formed under conditions which are very similar to those which produce arsine, and it closely resembles the latter compound, though it is still less stable. It is very poisonous.

Acids of antimony. The oxides Sb{2}O{3} and Sb{2}O{5} are weak acid anhydrides and are capable of forming two series of acids corresponding in formulas to the acids of phosphorus and arsenic. They are much weaker, however, and are of little practical importance.

Sulphides of antimony. Antimony resembles arsenic in that hydrogen sulphide precipitates it as a sulphide when conducted into an acidified solution containing an antimony compound:

2SbCl{3} + 3H{2}S = Sb{2}S{3} + 6HCl,

2SbCl{5} + 5H{2}S = Sb{2}S{5} + 10HCl.

The two sulphides of antimony are called the trisulphide and the pentasulphide respectively. When prepared in this way they are orange-colored substances, though the mineral stibnite is black.

Metallic properties of antimony. The physical properties of the element are those of a metal, and the fact that its sulphide is precipitated by hydrogen sulphide shows that it acts like a metal in a chemical way. Many other reactions show that antimony has more of the properties of a metal than of a non-metal. The compound Sb(OH)_{3}, corresponding to arsenious acid, while able to act as a weak acid is also able to act as a weak base with strong acids. For example, when treated with concentrated hydrochloric acid antimony chloride is formed:

Sb(OH)_{3} + 3HCl = SbCl_{3} + 3H_{2}O.

A number of elements act in this same way, their hydroxides under some conditions being weak acids and under others weak bases.

ALLOYS

Some metals when melted together thoroughly intermix, and on cooling form a homogeneous, metallic-appearing substance called an alloy. Not all metals will mix in this way, and in some cases definite chemical compounds are formed and separate out as the mixture solidifies, thus destroying the uniform quality of the alloy. In general the melting point of the alloy is below the average of the melting points of its constituents, and it is often lower than any one of them.

Antimony forms alloys with many of the metals, and its chief commercial use is for such purposes. It imparts to its alloys high density, rather low melting point, and the property of expanding on solidification. Such an alloy is especially useful in type founding, where fine lines are to be reproduced on a cast. Type metal consists of antimony, lead, and tin. Babbitt metal, used for journal bearings in machinery, contains the same metals in a different proportion together with a small percentage of copper.

BISMUTH

Occurrence. Bismuth is usually found in the uncombined form in nature. It also occurs as oxide and sulphide. Most of the bismuth of commerce comes from Saxony, and from Mexico and Colorado, but it is not an abundant element.

Preparation. It is prepared by merely heating the ore containing the native bismuth and allowing the melted metal to run out into suitable vessels. Other ores are converted into oxides and reduced by heating with carbon.

Physical properties. Bismuth is a heavy, crystalline, brittle metal nearly the color of silver, but with a slightly rosy tint which distinguishes it from other metals. It melts at a low temperature (270 deg.) and has a density of 9.8. It is not acted upon by the air at ordinary temperatures.

Chemical properties. When heated with the blowpipe on charcoal, bismuth gives a coating of the oxide Bi{2}O{3}. This has a yellowish-brown color which easily distinguishes it from the oxides formed by other metals. It combines very readily with the halogen elements, powdered bismuth burning readily in chlorine. It is not very easily acted upon by hydrochloric acid, but nitric and sulphuric acids act upon it in the same way that they do upon copper.

Uses. Bismuth finds its chief use as a constituent of alloys, particularly in those of low melting point. Some of these melt in hot water. For example, Wood's metal, consisting of bismuth, lead, tin, and cadmium, melts at 60.5 deg..

Compounds of bismuth. Unlike the other elements of this group, bismuth has almost no acid properties. Its chief oxide, Bi{2}O{3}, is basic in its properties. It dissolves in strong acids and forms salts of bismuth:

Bi{2}O{3} + 6HCl = 2BiCl{3} + 3H{2}O,

Bi{2}O{3} + 6HNO{3} = 2Bi(NO{3}){3} + 3H{2}O.

The nitrate and chloride of bismuth can be obtained as well-formed colorless crystals. When treated with water the salts are decomposed in the manner explained in the following paragraph.

HYDROLYSIS

Many salts such as those of antimony and bismuth form solutions which are somewhat acid in reaction, and must therefore contain hydrogen ions. This is accounted for by the same principle suggested to explain the fact that solutions of potassium cyanide are alkaline in reaction (p. 210). Water forms an appreciable number of hydrogen and hydroxyl ions, and very weak bases such as bismuth hydroxide are dissociated to but a very slight extent. When Bi^{+} ions from bismuth chloride, which dissociates very readily, are brought in contact with the OH^{-} ions from water, the two come to the equilibrium expressed in the equation

Bi^{+} + 3OH^{-} Bi(OH)_{3}.

For every hydroxyl ion removed from the solution in this way a hydrogen ion is left free, and the solution becomes acid in reaction.

Reactions of this kind and that described under potassium cyanide are called hydrolysis.

DEFINITION: Hydrolysis is the action of water upon a salt to form an acid and a base, one of which is very slightly dissociated.

Conditions favoring hydrolysis. While hydrolysis is primarily due to the slight extent to which either the acid or the base formed is dissociated, several other factors have an influence upon the extent to which it will take place.

1. Influence of mass. Since hydrolysis is a reversible reaction, the relative masses of the reacting substances influence the point at which equilibrium will be reached. In the equilibrium

BiCl_{3} + 3H_{2}O Bi(OH)_{3} + 3HCl

the addition of more water will result in the formation of more bismuth hydroxide and hydrochloric acid. The addition of more hydrochloric acid will convert some of the bismuth hydroxide into bismuth chloride.

2. Formation of insoluble substances. When one of the products of hydrolysis is nearly insoluble in water the solution will become saturated with it as soon as a very little has been formed. All in excess of this will precipitate, and the reaction will go on until the acid set free increases sufficiently to bring about an equilibrium. Thus a considerable amount of bismuth and antimony hydroxides are precipitated when water is added to the chlorides of these elements. The greater the dilution the more hydroxide precipitates. The addition of hydrochloric acid in considerable quantity will, however, redissolve the precipitate.

Partial hydrolysis. In many cases the hydrolysis of a salt is only partial, resulting in the formation of basic salts instead of the free base. Most of these basic salts are insoluble in water, which accounts for their ready formation. Thus bismuth chloride may hydrolyze by successive steps, as shown in the equations

BiCl_{3} + H_{2}O = Bi(OH)Cl_{2} + HCl,

BiCl_{3} + 2H_{2}O = Bi(OH)_{2}Cl + 2HCl,

BiCl_{3} + 3H_{2}O = Bi(OH)_{3} + 3HCl.

The basic salt so formed may also lose water, as shown in the equation

Bi(OH){2}Cl = BiOCl + H{2}O.

The salt represented in the last equation is sometimes called bismuth oxychloride, or bismuthyl chloride. The corresponding nitrate, BiONO_{3}, is largely used in medicine under the name of subnitrate of bismuth. In these two compounds the group of atoms, BiO, acts as a univalent metallic radical and is called _bismuthyl_. Similar basic salts are formed by the hydrolysis of antimony salts.

EXERCISES

1. Name all the elements so far studied which possess allotropic forms.

2. What compounds would you expect phosphorus to form with bromine and iodine? Write the equations showing the action of water on these compounds.

3. In the preparation of phosphine, why is coal gas passed into the flask? What other gases would serve the same purpose?

4. Give the formula for the salt which phosphine forms with hydriodic acid. Give the name of the compound.

5. Could phosphoric acid be substituted for sulphuric acid in the preparation of the common acids?

6. Write the equations for the preparation of the three sodium salts of orthophosphoric acid.

7. Why does a solution of disodium hydrogen phosphate react alkaline?

8. On the supposition that bone ash is pure calcium phosphate, what weight of it would be required in the preparation of 1 kg. of phosphorus?

9. If arsenopyrite is heated in a current of air, what products are formed?

10. (a) Write equations for the complete combustion of hydrosulphuric acid, methane, and arsine. (b) In what respects are the reactions similar?

11. Write the equations for all the reactions involved in Marsh's test for arsenic.

12. Write the names and formulas for the acids of antimony.

13. Write the equations showing the hydrolysis of antimony trichloride; of bismuth nitrate.

14. In what respects does nitrogen resemble the members of the phosphorus family?



CHAPTER XXI

SILICON, TITANIUM, BORON

================================================================= SYMBOL ATOMIC DENSITY CHLORIDES OXIDES WEIGHT Silicon Si 28.4 2.35 SiCl{4} SiO{2} Titanium Ti 48.1 3.5 TiCl{4} TiO{2} Boron B 11.0 2.45 BCl{3} B{2}O{3} =================================================================

General. Each of the three elements, silicon, titanium, and boron, belongs to a separate periodic family, but they occur near together in the periodic grouping and are very similar in both physical and chemical properties. Since the other elements in their families are either so rare that they cannot be studied in detail, or are best understood in connection with other elements, it is convenient to consider these three together at this point.

The three elements are very difficult to obtain in the free state, owing to their strong attraction for other elements. They can be prepared by the action of aluminium or magnesium on their oxides and in impure state by reduction with carbon in an electric furnace. They are very hard and melt only at the highest temperatures. At ordinary temperatures they are not attacked by oxygen, but when strongly heated they burn with great brilliancy. Silicon and boron are not attacked by acids under ordinary conditions; titanium is easily dissolved by them.

SILICON

Occurrence. Next to oxygen silicon is the most abundant element. It does not occur free in nature, but its compounds are very abundant and of the greatest importance. It occurs almost entirely in combination with oxygen as silicon dioxide (SiO_{2}), often called silica, or with oxygen and various metals in the form of salts of silicic acids, or silicates. These compounds form a large fraction of the earth's crust. Most plants absorb small amounts of silica from the soil, and it is also found in minute quantities in animal organisms.

Preparation. The element is most easily prepared by reducing pure powdered quartz with magnesium powder:

SiO_{2} + 2Mg = 2MgO + Si.

Properties. As would be expected from its place in the periodic table, silicon resembles carbon in many respects. It can be obtained in several allotropic forms, corresponding to those of carbon. The crystallized form is very hard, and is inactive toward reagents. The amorphous variety has, in general, properties more similar to charcoal.

Compounds of silicon with hydrogen and the halogens. Silicon hydride (SiH_{4}) corresponds in formula to methane (CH_{4}), but its properties are more like those of phosphine (PH_{3}). It is a very inflammable gas of disagreeable odor, and, as ordinarily prepared, takes fire spontaneously on account of the presence of impurities.

Silicon combines with the elements of the chlorine family to form such compounds as SiCl{4} and SiF{4}. Of these silicon fluoride is the most familiar and interesting. As stated in the discussion of fluorine, it is formed when hydrofluoric acid acts upon silicon dioxide or a silicate. With silica the reaction is thus expressed:

SiO_{2} + 4HF = SiF_{4} + 2H_{2}O.

It is a very volatile, invisible, poisonous gas. In contact with water it is partially decomposed, as shown in the equation

SiF_{4} + 4H_{2}O = 4HF + Si(OH)_{4}.

The hydrofluoric acid so formed combines with an additional amount of silicon fluoride, forming the complex fluosilicic acid (H{2}SiF{6}), thus:

2HF + SiF_{4} = H_{2}SiF_{6}.

Silicides. As the name indicates, silicides are binary compounds consisting of silicon and some other element. They are very stable at high temperatures, and are usually made by heating the appropriate substances in an electric furnace. The most important one is carborundum, which is a silicide of carbon of the formula CSi. It is made by heating coke and sand, which is a form of silicon dioxide, in an electric furnace, the process being extensively carried on at Niagara Falls. The following equation represents the reaction

SiO_{2} + 3C = CSi + 2CO.

The substance so prepared consists of beautiful purplish-black crystals, which are very hard. Carborundum is used as an abrasive, that is, as a material for grinding and polishing very hard substances. Ferrosilicon is a silicide of iron alloyed with an excess of iron, which finds extensive use in the manufacture of certain kinds of steel.

Manufacture of carborundum. The mixture of materials is heated in a large resistance furnace for about thirty-six hours. After the reaction is completed there is left a core of graphite G. Surrounding this core is a layer of crystallized carborundum C, about 16 in. thick. Outside this is a shell of amorphous carborundum A. The remaining materials M are unchanged and are used for a new charge.



Silicon dioxide (_silica_) (SiO_{2}). This substance is found in a great variety of forms in nature, both in the amorphous and in the crystalline condition. In the form of quartz it is found in beautifully formed six-sided prisms, sometimes of great size. When pure it is perfectly transparent and colorless. Some colored varieties are given special names, as amethyst (violet), rose quartz (pale pink), smoky or milky quartz (colored and opaque). Other varieties of silicon dioxide, some of which also contain water, are chalcedony, onyx, jasper, opal, agate, and flint. Sand and sandstone are largely silicon dioxide.

Properties. As obtained by chemical processes silicon dioxide is an amorphous white powder. In the crystallized state it is very hard and has a density of 2.6. It is insoluble in water and in most chemical reagents, and requires the hottest oxyhydrogen flame for fusion. Acids, excepting hydrofluoric acid, have little action on it, and it requires the most energetic reducing agents to deprive it of oxygen. It is the anhydride of an acid, and consequently it dissolves in fused alkalis to form silicates. Being nonvolatile, it will drive out most other anhydrides when heated to a high temperature with their salts, especially when the silicates so formed are fusible. The following equations illustrate this property:

Na{2}CO{3} + SiO{2} = Na{2}SiO{3} + CO{2},

Na{2}SO{4} + SiO{2} = Na{2}SiO{3} + SO{3}.

Silicic acids. Silicon forms two simple acids, orthosilicic acid (H{4}SiO{4}) and metasilicic acid (H{2}SiO{3}). Orthosilicic acid is formed as a jelly-like mass when orthosilicates are treated with strong acids such as hydrochloric. On attempting to dry this acid it loses water, passing into metasilicic or common silicic acid:

H_{4}SiO_{4} = H_{2}SiO_{3} + H_{2}O.

Metasilicic acid when heated breaks up into silica and water, thus:

H{2}SiO{3} = H{2}O + SiO{2}.

Salts of silicic acids,—silicates. A number of salts of the orthosilicic and metasilicic acids occur in nature. Thus mica (KAlSiO_{4}) is a salt of orthosilicic acid.

Polysilicic acids. Silicon has the power to form a great many complex acids which may be regarded as derived from the union of several molecules of the orthosilicic acid, with the loss of water. Thus we have

3H{4}SiO{4} = H{4}Si{3}O{8} + 4H{2}O.

These acids cannot be prepared in the pure state, but their salts form many of the crystalline rocks in nature. Feldspar, for example, has the formula KAlSi_{3}O_{8}, and is a mixed salt of the acid H_{4}Si_{3}O_{8}, whose formation is represented in the equation above. Kaolin has the formula Al_{2}Si_{2}O_{7}.2H_{2}O. Many other examples will be met in the study of the metals.

Glass. When sodium and calcium silicates, together with silicon dioxide, are heated to a very high temperature, the mixture slowly fuses to a transparent liquid, which on cooling passes into the solid called glass. Instead of starting with sodium and calcium silicates it is more convenient and economical to heat sodium carbonate (or sulphate) and lime with an excess of clean sand, the silicates being formed during the heating:

Na{2}CO{3} + SiO{2} = Na{2}SiO{3} + CO{2},

CaO + SiO{2} = CaSiO{3}.



The mixture is heated below the fusing point for some time, so that the escaping carbon dioxide may not spatter the hot liquid; the heat is then increased and the mixture kept in a state of fusion until all gases formed in the reaction have escaped.

Molding and blowing of glass. The way in which the melted mixture is handled in the glass factory depends upon the character of the article to be made. Many articles, such as bottles, are made by blowing the plastic glass into hollow molds of the desired shape. The mold is first opened, as shown in Fig. 74. A lump of plastic glass A on the hollow rod B is lowered into the mold, which is then closed by the handles C. By blowing into the tube the glass is blown into the shape of the mold. The mold is then opened and the bottle lifted out. The neck of the bottle must be cut off at the proper place and the sharp edges rounded off in a flame.

Other objects, such as lamp chimneys, are made by getting a lump of plastic glass on the end of a hollow iron rod and blowing it into the desired shape without the help of a mold, great skill being required in the manipulation of the glass. Window glass is made by blowing large hollow cylinders about 6 ft. long and 1-1/2 ft. in diameter. These are cut longitudinally, and are then placed in an oven and heated until they soften, when they are flattened out into plates (Fig. 75). Plate glass is cast into flat slabs, which are then ground and polished to perfectly plane surfaces.

Varieties of glass. The ingredients mentioned above make a soft, easily fusible glass. If potassium carbonate is substituted for the sodium carbonate, the glass is much harder and less easily fused; increasing the amount of sand has somewhat the same effect. Potassium glass is largely used in making chemical glassware, since it resists the action of reagents better than the softer sodium glass. If lead oxide is substituted for the whole or a part of the lime, the glass is very soft, but has a high index of refraction and is valuable for making optical instruments and artificial jewels.



Coloring of glass. Various substances fused along with the glass mixture give characteristic colors. The amber color of common bottles is due to iron compounds in the glass; in other cases iron colors the glass green. Cobalt compounds color it deep blue; those of manganese give it an amethyst tint and uranium compounds impart a peculiar yellowish green color. Since iron is nearly always present in the ingredients, glass is usually slightly yellow. This color can be removed by adding the proper amount of manganese dioxide, for the amethyst color of manganese and the yellow of iron together produce white light.

Nature of glass. Glass is not a definite chemical compound and its composition varies between wide limits. Fused glass is really a solution of various silicates, such as those of calcium and lead, in fused sodium or potassium silicate. A certain amount of silicon dioxide is also present. This solution is then allowed to solidify under such conditions of cooling that the dissolved substances do not separate from the solvent. The compounds which are used to color the glass are sometimes converted into silicates, which then dissolve in the glass, giving it a uniform color. In other cases, as in the milky glasses which resemble porcelain in appearance, the color or opaqueness is due to the finely divided color material evenly distributed throughout the glass, but not dissolved in it. Milky glass is made by mixing calcium fluoride, tin oxide, or some other insoluble substance in the melted glass. Copper or gold in metallic form scattered through glass gives it shades of red.

TITANIUM

Titanium is a very widely distributed element in nature, being found in almost all soils, in many rocks, and even in plant and animal tissues. It is not very abundant in any one locality, and it possesses little commercial value save in connection with the iron industry. Its most common ore is rutile (TiO_{2}), which resembles silica in many respects.

In both physical and chemical properties titanium resembles silicon, though it is somewhat more metallic in character. This resemblance is most marked in the acids of titanium. It not only forms metatitanic and orthotitanic acids but a great variety of polytitanic acids as well.

BORON

Occurrence. Boron is never found free in nature. It occurs as boric acid (H{3}BO{3}), and in salts of polyboric acids, which usually have very complicated formulas.

Preparation and properties. Boron can be prepared from its oxide by reduction with magnesium, exactly as in the case of silicon. It resembles silicon very strikingly in its properties. It occurs in several allotropic forms, is very hard when crystallized, and is rather inactive toward reagents. It forms a hydride, BH{3}, and combines directly with the elements of the chlorine family. Boron fluoride (BF{3}) is very similar to silicon fluoride in its mode of formation and chemical properties.

Boric oxide (B{2}O{3}). Boron forms one well-known oxide, B{2}O{3}, called boric anhydride. It is formed as a glassy mass by heating boric acid to a high temperature. It absorbs water very readily, uniting with it to form boric acid again:

B_{2}O_{3} + 3H_{2}O = 2H_{3}BO_{3}.

In this respect it differs from silicon dioxide, which will not combine directly with water.

Boric acid (H{3}BO{3}). This is found in nature in considerable quantities and forms one of the chief sources of boron compounds. It is found dissolved in the water of hot springs in some localities, particularly in Italy. Being volatile with steam, the vapor which escapes from these springs has some boric acid in it. It is easily obtained from these sources by condensation and evaporation, the necessary heat being supplied by other hot springs.

Boric acid crystallizes in pearly flakes, which are greasy to the touch. In the laboratory it is easily prepared by treating a strong, hot solution of borax with sulphuric acid. Boric acid being sparingly soluble in water crystallizes out on cooling:

Na{2}B{4}O{7} + 5H{2}O + H{2}SO{4} = Na{2}SO{4} + 4H{3}BO{3}.

The substance is a mild antiseptic, and on this account is often used in medicine and as a preservative for canned foods and milk.

Metaboric and polyboric acids. When boric acid is gently heated it is converted into metaboric acid (HBO_{2}):

H{3}BO{3} = HBO{2} + H{2}O.

On heating metaboric acid to a somewhat higher temperature tetraboric acid (H_{2}B_{4}O_{7}) is formed:

4HBO_{2} = H_{2}B_{4}O_{7} + H_{2}O.

Many other complex acids of boron are known.

Borax. Borax is the sodium salt of tetraboric acid, having the formula Na{2}B{4}O{7}.10 H{2}O. It is found in some arid countries, as southern California and Tibet, but is now made commercially from the mineral colemanite, which is the calcium salt of a complex boric acid. When this is treated with a solution of sodium carbonate, calcium carbonate is precipitated and borax crystallizes from the solution.

When heated borax at first swells up greatly, owing to the expulsion of the water of crystallization, and then melts to a clear glass. This glass has the property of easily dissolving many metallic oxides, and on this account borax is used as a flux in soldering, for the purpose of removing from the metallic surfaces to be soldered the film of oxide with which they are likely to be covered. These oxides often give a characteristic color to the clear borax glass, and borax beads are therefore often used in testing for the presence of metals, instead of the metaphosphoric acid bead already described.

The reason that metallic oxides dissolve in borax is that borax contains an excess of acid anhydride, as can be more easily seen if its formula is written 2NaBO_{2} + B_{2}O_{3}. The metallic oxide combines with this excess of acid anhydride, forming a mixed salt of metaboric acid.