Comparison of sulphur and oxygen. A comparison of the formulas and the chemical properties of corresponding compounds of oxygen and sulphur brings to light many striking similarities. The conduct of hydrosulphuric acid and water toward many substances has been seen to be very similar; the oxides and sulphides of the metals have analogous formulas and undergo many parallel reactions. Carbon dioxide and disulphide are prepared in similar ways and undergo many analogous reactions. It is clear, therefore, that these two elements are far more closely related to each other than to any of the other elements so far studied.

Selenium and tellurium. These two very uncommon elements are still more closely related to sulphur than is oxygen. They occur in comparatively small quantities and are usually found associated with sulphur and sulphides, either as the free elements or more commonly in combination with metals. They form compounds with hydrogen of the formulas H_{2}Se and H_{2}Te; these bodies are gases with properties very similar to those of H_{2}S. They also form oxides and oxygen acids which resemble the corresponding sulphur compounds. The elements even have allotropic forms corresponding very closely to those of sulphur. Tellurium is sometimes found in combination with gold and copper, and occasions some difficulties in the refining of these metals. The elements have very few practical applications.

Crystallography. In order to understand the difference between the two kinds of sulphur crystals, it is necessary to know something about crystals in general and the forms which they may assume. An examination of a large number of crystals has shown that although they may differ much in geometric form, they can all be considered as modifications of a few simple plans. The best way to understand the relation of one crystal to another is to look upon every crystal as having its faces and angles arranged in definite fashion about certain imaginary lines drawn through the crystal. These lines are called axes, and bear much the same relation to a crystal as do the axis and parallels of latitude and longitude to the earth and a geographical study of it. All crystals can be referred to one of six simple plans or systems, which have their axes as shown in the following drawings.

The names and characteristics of these systems are as follows:

1. Isometric or regular system (Fig. 46). Three equal axes, all at right angles.



2. Tetragonal system (Fig. 47). Two equal axes and one of different length, all at right angles to each other.



3. Orthorhombic system (Fig. 48). Three unequal axes, all at right angles to each other.



4. Monoclinic system (Fig. 49). Two axes at right angles, and a third at right angles to one of these, but inclined to the other.



5. Triclinic system (Fig. 50). Three axes, all inclined to each other.



6. Hexagonal system (Fig. 51). Three equal axes in the same plane intersecting at angles of 60 deg., and a fourth at right angles to all of these.



Every crystal can be imagined to have its faces and angles arranged in a definite way around one of these systems of axes. A cube, for instance, is referred to Plan 1, an axis ending in the center of each face; while in a regular octohedron an axis ends in each solid angle. These forms are shown in Fig. 46. It will be seen that both of these figures belong to the same system, though they are very different in appearance. In the same way, many geometric forms may be derived from each of the systems, and the light lines about the axes in the drawings show two of the simplest forms of each of the systems.

In general a given substance always crystallizes in the same system, and two corresponding faces of each crystal of it always make the same angle with each other. A few substances, of which sulphur is an example, crystallize in two different systems, and the crystals differ in such physical properties as melting point and density. Such substances are said to be dimorphous.

EXERCISES

1. (a) Would the same amount of heat be generated by the combustion of 1 g. of each of the allotropic modifications of sulphur? (b) Would the same amount of sulphur dioxide be formed in each case?

2. Is the equation for the preparation of hydrosulphuric acid a reversible one? As ordinarily carried out, does the reaction complete itself?

3. Suppose that hydrosulphuric acid were a liquid, would it be necessary to modify the method of preparation?

4. Can sulphuric acid be used to dry hydrosulphuric acid? Give reason for answer.

5. Does dry hydrosulphuric acid react with litmus paper? State reason for answer.

6. How many grams of iron sulphide are necessary to prepare 100 l. of hydrosulphuric acid when the laboratory conditions are 17 deg. and 740 mm. pressure?

7. Suppose that the hydrogen in 1 l. of hydrosulphuric acid were liberated; what volume would it occupy, the gases being measured under the same conditions?

8. Write the equations representing the reaction between hydrosulphuric acid and sodium hydroxide and ammonium hydroxide respectively.

9. Show that the preparation of sulphur dioxide from a sulphite is similar in principle to the preparation of hydrogen sulphide.

10. (a) Does dry sulphur dioxide react with litmus paper? (b) How can it be shown that a solution of sulphur dioxide in water acts like an acid?

11. (a) Calculate the percentage composition of sulphurous anhydride and sulphuric anhydride. (b) Show how these two substances are in harmony with the law of multiple proportion.

12. How many pounds of sulphur would be necessary in the preparation of 100 lb. of 98% sulphuric acid?

13. What weight of sulphur dioxide is necessary in the preparation of 1 kg. of sodium sulphite?

14. What weight of copper sulphate crystals can be obtained by dissolving 1 kg. of copper in sulphuric acid and crystallizing the product from water?

15. Write the names and formulas of the oxides and oxygen acids of selenium and tellurium.

16. In the commercial preparation of carbon disulphide, what is the function of the electric current?

17. If the Gay-Lussac tower were omitted from the sulphuric acid factory, what effect would this have on the cost of production of sulphuric acid?



CHAPTER XV

PERIODIC LAW

A number of the elements have now been studied somewhat closely. The first three of these, oxygen, hydrogen, and nitrogen, while having some physical properties in common with each other, have almost no point of similarity as regards their chemical conduct. On the other hand, oxygen and sulphur, while quite different physically, have much in common in their chemical properties.

About eighty elements are now known. If all of these should have properties as diverse as do oxygen, hydrogen, and nitrogen, the study of chemistry would plainly be a very difficult and complicated one. If, however, the elements can be classified in groups, the members of which have very similar properties, the study will be very much simplified.

Earlier classification of the elements. Even at an early period efforts were made to discover some natural principle in accordance with which the elements could be classified. Two of these classifications may be mentioned here.

1. Classification into metals and non-metals. The classification into metals and non-metals most naturally suggested itself. This grouping was based largely on physical properties, the metals being heavy, lustrous, malleable, ductile, and good conductors of heat and electricity. Elements possessing these properties are usually base-forming in character, and the ability to form bases came to be regarded as a characteristic property of the metals. The non-metals possessed physical properties which were the reverse of those of the metals, and were acid-forming in character.

Not much was gained by this classification, and it was very imperfect. Some metals, such as potassium, are very light; some non-metals, such as iodine, have a high luster; some elements can form either an acid or a base.

2. Classification into triad families. In 1825 Doebereiner observed that an interesting relation exists between the atomic weights of chemically similar elements. To illustrate, lithium, sodium, and potassium resemble each other very closely, and the atomic weight of sodium is almost exactly an arithmetical mean between those of the other two: (7.03 + 39.15)/2 = 23.09. In many chemical and physical properties sodium is midway between the other two.

A number of triad families were found, but among eighty elements, whose atomic weights range all the way from 1 to 240, such agreements might be mere chance. Moreover many elements did not appear to belong to such families.

Periodic division. In 1869 the Russian chemist Mendeleeff devised an arrangement of the elements based on their atomic weights, which has proved to be of great service in the comparative study of the elements. A few months later the German, Lothar Meyer, independently suggested the same ideas. This arrangement brought to light a great generalization, now known as the periodic law. An exact statement of the law will be given after the method of arranging the elements has been described.



Arrangement of the periodic table. The arrangement suggested by Mendeleeff, modified somewhat by more recent investigations, is as follows: Beginning with lithium, which has an atomic weight of 7, the elements are arranged in a horizontal row in the order of their atomic weights, thus:

Li (7.03), Be (9.1), B (11), C (12), N (14.04), O (16), F (19).

These seven elements all differ markedly from each other. The eighth element, sodium, is very similar to lithium. It is placed just under lithium, and a new row follows:

Na(23.05), Mg (24.36), Al (27.1), Si (28.4), P (31), S (32.06), Cl(35.45).

When the fifteenth element, potassium, is reached, it is placed under sodium, to which it is very similar, and serves to begin a third row:

K (39.15), Ca (40.1), Sc (44.1,) Ti (48.1), V (51.2), Cr (52.1), Mn(55).

Not only is there a strong similarity between lithium, sodium, and potassium, which have been placed in a vertical row because of this resemblance, but the elements in the other vertical rows exhibit much of the same kind of similarity among themselves, and evidently form little natural groups.

The three elements following manganese, namely, iron, nickel, and cobalt, have atomic weights near together, and are very similar chemically. They do not strongly resemble any of the elements so far considered, and are accordingly placed in a group by themselves, following manganese. A new row is begun with copper, which somewhat resembles the elements of the first vertical column. Following the fifth and seventh rows are groups of three closely related elements, so that the completed arrangement has the appearance represented in the table on page 168.

THE PERIODIC ARRANGEMENT OF THE ELEMENTS

- - - - - Periods GROUP GROUP GROUP GROUP GROUP 0 I II III IV A B A B A B A B A B - - - - - 1 H==1.008 2 He=4 Li=7.03 Be=9.1 B=11 C=12 - - - - - 3 Ne=20 Na=23.05 Mg=24.36 AL=27.1 Si=28.4 - - - - - 4 A=39.9 K=39.15 Ca=40.1 Sc=44.1 Ti=48.1 - - - - - 5 Cu=63.6 Zn=65.4 Ga=70 Ge=72.5 - - - - - 6 Kr=81.8 Rb=85.5 Sr=87.6 Y=89 Zr=90.6 - - - - - 7 Ag=107.93 Cd=112.4 In=115 Sn=119 - - - - - 8 X=128 Cs=132.9 Ba=137.4 La=138.9 Ce=Yb* 140.25-173 - - - - - 9 Au=197.2 Hg=200 Tl=204.1 Pb=206.9 Bi=208.5 - - - - - 10 Ra=225 Th=232.5 - - - - - R_{2}O RO R_{2}O_{3} RO_{2} RH RH_{2} RH_{3} RH_{4} - - - - -

================part 2============

- - - -+ Periods GROUP GROUP GROUP GROUP V VI VII VIII A B A B A B + - - - - 1 2 N=14.04 O=16 F=19 - - - -+ 3 P=31 S=32.06 Cl=35.45 + - - - - 4 V=51.2 Cr=52.1 Mn=55 Fe=55.9 Ni=58.7 Co=59 - - - -+ 5 As=75 Se=79.2 Br=79.96 + - - - - 6 Cb=94 Mo=96 Ru=101.7 Rh=103 Pd=106.5 - - - -+ 7 Sb=120.2 Te=127.6 I=126.97 + - - - - 8 Ta=183 W=184 Os=191 Ir=193 Pt=194.8 - - - -+ 9 + - - - - 10 U=238.5 - - - -+ R{2}O{5} RO{3} R{2}O{7} RO{4} RH{3} RH{2} RH + - - - -

[* This includes a number of elements whose atomic weights lie between 140 and 173, but which have not been accurately studied, and so their proper arrangement is uncertain.]

Place of the atmospheric elements. When argon was discovered it was seen at once that there was no place in the table for an element of atomic weight approximately 40. When the other inactive elements were found, however, it became apparent that they form a group just preceding Group 1. They are accordingly arranged in this way in Group 0 (see table on opposite page). A study of this table brings to light certain very striking facts.

Properties of elements vary with atomic weights. There is evidently a close relation between the properties of an element and its atomic weight. Lithium, at the beginning of the first group, is a very strong base-forming element, with pronounced metallic properties. Beryllium, following lithium, is less strongly base-forming, while boron has some base-forming and some acid-forming properties. In carbon all base-forming properties have disappeared, and the acid-forming properties are more marked than in boron. These become still more emphasized as we pass through nitrogen and oxygen, until on reaching fluorine we have one of the strongest acid-forming elements. The properties of these seven elements therefore vary regularly with their atomic weights, or, in mathematical language, are regular functions of them.

Periodic law. The properties of the first seven elements vary continuously—that is steadily—away from base-forming and toward acid-forming properties. If lithium had the smallest atomic weight of any of the elements, and fluorine the greatest, so that in passing from one to the other we had included all the elements, we could say that the properties of elements are continuous functions of their atomic weights. But fluorine is an element of small atomic weight, and the one following it, sodium, breaks the regular order, for in it reappear all the characteristic properties of lithium. Magnesium, following sodium, bears much the same relation to beryllium that sodium does to lithium, and the properties of the elements in the second row vary much as they do in the first row until potassium is reached, when another repetition begins. The properties of the elements do not vary continuously, therefore, with atomic weights, but at regular intervals there is a repetition, or period. This generalization is known as the periodic law, and may be stated thus: The properties of elements are periodic functions of their atomic weights.

The two families in a group. While all the elements in a given vertical column bear a general resemblance to each other, it has been noticed that those belonging to periods having even numbers are very strikingly similar to each other. They are placed at the left side of the group columns. In like manner, the elements belonging to the odd periods are very similar and are arranged at the right side of the group columns. Thus calcium, strontium, and barium are very much alike; so, too, are magnesium, zinc, and cadmium. The resemblance between calcium and magnesium, or strontium and zinc, is much less marked. This method of arrangement therefore divides each group into two families, each containing four or five members, between which there is a great similarity.

Family resemblances. Let us now inquire more closely in what respects the elements of a family resemble each other.

1. Valence. In general the valence of the elements in a family is the same, and the formulas of their compounds are therefore similar. If we know that the formula of sodium chloride is NaCl, it is pretty certain that the formula of potassium chloride will be KCl—not KCl{2} or KCl{3}. The general formulas R{2}O, RO, etc., placed below the columns show the formulas of the oxides of the elements in the column provided they form oxides. In like manner the formulas RH, RH{2}, etc., show the composition of the compounds formed with hydrogen or chlorine.

2. Chemical properties. The chemical properties of the members of a family are quite similar. If one member is a metal, the others usually are; if one is a non-metal, so, too, are the others. The families in the first two columns consist of metals, while the elements found in the last two columns form acids. There is in addition a certain regularity in properties of the elements in each family. If the element at the head of the family is a strong acid-forming element, this property is likely to diminish gradually, as we pass to the members of the family with higher atomic weights. Thus phosphorus is strongly acid-forming, arsenic less so, antimony still less so, while bismuth has almost no acid-forming properties. We shall meet with many illustrations of this fact.

3. Physical properties. In the same way, the physical properties of the members of a family are in general somewhat similar, and show a regular gradation as we pass from element to element in the family. Thus the densities of the members of the magnesium family are

Mg = 1.75, Zn = 7.00, Cd = 8.67, Hg = 13.6.

Their melting points are

Mg = 750 deg., Zn = 420 deg., Cd = 320 deg., Hg = -39.5 deg..

Value of the periodic law. The periodic law has proved of much value in the development of the science of chemistry.

1. It simplifies study. It is at once evident that such regularities very much simplify the study of chemistry. A thorough study of one element of a family makes the study of the other members a much easier task, since so many of the properties and chemical reactions of the elements are similar. Thus, having studied the element sulphur in some detail, it is not necessary to study selenium and tellurium so closely, for most of their properties can be predicted from the relation which they sustain to sulphur.

2. It predicts new elements. When the periodic law was first formulated there were a number of vacant places in the table which evidently belonged to elements at that time unknown. From their position in the table, Mendeleeff predicted with great precision the properties of the elements which he felt sure would one day be discovered to fill these places. Three of them, scandium, germanium, and gallium, were found within fifteen years, and their properties agreed in a remarkable way with the predictions of Mendeleeff. There are still some vacant places in the table, especially among the heavier elements.

3. It corrects errors. The physical constants of many of the elements did not at first agree with those demanded by the periodic law, and a further study of many such cases showed that errors had been made. The law has therefore done much service in indicating probable error.

Imperfections of the law. There still remain a good many features which must be regarded as imperfections in the law. Most conspicuous is the fact that the element hydrogen has no place in the table. In some of the groups elements appear in one of the families, while all of their properties show that they belong in the other. Thus sodium belongs with lithium and not with copper; fluorine belongs with chlorine and not with manganese. There are two instances where the elements must be transposed in order to make them fit into their proper group. According to their atomic weights, tellurium should follow iodine, and argon should follow potassium. Their properties show in each case that this order must be reversed. The table separates some elements altogether which, in many respects have closely agreeing properties. Iron, chromium, and manganese are all in different groups, although they are similar in many respects.

The system is therefore to be regarded as but a partial and imperfect expression of some very important and fundamental relation between the substances which we know as elements, the exact nature of this relation being as yet not completely clear to us.

EXERCISES

1. Suppose that an element were discovered that filled the blank in Group O, Period 5; what properties would it probably have?

2. Suppose that an element were discovered that filled the blank in Group VI, Period 9, family B; what properties would it have?

3. Sulphur and oxygen both belong in Group VI, although in different families; in what respects are the two similar?



CHAPTER XVI

THE CHLORINE FAMILY

================================================================== ATOMIC MELTING BOILING COLOR AND STATE WEIGHT POINT POINT ___ __ __ __ ____ Fluorine (F) 19.00 -223 deg. -187 deg. Pale yellowish gas. Chlorine (Cl) 35.45 -102 deg. -33.6 deg. Greenish-yellow gas. Bromine (Br) 79.96 -7 deg. 59 deg. Red liquid. Iodine (I) 126.97 107 deg. 175 deg. Purplish-black solid. ==================================================================

The family. The four elements named in the above table form a strongly marked family of elements and illustrate very clearly the way in which the members of a family in a periodic group resemble each other, as well as the character of the differences which we may expect to find between the individual members.

1. Occurrence. These elements do not occur in nature in the free state. The compounds of the last three elements of the family are found extensively in sea water, and on this account the name halogens, signifying "producers of sea salt," is sometimes applied to the family.

2. Properties. As will be seen by reference to the table, the melting points and boiling points of the elements of the family increase with their atomic weights. A somewhat similar gradation is noted in their color and state. One atom of each of the elements combines with one atom of hydrogen to form acids, which are gases very soluble in water. The affinity of the elements for hydrogen is in the inverse order of their atomic weights, fluorine having the strongest affinity and iodine the weakest. Only chlorine and iodine form oxides, and those of the former element are very unstable. The elements of the group are univalent in their compounds with hydrogen and the metals.

FLUORINE

Occurrence. The element fluorine occurs in nature most abundantly as the mineral fluorspar (CaF_{2}), as cryolite (Na_{3}AlF_{6}), and in the complex mineral apatite (3 Ca_{3}(PO_{4})_{2}.CaF_{2}).

Preparation. All attempts to isolate the element resulted in failure until recent years. Methods similar to those which succeed in the preparation of the other elements of the family cannot be used; for as soon as the fluorine is liberated it combines with the materials of which the apparatus is made or with the hydrogen of the water which is always present. The preparation of fluorine was finally accomplished by the French chemist Moissan by the electrolysis of hydrofluoric acid. Perfectly dry hydrofluoric acid (HF) was condensed to a liquid and placed in a U-shaped tube made of platinum (or copper), which was furnished with electrodes and delivery tubes, as shown in Fig. 52. This liquid is not an electrolyte, but becomes such when potassium fluoride is dissolved in it. When this solution was electrolyzed hydrogen was set free at the cathode and fluorine at the anode.



Properties. Fluorine is a gas of slightly yellowish color, and can be condensed to a liquid boiling at -187 deg. under atmospheric pressure. It solidifies at -223 deg.. It is extremely active chemically, being the most active of all the elements at ordinary temperatures.

It combines with all the common elements save oxygen, very often with incandescence and the liberation of much heat. It has a strong affinity for hydrogen and is able to withdraw it from its compounds with other elements. Because of its great activity it is extremely poisonous. Fluorine does not form any oxides, neither does it form any oxygen acids, in which respects it differs from the other members of the family.

Hydrofluoric acid (HF). Hydrofluoric acid is readily obtained from fluorspar by the action of concentrated sulphuric acid. The equation is

CaF{2} + H{2}SO{4} = CaSO{4} + 2HF.

In its physical properties it resembles the binary acids of the other elements of this family, being, however, more easily condensed to a liquid. The anhydrous acid boils at 19 deg. and can therefore be prepared at ordinary pressures. It is soluble in all proportions in water, and a concentrated solution—about 50%—is prepared for the market. Its fumes are exceedingly irritating to the respiratory organs, and several chemists have lost their lives by accidentally breathing them.



Chemical properties. Hydrofluoric acid, like other strong acids, readily acts on bases and metallic oxides and forms the corresponding fluorides. It also dissolves certain metals such as silver and copper. It acts very vigorously upon organic matter, a single drop of the concentrated acid making a sore on the skin which is very painful and slow in healing. Its most characteristic property is its action upon silicon dioxide (SiO{2}), with which it forms water and the gas silicon tetrafluoride (SiF{4}), as shown in the equation

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

Glass consists of certain compounds of silicon, which are likewise acted on by the acid so that it cannot be kept in glass bottles. It is preserved in flasks made of wax or gutta-percha.

Etching. Advantage is taken of this reaction in etching designs upon glass. The glass vessel is painted over with a protective paint upon which the acid will not act, the parts which it is desired to make opaque being left unprotected. A mixture of fluorspar and sulphuric acid is then painted over the vessel and after a few minutes the vessel is washed clean. Wherever the hydrofluoric acid comes in contact with the glass it acts upon it, destroying its luster and making it opaque, so that the exposed design will be etched upon the clear glass. Frosted glass globes are often made in this way.

The etching may also be effected by covering the glass with a thin layer of paraffin, cutting the design through the wax and then exposing the glass to the fumes of the acid.

Salts of hydrofluoric acid,—fluorides. A number of the fluorides are known, but only one of them, calcium fluoride (CaF_{2}), is of importance. This is the well-known mineral fluorspar.

CHLORINE

Historical. While studying the action of hydrochloric acid upon the mineral pyrolusite, in 1774, Scheele obtained a yellowish, gaseous substance to which he gave a name in keeping with the phlogiston theory then current. Later it was supposed to be a compound containing oxygen. In 1810, however, the English chemist Sir Humphry Davy proved it to be an element and named it chlorine.

Occurrence. Chlorine does not occur free in nature, but its compounds are widely distributed. For the most part it occurs in combination with the metals in the form of chlorides, those of sodium, potassium, and magnesium being most abundant. Nearly all salt water contains these substances, particularly sodium chloride, and very large salt beds consisting of chlorides are found in many parts of the world.

Preparation. Two general methods of preparing chlorine may be mentioned, namely, the laboratory method and the electrolytic method.

1. _Laboratory method._ In the laboratory chlorine is made by warming the mineral pyrolusite (manganese dioxide, MnO_{2}) with concentrated hydrochloric acid. The first reaction, which seems to be similar to the action of acids upon oxides in general, is expressed in the equation

MnO_{2} + 4HCl = MnCl_{4} + 2H_{2}O.

The manganese compound so formed is very unstable, however, and breaks clown according to the equation

MnCl{4} = MnCl{2} + 2Cl.

Instead of using hydrochloric acid in the preparation of chlorine it will serve just as well to use a mixture of sodium chloride and sulphuric acid, since these two react to form hydrochloric acid. The following equations will then express the changes:

(1) 2NaCl + H{2}SO{4} = Na{2}SO{4} + 2HCl.

(2) MnO_{2} + 4 HCl = MnCl_{2} + 2Cl + 2H_{2}O.

(3) MnCl{2} + H{2}SO{4} = MnSO{4} + 2HCl.

Combining these equations, the following equation expressing the complete reaction is obtained:

2NaCl + MnO_{2} + 2H_{2}SO_{4} = MnSO_{4} + Na_{2}SO_{4} + 2H_{2}O + 2Cl.

Since the hydrochloric acid liberated in the third equation is free to act upon manganese dioxide, it will be seen that all of the chlorine originally present in the sodium chloride is set free.

The manganese dioxide and the hydrochloric acid are brought together in a flask, as represented in Fig. 53, and a gentle heat is applied. The rate of evolution of the gas is regulated by the amount of heat applied, and the gas is collected by displacement of air. As the equations show, only half of the chlorine present in the hydrochloric acid is liberated.



2. Electrolytic method. Under the discussion of electrolysis (p. 102) it was shown that when a solution of sodium chloride is electrolyzed chlorine is evolved at the anode, while the sodium set free at the cathode reacts with the water to form hydrogen, which is evolved, and sodium hydroxide, which remains in solution. A great deal of the chlorine required in the chemical industries is now made in this way in connection with the manufacture of sodium hydroxide.

Physical properties. Chlorine is a greenish-yellow gas, which has a peculiar suffocating odor and produces a very violent effect upon the throat and lungs. Even when inhaled in small quantities it often produces all the symptoms of a hard cold, and in larger quantities may have serious and even fatal action. It is quite heavy (density = 2.45) and can therefore be collected by displacement of air. One volume of water under ordinary conditions dissolves about three volumes of chlorine. The gas is readily liquefied, a pressure of six atmospheres serving to liquefy it at 0 deg.. It forms a yellowish liquid which solidifies at -102 deg..

Chemical properties. At ordinary temperatures chlorine is far more active chemically than any of the elements we have so far considered, with the exception of fluorine; indeed, it is one of the most active of all elements.

1. Action on metals. A great many metals combine directly with chlorine, especially when hot. A strip of copper foil heated in a burner flame and then dropped into chlorine burns with incandescence. Sodium burns brilliantly when heated strongly in slightly moist chlorine. Gold and silver are quickly tarnished by the gas.

2. Action on non-metals. Chlorine has likewise a strong affinity for many of the non-metals. Thus phosphorus burns in a current of the gas, while antimony and arsenic in the form of a fine powder at once burst into flame when dropped into jars of the gas. The products formed in all cases where chlorine combines with another element are called chlorides.

3. Action on hydrogen. Chlorine has a strong affinity for hydrogen, uniting with it to form hydrochloric acid. A jet of hydrogen burning in the air continues to burn when introduced into a jar of chlorine, giving a somewhat luminous flame. A mixture of the two gases explodes violently when a spark is passed through it or when it is exposed to bright sunlight. In the latter case it is the light and not the heat which starts the action.

4. Action on substances containing hydrogen. Not only will chlorine combine directly with free hydrogen but it will often abstract the element from its compounds. Thus, when chlorine is passed into a solution containing hydrosulphuric acid, sulphur is precipitated and Hydrochloric acid formed. The reaction is shown by the following equation:

H_{2}S + 2Cl = 2HCl + S.

With ammonia the action is similar:

NH_{3} + 3Cl = 3HCl + N.

The same tendency is very strikingly seen in the action of chlorine upon turpentine. The latter substance is largely made up of compounds having the composition represented by the formula C{10}H{16}. When a strip of paper moistened with warm turpentine is placed in a jar of chlorine dense fumes of hydrochloric acid appear and a black deposit of carbon is formed. Even water, which is a very stable compound, can be decomposed by chlorine, the oxygen being liberated. This may be shown in the following way:



If a long tube of rather large diameter is filled with a strong solution of chlorine in water and inverted in a vessel of the same solution, as shown in Fig. 54, and the apparatus is placed in bright sunlight, very soon bubbles of a gas will be observed to rise through the solution and collect in the tube. An examination of this gas will show that it is oxygen. It is liberated from water in accordance with the following equation:

H_{2}O + 2Cl = 2HCl + O.

5. Action on color substances,—bleaching action. If strips of brightly colored cloth or some highly colored flowers are placed in quite dry chlorine, no marked change in color is noticed as a rule. If, however, the cloth and flowers are first moistened, the color rapidly disappears, that is, the objects are bleached. Evidently the moisture as well as the chlorine is concerned in the action, and a study of the case shows that the chlorine has combined with the hydrogen of the water. The oxygen set free oxidizes the color substance, converting it into a colorless compound. It is evident from this explanation that chlorine will only bleach those substances which are changed into colorless compounds by oxidation.

6. Action as a disinfectant. Chlorine has also marked germicidal properties, and the free element, as well as compounds from which it is easily liberated, are used as disinfectants.

Nascent state. It will be noticed that oxygen when set free from water by chlorine is able to do what ordinary oxygen cannot do, for both the cloth and the flowers are unchanged in the air which contains oxygen. It is generally true that the activity of an element is greatest at the instant of liberation from its compounds. To express this fact elements at the instant of liberation are said to be in the nascent state. It is nascent oxygen which does the bleaching.

Hydrochloric acid (muriatic acid) (HCl). The preparation of hydrochloric acid may be discussed under two general heads:

1. Laboratory preparation. The product formed by the burning of hydrogen in chlorine is the gas hydrochloric acid. This substance is much more easily obtained, however, by treating common salt (sodium chloride) with sulphuric acid. The following equation shows the reaction:

2NaCl + H{2}SO{4} = Na{2}SO{4} + 2HCl.

The dry salt is placed in a flask furnished with a funnel tube and an exit tube, the sulphuric acid is added, and the flask gently warmed. The hydrochloric acid gas is rapidly given off and can be collected by displacement of air. The same apparatus can be used as was employed in the preparation of chlorine (Fig. 53).

When a solution of salt is treated with sulphuric acid there is no very marked action. The hydrochloric acid formed is very soluble in water, and so does not escape from the solution; hence a state of equilibrium is soon reached between the four substances represented in the equation. When concentrated sulphuric acid, in which hydrochloric acid is not soluble, is poured upon dry salt the reaction is complete.

2. Commercial preparation. Commercially, hydrochloric acid is prepared in connection with the manufacture of sodium sulphate, the reaction being the same as that just given. The reaction is carried out in a furnace, and the hydrochloric acid as it escapes in the form of gas is passed into water in which it dissolves, the solution forming the hydrochloric acid of commerce. When the materials are pure a colorless solution is obtained. The most concentrated solution has a density of 1.2 and contains 40% HCl. The commercial acid, often called muriatic acid, is usually colored yellow by impurities.

Composition of hydrochloric acid. When a solution of hydrochloric acid is electrolyzed in an apparatus similar to the one in which water was electrolyzed (Fig. 18), chlorine collects at the anode and hydrogen at the cathode. At first the chlorine dissolves in the water, but soon the water in the one tube becomes saturated with it, and if the stopcocks are left open until this is the case, and are then closed, it will be seen that the two gases are set free in equal volumes.

When measured volumes of the two gases are caused to unite it is found that one volume of hydrogen combines with one of chlorine. Other experiments show that the volume of hydrochloric acid formed is just equal to the sum of the volumes of hydrogen and chlorine. Therefore one volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric acid gas. Since chlorine is 35.18 times as heavy as hydrogen, it follows that one part of hydrogen by weight combines with 35.18 parts of chlorine to form 36.18 parts of hydrochloric acid.

Physical properties. Hydrochloric acid is a colorless gas which has an irritating effect when inhaled, and possesses a sour, biting taste, but no marked odor. It is heavier than air (density = 1.26) and is very soluble in water. Under standard conditions 1 volume of water dissolves about 500 volumes of the gas. On warming such a solution the gas escapes, until at the boiling point the solution contains about 20% by weight of HCl. Further boiling will not drive out any more acid, but the solution will distill with unchanged concentration. A more dilute solution than this will lose water on boiling until it has reached the same concentration, 20%, and will then distill unchanged. Under high pressure the gas can be liquefied, 28 atmospheres being required at 0 deg.. Under these conditions it forms a colorless liquid which is not very active chemically. It boils at -80 deg. and solidifies at -113 deg.. The solution of the gas in water is used almost entirely in the place of the gas itself, since it is not only far more convenient but also more active.

Chemical properties. The most important chemical properties of hydrochloric acid are the following:

1. Action as an acid. In aqueous solution hydrochloric acid has very strong acid properties; indeed, it is one of the strongest acids. It acts upon oxides and hydroxides, converting them into salts:

NaOH + HCl = NaCl + H_{2}O, CuO + 2HCl = CuCl_{2} + H_{2}O.

It acts upon many metals, forming chlorides and liberating hydrogen:

Zn + 2HCl = ZnCl{2} + 2H, Al + 3HCl = AlCl{3} + 3H.

Unlike nitric and sulphuric acids it has no oxidizing action, so that when it acts on metals hydrogen is always given off.

2. Relation to combustion. Hydrochloric acid gas is not readily decomposed, and is therefore neither combustible nor a supporter of combustion.

3. Action on oxidizing agents. Although hydrochloric acid is incombustible, it can be oxidized under some circumstances, in which case the hydrogen combines with oxygen, while the chlorine is set free. Thus, when a solution of hydrochloric acid acts upon manganese dioxide part of the chlorine is set free:

MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + 2Cl.

Aqua regia. It has been seen that when nitric acid acts as an oxidizing agent it usually decomposes, as represented in the equation

2HNO{3} = H{2}O + 2NO + 3O.

The oxygen so set free may act on hydrochloric acid:

6HCl + 3O = 3H_{2}O + 6Cl.

The complete equation therefore is

2HNO{3} + 6HCl = 4H{2}O + 2NO + 6Cl.

When concentrated nitric and hydrochloric acids are mixed this reaction goes on slowly, chlorine and some other substances not represented in the equation being formed. The mixture is known as aqua regia and is commonly prepared by adding one volume of nitric acid to three volumes of hydrochloric acid. It acts more powerfully upon metals and other substances than either of the acids separately, and owes its strength not to acid properties but to the action of the nascent chlorine which it liberates. Consequently, when it acts upon metals such as gold it converts them into chlorides, and the reaction can be represented by such equations as

Au + 3Cl = AuCl_{3}.

Salts of hydrochloric acid,—chlorides. The chlorides of all the metals are known and many of them are very important compounds. Some of them are found in nature, and all can be prepared by the general method of preparing salts. Silver chloride, lead chloride, and mercurous chloride are insoluble in water and acids, and can be prepared by adding hydrochloric acid to solutions of compounds of the respective elements. While the chlorides have formulas similar to the fluorides, their properties are often quite different. This is seen in the solubility of the salts. Those metals whose chlorides are insoluble form soluble fluorides, while many of the metals which form soluble chlorides form insoluble fluorides.

Compounds of chlorine with oxygen and hydrogen. Chlorine combines with oxygen and hydrogen to form four different acids. They are all quite unstable, and most of them cannot be prepared in pure form; their salts can easily be made, however, and some of them will be met with in the study of the metals. The formulas and names of these acids are as follows:

HClO hypochlorous acid.

HClO_{2} chlorous acid.

HClO_{3} chloric acid.

HClO_{4} perchloric acid.

Oxides of chlorine. Two oxides are known, having the formulas Cl{2}O and ClO{2}. They decompose very easily and are good oxidizing agents.

BROMINE

Historical. Bromine was discovered in 1826 by the French chemist Ballard, who isolated it from sea salt. He named it bromine (stench) because of its unbearable fumes.

Occurrence. Bromine occurs almost entirely in the form of bromides, especially as sodium bromide and magnesium bromide, which are found in many salt springs and salt deposits. The Stassfurt deposits in Germany and the salt waters of Ohio and Michigan are especially rich in bromides.

Preparation of bromine. The laboratory method of preparing bromine is essentially different from the commercial method.



1. Laboratory method. As in the case of chlorine, bromine can be prepared by the action of hydrobromic acid (HBr) on manganese dioxide. Since hydrobromic acid is not an article of commerce, a mixture of sulphuric acid and a bromide is commonly substituted for it. The materials are placed in a retort arranged as shown in Fig. 55. The end of the retort just touches the surface of the water in the test tube. On heating, the bromine distills over and is collected in the cold receiver. The equation is

2NaBr + 2H_{2}SO_{4} + MnO_{2} = Na_{2}SO_{4} + MnSO_{4} + 2H_{2}O + 2Br.

2. Commercial method. Bromine is prepared commercially from the waters of salt wells which are especially rich in bromides. On passing a current of electricity through such waters the bromine is first liberated. Any chlorine liberated, however, will assist in the reaction, since free chlorine decomposes bromides, as shown in the equation

NaBr + Cl = NaCl + Br.

When the water containing the bromine is heated, the liberated bromine distills over into the receiver.

Physical properties. Bromine is a dark red liquid about three times as heavy as water. Its vapor has a very offensive odor and is most irritating to the eyes and throat. The liquid boils at 59 deg. and solidifies at -7 deg.; but even at ordinary temperatures it evaporates rapidly, forming a reddish-brown gas very similar to nitrogen peroxide in appearance. Bromine is somewhat soluble in water, 100 volumes of water under ordinary conditions dissolving 1 volume of the liquid. It is readily soluble in carbon disulphide, forming a yellow solution.

Chemical properties and uses. In chemical action bromine is very similar to chlorine. It combines directly with many of the same elements with which chlorine unites, but with less energy. It combines with hydrogen and takes away the latter element from some of its compounds, but not so readily as does chlorine. Its bleaching properties are also less marked.

Bromine finds many uses in the manufacture of organic drugs and dyestuffs and in the preparation of bromides.

Hydrobromic acid (HBr). When sulphuric acid acts upon a bromide hydrobromic acid is set free:

2NaBr + H{2}SO{4} = Na{2}SO{4} + 2HBr.

At the same time some bromine is set free, as may be seen from the red fumes which appear, and from the odor. The explanation of this is found in the fact that hydrobromic acid is much less stable than hydrochloric acid, and is therefore more easily oxidized. Concentrated sulphuric acid is a good oxidizing agent, and oxidizes a part of the hydrobromic acid, liberating bromine:

H{2}SO{4} + 2HBr = 2H{2}O + SO{2} + 2Br.

Preparation of pure hydrobromic acid. A convenient way to make pure hydrobromic acid is by the action of bromine upon moist red phosphorus. This can be done with the apparatus shown in Fig. 56. Bromine is put into the dropping funnel A, and red phosphorus, together with enough water to cover it, is placed in the flask B. By means of the stopcock the bromine is allowed to flow drop by drop into the flask, the reaction taking place without the application of heat. The equations are

(1) P + 3Br = PBr_{3},

(2) PBr_{3} + 3H_{2}O = P(OH)_{3} + 3HBr.



The U-tube C contains glass beads which have been moistened with water and rubbed in red phosphorus. Any bromine escaping action in the flask acts upon the phosphorus in the U-tube. The hydrobromic acid is collected in the same way as hydrochloric acid.

Properties. Hydrobromic acid very strikingly resembles hydrochloric acid in physical and chemical properties. It is a colorless, strongly fuming gas, heavier than hydrochloric acid and, like it, is very soluble in water. Under standard conditions 1 volume of water dissolves 610 volumes of the gas. Chemically, the chief point in which it differs from hydrochloric acid is in the fact that it is much more easily oxidized, so that bromine is more readily set free from it than chlorine is from hydrochloric acid.

Salts of hydrobromic acid,—bromides. The bromides are very similar to the chlorides in their properties. Chlorine acts upon both bromides and free hydrobromic acid, liberating bromine from them:

KBr + Cl = KCl + Br,

HBr + Cl = HCl + Br.

Silver bromide is extensively used in photography, and the bromides of sodium and potassium are used as drugs.

Oxygen compounds. No oxides of bromine are surely known, and bromine does not form so many oxygen acids as chlorine does. Salts of hypobromous acid (HBrO) and bromic acid (HBrO_{3}) are known.

IODINE

Historical. Iodine was discovered in 1812 by Courtois in the ashes of certain sea plants. Its presence was revealed by its beautiful violet vapor, and this suggested the name iodine (from the Greek for violet appearance).

Occurrence. In the combined state iodine occurs in very small quantities in sea water, from which it is absorbed by certain sea plants, so that it is found in their ashes. It occurs along with bromine in salt springs and beds, and is also found in Chili saltpeter.

Preparation. Iodine may be prepared in a number of ways, the principal methods being the following:

1. Laboratory method. Iodine can readily be prepared in the laboratory from an iodide by the method used in preparing bromine, except that sodium iodide is substituted for sodium bromide. It can also be made by passing chlorine into a solution of an iodide.



2. _Commercial method._ Commercially iodine was formerly prepared from seaweed (kelp), but is now obtained almost entirely from the deposits of Chili saltpeter. The crude saltpeter is dissolved in water and the solution evaporated until the saltpeter crystallizes. The remaining liquors, known as the "mother liquors," contain sodium iodate (NaIO_{3}), in which form the iodine is present in the saltpeter. The chemical reaction by which the iodine is liberated from this compound is a complicated one, depending on the fact that sulphurous acid acts upon iodic acid, setting iodine free. This reaction is shown as follows:

2HIO{3} + 5H{2}SO{3} = 5H{2}SO{4} + H{2}O + 2I.

Purification of iodine. Iodine can be purified very conveniently in the following way. The crude iodine is placed in an evaporating dish E (Fig. 57), and the dish is set upon the sand bath S. The iodine is covered with the inverted funnel F, and the sand bath is gently heated with a Bunsen burner. As the dish becomes warm the iodine rapidly evaporates and condenses again on the cold surface of the funnel in shining crystals.

This process, in which a solid is converted into a vapor and is again condensed into a solid without passing through the liquid state, is called sublimation.

Physical properties. Iodine is a purplish-black, shining, heavy solid which crystallizes in brilliant plates. Even at ordinary temperatures it gives off a beautiful violet vapor, which increases in amount as heat is applied. It melts at 107 deg. and boils at 175 deg.. It is slightly soluble in water, but readily dissolves in alcohol, forming a brown solution (tincture of iodine), and in carbon disulphide, forming a violet solution. The element has a strong, unpleasant odor, though by no means as irritating as that of chlorine and bromine.

Chemical properties. Chemically iodine is quite similar to chlorine and bromine, but is still less active than bromine. It combines directly with many elements at ordinary temperatures. At elevated temperatures it combines with hydrogen, but the reaction is reversible and the compound formed is quite easily decomposed. Both chlorine and bromine displace it from its salts:

KI + Br = KBr + I,

KI + Cl = KCl + I.

When even minute traces of iodine are added to thin starch paste a very intense blue color develops, and this reaction forms a delicate test for iodine. Iodine is extensively used in medicine, especially in the form of a tincture. It is also largely used in the preparation of dyes and organic drugs, iodoform, a substance used as an antiseptic, has the formula CHI_{3}.

Hydriodic acid (HI). This acid cannot be prepared in pure condition by the action of sulphuric acid upon an iodide, since the hydriodic acid set free is oxidized by the sulphuric acid just as in the case of hydrobromic acid, but to a much greater extent. It can be prepared in exactly the same way as hydrobromic acid, iodine being substituted for bromine. It can also be prepared by passing hydrosulphuric acid into water in which iodine is suspended. The equation is

H_{2}S + 2I = 2HI + S.

The hydriodic acid formed in this way dissolves in the water.

Properties and uses. Hydriodic acid resembles the corresponding acids of chlorine and bromine in physical properties, being a strongly fuming, colorless gas, readily soluble in water. Under standard conditions 1 volume of water dissolves about 460 volumes of the gas. It is, however, more unstable than either hydrochloric or hydrobromic acids, and on exposure to the air it gradually decomposes in accordance with the equation

2HI + O = H_{2}O + 2I.

Owing to the slight affinity between iodine and hydrogen the acid easily gives up its hydrogen and is therefore a strong reducing agent. This is seen in its action on sulphuric acid.

The salts of hydriodic acid, the iodides, are, in general, similar to the chlorides and bromides. Potassium iodide (KI) is the most familiar of the iodides and is largely used in medicine.

Oxygen compounds. Iodine has a much greater affinity for oxygen than has either chlorine or bromine. When heated with nitric acid it forms a stable oxide (I{2}O{5}). Salts of iodic acid (HIO{3}) and periodic acid (HIO{4}) are easily prepared, and the free acids are much more stable than the corresponding acids of the other members of this family.

GAY-LUSSAC'S LAW OF VOLUMES

In the discussion of the composition of hydrochloric acid it was stated that one volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric acid. With bromine and iodine similar combining ratios hold good. These facts recall the simple volume relations already noted in the study of the composition of steam and ammonia. These relations may be represented graphically in the following way:

- H + Cl = H Cl + H Cl -

- - - H H + O = H{2}O + H{2}O - - -

- - - - H H H + N = NH{3} + NH{3} - - - -

In the early part of the past century Gay-Lussac, a distinguished French chemist, studied the volume relations of many combining gases, and concluded that similar relations always hold. His observations are summed up in the following law: When two gases combine chemically there is always a simple ratio between their volumes, and between the volume of either one of them and that of the product, provided it is a gas. By a simple ratio is meant of course the ratio of small whole numbers, as 1 : 2, 2 : 3.

EXERCISES

1. How do we account for the fact that liquid hydrofluoric acid is not an electrolyte?

2. Why does sulphuric acid liberate hydrofluoric acid from its salts?

3. In the preparation of chlorine, what advantages are there in treating manganese dioxide with a mixture of sodium chloride and sulphuric acid rather than with hydrochloric acid?

4. Why must chlorine water be kept in the dark?

5. What is the derivation of the word nascent?

6. What substances studied are used as bleaching agents? To what is the bleaching action due in each case?

7. What substances studied are used as disinfecting agents?

8. What is meant by the statement that hydrochloric acid is one of the strongest acids?

9. What is the meaning of the phrase aqua regia?

10. Cl_{2}O is the anhydride of what acid?

11. A solution of hydriodic acid on standing turns brown. How is this accounted for?

12. How can bromine vapor and nitrogen peroxide be distinguished from each other?

13. Write the equations for the reaction taking place when hydriodic acid is prepared from iodine, phosphorus, and water.

14. From their behavior toward sulphuric acid, to what class of agents do hydrobromic and hydriodic acids belong?

15. Give the derivation of the names of the elements of the chlorine family.

16. Write the names and formulas for the binary acids of the group in the order of the stability of the acids.

17. What is formed when a metal dissolves in each of the following? nitric acid; dilute sulphuric acid; concentrated sulphuric acid; hydrochloric acid; aqua regia.

18. How could you distinguish between a chloride, a bromide, and an iodide?

19. What weight of sodium chloride is necessary to prepare sufficient hydrochloric acid to saturate 1 l. of water under standard conditions?

20. On decomposition 100 l. of hydrochloric acid would yield how many liters of hydrogen and chlorine respectively, the gases being measured under the same conditions? Are your results in accord with the experimental facts?



CHAPTER XVII

CARBON AND SOME OF ITS SIMPLER COMPOUNDS

The family. Carbon stands at the head of a family of elements in the fourth group in the periodic table. The resemblances between the elements of this family, while quite marked, are not so striking as in the case of the elements of the chlorine family. With the exception of carbon, these elements are comparatively rare, and need not be taken up in detail in this chapter. Titanium will be referred to again in connection with silicon which it very closely resembles.

Occurrence. Carbon is found in nature in the uncombined state in several forms. The diamond is practically pure carbon, while graphite and coal are largely carbon, but contain small amounts of other substances. Its natural compounds are exceedingly numerous and occur as gases, liquids, and solids. Carbon dioxide is its most familiar gaseous compound. Natural gas and petroleum are largely compounds of carbon with hydrogen. The carbonates, especially calcium carbonate, constitute great strata of rocks, and are found in almost every locality. All living organisms, both plant and animal, contain a large percentage of this element, and the number of its compounds which go to make up all the vast variety of animate nature is almost limitless. Over one hundred thousand definite compounds containing carbon have been prepared. In the free state carbon occurs in three allotropic forms, two of which are crystalline and one amorphous.

Crystalline carbon. Crystalline carbon occurs in two forms,—diamond and graphite.

1. Diamond. Diamonds are found in considerable quantities in several localities, especially in South Africa, the East Indies, and Brazil. The crystals belong to the regular system, but the natural stones do not show this very clearly. When found they are usually covered with a rough coating which is removed in the process of cutting. Diamond cutting is carried on most extensively in Holland.

The density of the diamond is 3.5, and, though brittle, it is one of the hardest of substances. Black diamonds, as well as broken and imperfect stones which are valueless as gems, are used for grinding hard substances. Few chemical reagents have any action on the diamond, but when heated in oxygen or the air it blackens and burns, forming carbon dioxide.

Lavoisier first showed that carbon dioxide is formed by the combustion of the diamond; and Sir Humphry Davy in 1814 showed that this is the only product of combustion, and that the diamond is pure carbon.

The diamond as a gem. The pure diamond is perfectly transparent and colorless, but many are tinted a variety of colors by traces of foreign substances. Usually the colorless ones are the most highly prized, although in some instances the color adds to the value; thus the famous Hope diamond is a beautiful blue. Light passing through a diamond is very much refracted, and to this fact the stone owes its brilliancy and sparkle.

Artificial preparation of diamonds. Many attempts have been made to produce diamonds artificially, but for a long time these always ended in failure, graphite and not diamonds being the product obtained. The French chemist Moissan, in his extended study of chemistry at high temperatures, finally succeeded (1893) in making some small ones. He accomplished this by dissolving carbon in boiling iron and plunging the crucible containing the mixture into water, as shown in Fig. 58. Under these conditions the carbon crystallized in the iron in the form of the diamond. The diamonds were then obtained by dissolving away the iron in hydrochloric acid.



2. Graphite. This form of carbon is found in large quantities, especially in Ceylon, Siberia, and in some localities of the United States and Canada. It is a shining black substance, very soft and greasy to the touch. Its density is about 2.15. It varies somewhat in properties according to the locality in which it is found, and is more easily attacked by reagents than is the diamond. It is also manufactured by heating carbon with a small amount of iron (3%) in an electric furnace. It is used in the manufacture of lead pencils and crucibles, as a lubricant, and as a protective covering for iron in the form of a polish or a paint.

Amorphous carbon. Although there are many varieties of amorphous carbon known, they are not true allotropic modifications. They differ merely in their degree of purity, their fineness of division, and in their mode of preparation. These substances are of the greatest importance, owing to their many uses in the arts and industries. As they occur in nature, or are made artificially, they are nearly all impure carbon, the impurity depending on the particular substance in question.

1. Pure carbon. Pure amorphous carbon is best prepared by charring sugar. This is a substance consisting of carbon, hydrogen, and oxygen, the latter two elements being present in the ratio of one oxygen atom to two of hydrogen. When sugar is strongly heated the oxygen and hydrogen are driven off in the form of water and pure carbon is left behind. Prepared in this way it is a soft, lustrous, very bulky, black powder.

2. Coal and coke. Coals of various kinds were probably formed from vast accumulations of vegetable matter in former ages, which became covered over with earthy material and were thus protected from rapid decay. Under various natural agencies the organic matter was slowly changed into coal. In anthracite these changes have gone the farthest, and this variety of coal is nearly pure carbon. Soft or bituminous coals contain considerable organic matter besides carbon and mineral substances. When heated strongly out of contact with air the organic matter is decomposed and the resulting volatile matter is driven off in the form of gases and vapors, and only the mineral matter and carbon remain behind. The gaseous product is chiefly illuminating gas and the solid residue is coke. Some of the coke is found as a dense cake on the sides and roof of the retort. This is called retort carbon and is quite pure.

3. Charcoal. This is prepared from wood in the same way that coke is made from coal. When the process is carried on in retorts the products expelled by the heat are saved. Among these are many valuable substances such as wood alcohol and acetic acid. Where timber is abundant the process is carried out in a wasteful way, by merely covering piles of wood with sod and setting the wood on fire. Some wood burns and the heat from this decomposes the wood not burned, forming charcoal from it. The charcoal, of course, contains the mineral part of the wood from which it is formed.

4. Bone black. This is sometimes called animal charcoal, and is made by charring bones and animal refuse. The organic part of the materials is thus decomposed and carbon is left in a very finely divided state, scattered through the mineral part which consists largely of calcium phosphate. For some uses this mineral part is removed by treatment with hydrochloric acid and prolonged washing.

5. Lampblack. Lampblack and soot are products of imperfect combustion of oil and coal, and are deposited from a smoky flame on a cold surface. The carbon in this form is very finely divided and usually contains various oily materials.

Properties. While the various forms of carbon differ in many properties, especially in color and hardness, yet they are all odorless, tasteless solids, insoluble in water and characterized by their stability towards heat. Only in the intense heat of the electric arc does carbon volatilize, passing directly from the solid state into a vapor. Owing to this fact the inside surface of an incandescent light bulb after being used for some time becomes coated with a dark film of carbon. It is not acted on at ordinary temperatures by most reagents, but at a higher temperature it combines directly with many of the elements, forming compounds called carbides. When heated in the presence of sufficient oxygen it burns, forming carbon dioxide.

Uses of carbon. The chief use of amorphous carbon is for fuel to furnish heat and power for all the uses of civilization. An enormous quantity of carbon in the form of the purer coals, coke, and charcoal is used as a reducing agent in the manufacture of the various metals, especially in the metallurgy of iron. Most of the metals are found in nature as oxides, or in forms which can readily be converted into oxides. When these oxides are heated with carbon the oxygen is abstracted, leaving the metal. Retort carbon and coke are used to make electric light carbons and battery plates, while lampblack is used for indelible inks, printer's ink, and black varnishes. Bone black and charcoal have the property of absorbing large volumes of certain gases, as well as smaller amounts of organic matter; hence they are used in filters to remove noxious gases and objectionable colors and odors from water. Bone black is used extensively in the sugar refineries to remove coloring matter from the impure sugars.

Chemistry of carbon compounds. Carbon is remarkable for the very large number of compounds which it forms with the other elements, especially with oxygen and hydrogen. Compounds containing carbon are more numerous than all others put together, and the chemistry of these substances presents peculiarities not met with in the study of other substances. For these reasons the systematic study of carbon compounds, or of organic chemistry as it is usually called, must be deferred until the student has gained some knowledge of the chemistry of other elements. An acquaintance with a few of the most familiar carbon compounds is, however, essential for the understanding of the general principles of chemistry.

Compounds of carbon with hydrogen,—the hydrocarbons. Carbon unites with hydrogen to form a very large number of compounds called hydrocarbons. Petroleum and natural gas are essentially mixtures of a great variety of these hydrocarbons. Many others are found in living plants, and still others are produced by the decay of organic matter in the absence of air. Only two of them, methane and acetylene, will be discussed here.

Methane (_marsh gas_) (CH_{4}). This is one of the most important of these hydrocarbons, and constitutes about nine tenths of natural gas. As its name suggests, it is formed in marshes by the decay of vegetable matter under water, and bubbles of the gas are often seen to rise when the dead leaves on the bottom of pools are stirred. It also collects in mines, and, when mixed with air, is called _fire damp_ by the miners because of its great inflammability, damp being an old name for a gas. It is formed when organic matter, such as coal or wood, is heated in closed vessels, and is therefore a principal constituent of coal gas.

Preparation. Methane is prepared in the laboratory by heating sodium or calcium acetate with soda-lime. Equal weights of fused sodium acetate and soda-lime are thoroughly dried, then mixed and placed in a good-sized, hard-glass test tube fitted with a one-holed stopper and delivery tube. The mixture is gradually heated, and when the air has been displaced from the tube the gas is collected in bottles by displacement of water. Soda-lime is a mixture of sodium and calcium hydroxides. Regarding it as sodium hydroxide alone, the equation is

NaC{2}H{3}O{2} + NaOH = Na{2}CO{3} + CH{4}.

Properties. Methane is a colorless, odorless gas whose density is 0.55. It is difficult to liquefy, boiling at -155 deg. under standard pressure, and is almost insoluble in water. It burns with a pale blue flame, liberating much heat, and when mixed with oxygen is very explosive.

Davy's safety lamp. In 1815 Sir Humphry Davy invented a lamp for the use of miners, to prevent the dreadful mine explosions then common, due to methane mixed with air. The invention consisted in surrounding the upper part of the common miner's lamp with a mantle of wire gauze and the lower part with glass (Fig. 59). It has been seen that two gases will not combine until raised to their kindling temperature, and if while combining they are cooled below this point, the combination ceases. A flame will not pass through a wire gauze because the metal, being a good conductor of heat, takes away so much heat from the flame that the gases are cooled below the kindling temperature. When a lamp so protected is brought into an explosive mixture the gases inside the wire mantle burn in a series of little explosions, giving warning to the miner that the air is unsafe.



Acetylene (C_{2}H_{2}). This is a colorless gas usually having a disagreeable odor due to impurities. It is now made in large quantities from calcium carbide (CaC_{2}). This substance is formed when coal and lime are heated together in an electric furnace. When treated with water the carbide is decomposed, yielding acetylene:

CaC_{2} + 2H_{2}O = C_{2}H_{2} + Ca(OH)_{2}.

Under ordinary conditions the gas burns with a very smoky flame; in burners constructed so as to secure a large amount of oxygen it burns with a very brilliant white light, and hence is used as an illuminant.

Laboratory preparation. The gas can be prepared readily in a generator such as is shown in Fig. 60. The inner tube contains fragments of calcium carbide, while the outer one is filled with water. As long as the stopcock is closed the water cannot rise in the inner tube. When the stopcock is open the water rises, and, coming into contact with the carbide in the inner tube, generates acetylene. This escapes through the stopcock, and after the air has been expelled may be lighted as it issues from the burner.



Carbon forms two oxides, namely, carbon dioxide (CO_{2}) and carbon monoxide (CO).

Carbon dioxide (CO_{2}). Carbon dioxide is present in the air to the extent of about 3 parts in 10,000, and this apparently small amount is of fundamental importance in nature. In some localities it escapes from the earth in great quantities, and many spring waters carry large amounts of it in solution. When these highly charged spring waters reach the surface of the earth, and the pressure on them is removed, the carbon dioxide escapes with effervescence. It is a product of the oxidation of all organic matter, and is therefore formed in fires as well as in the process of decay. It is thrown off from the lungs of all animals in respiration, and is a product of many fermentation processes such as vinegar making and brewing. Combined with metallic oxides it forms vast deposits of carbonates in nature.

Preparation. In the laboratory carbon dioxide is always prepared by the action of an acid upon a carbonate, usually calcium carbonate, the apparatus shown in Fig. 39 serving the purpose very well. This reaction might be expected to produce carbonic acid, thus:

CaCO{3} + 2HCl = CaCl{2} + H{2}CO{3}.

Carbonic acid is very unstable, however, and decomposes into its anhydride, CO_{2}, and water, thus:

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

The complete reaction is represented by the equation

CaCO{3} + 2HCl = CaCl{2} + CO{2} + H{2}O.

Physical properties. Carbon dioxide is a colorless, practically odorless gas whose density is 1.5. Its weight may be inferred from the fact that it can be siphoned, or poured like water, from one vessel downward into another. At 15 deg. and under ordinary pressure it dissolves in its own volume of water and imparts a somewhat biting, pungent taste to it. It is easily condensed, and is now prepared commercially in this form by pumping the gas into steel cylinders (see Fig. 6) which are kept cold during the process. When the liquid is permitted to escape into the air part of it instantly evaporates, and in so doing absorbs so much heat that another portion is solidified, the solid form strikingly resembling snow in appearance. This snow is very cold and mercury can easily be frozen with it.

Solid carbon dioxide. Cylinders of liquid carbon dioxide are inexpensive, and should be available in every school. To demonstrate the properties of solid carbon dioxide, the cylinder should be placed across the table and supported in such a way that the stopcock end is several inches lower than the other end. A loose bag is made by holding the corners of a handkerchief around the neck of the stopcock, and the cock is then turned on so that the gas rushes out in large quantities. Very quickly a considerable quantity of the snow collects in the handkerchief. To freeze mercury, press a piece of filter paper into a small evaporating dish and pour the mercury upon it. Coil a flat spiral upon the end of a wire, and dip the spiral into the mercury. Place a quantity of solid carbon dioxide upon the mercury and pour 10 cc.-15 cc. of ether over it. In a minute or two the mercury will solidify and may be removed from the dish by the wire serving as a handle. The filter paper is to prevent the mercury from sticking to the dish; the ether dissolves the solid carbon dioxide and promotes its rapid conversion into gas.

Chemical properties. Carbon dioxide is incombustible, since it is, like water, a product of combustion. It does not support combustion, as does nitrogen peroxide, because the oxygen in it is held in very firm chemical union with the carbon. Very strong reducing agents, such as highly heated carbon, can take away half of its oxygen:

CO_{2} + C = 2CO.

Uses. The relation of carbon dioxide to plant life has been discussed in a previous chapter. Water highly charged with carbon dioxide is used for making soda water and similar beverages. Since it is a non-supporter of combustion and can be generated readily, carbon dioxide is also used as a fire extinguisher. Some of the portable fire extinguishers are simply devices for generating large amounts of the gas. It is not necessary that all the oxygen should be kept away from the fire in order to smother it. A burning candle is extinguished in air which contains only 2.5% of carbon dioxide.

Carbonic acid (H{2}CO{3}). Like most of the oxides of the non-metallic elements, carbon dioxide is an acid anhydride. It combines with water to form an acid of the formula H{2}CO{3}, called carbonic acid:

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

The acid is, however, very unstable and cannot be isolated. Only a very small amount of it is actually formed when carbon dioxide is passed into water, as is evident from the small solubility of the gas. If, however, a base is present in the water, salts of carbonic acid are formed, and these are quite stable:

2NaOH + H_{2}O + CO_{2} = Na_{2}CO_{3} + 2H_{2}O.

Action of carbon dioxide on bases. This conduct is explained by the principles of reversible reactions. The equation

H{2}O +CO{2} H{2}CO{3}

is a reversible equation, and the extent to which the reaction progresses depends upon the relative concentrations of each of the three factors in it. Equilibrium is ordinarily reached when very little H_{2}CO_{3} is formed. If a base is present in the water to combine with the H_{2}CO_{3} as fast as it is formed, all of the CO_{2} is converted into H_{2}CO_{3}, and thence into a carbonate.

Salts of carbonic acid,—carbonates. The carbonates form a very important class of salts. They are found in large quantities in nature, and are often used in chemical processes. Only the carbonates of sodium, potassium, and ammonium are soluble, and these can be made by the action of carbon dioxide on solutions of the bases, as has just been explained.

The insoluble carbonates are formed as precipitates when soluble salts are treated with a solution of a soluble carbonate. Thus the insoluble calcium carbonate can be made by bringing together solutions of calcium chloride and sodium carbonate:

CaCl{2} + Na{2}CO{3} = CaCO{3} + 2NaCl.

Most of the carbonates are decomposed by heat, yielding an oxide of the metal and carbon dioxide. Thus lime (calcium oxide) is made by strongly heating calcium carbonate:

CaCO{3} = CaO + CO{2}.

Acid carbonates. Like all acids containing two acid hydrogen atoms, carbonic acid can form both normal and acid salts. The acid carbonates are made by treating a normal carbonate with an excess of carbonic acid. With few exceptions they are very unstable, heat decomposing them even when in solution.

Action of carbon dioxide on calcium hydroxide. If carbon dioxide is passed into clear lime water, calcium carbonate is at first precipitated:

H{2}O + CO{2} = H{2}CO{3},

Ca(OH)_{2} + H_{2}CO_{3} = CaCO_{3} + 2H_{2}O.

Advantage is taken of this reaction in testing for the presence of carbon dioxide, as already explained in the chapter on the atmosphere. If the current of carbon dioxide is continued, the precipitate soon dissolves, because the excess of carbonic acid forms calcium acid carbonate which is soluble:

CaCO_{3} + H_{2}CO_{3} = Ca(HCO_{3})_{2}.

If now the solution is heated, the acid carbonate is decomposed and calcium carbonate once more precipitated:

Ca(HCO_{3})_{2} = CaCO_{3} + H_{2}CO_{3}.

Carbon monoxide (CO). Carbon monoxide can be made in a number of ways, the most important of which are the three following:

1. By the partial oxidation of carbon. If a slow current of air is conducted over highly heated carbon, the monoxide is formed, thus:

C + O = CO

It is therefore often formed in stoves when the air draught is insufficient. Water gas, which contains large amounts of carbon monoxide, is made by partially oxidizing carbon with steam:

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

2. By the partial reduction of carbon dioxide. When carbon dioxide is conducted over highly heated carbon it is reduced to carbon monoxide by the excess of carbon:

CO_{2} + C = 2CO.

When coal is burning in a stove or grate carbon dioxide is at first formed in the free supply of air, but as the hot gas rises through the glowing coal it is reduced to carbon monoxide. When the carbon monoxide reaches the free air above the coal it takes up oxygen to form carbon dioxide, burning with the blue flame so familiar above a bed of coals, especially in the case of hard coals.

3. _By the decomposition of oxalic acid._ In the laboratory carbon monoxide is usually prepared by the action of concentrated sulphuric acid upon oxalic acid. The latter substance has the formula C_{2}H_{2}O_{4}. The sulphuric acid, owing to its affinity for water, decomposes the oxalic acid, as represented in the equation

C_{2}H_{2}O_{4} + (H_{2}SO_{4}) = (H_{2}SO_{4}) + H_{2}O + CO_{2} + CO.

Properties. Carbon monoxide is a light, colorless, almost odorless gas, very difficult to liquefy. Chemically it is very active, combining directly with a great many substances. It has a great affinity for oxygen and is therefore combustible and a good reducing agent. Thus, if carbon monoxide is passed over hot copper oxide, the copper is reduced to the metallic state:

CuO + CO = Cu + CO_{2}.

When inhaled it combines with the red coloring matter of the blood and in this way prevents the absorption of oxygen, so that even a small quantity of the gas may prove fatal.



The reducing power of carbon monoxide. Fig. 61 illustrates a method of showing the reducing power of carbon monoxide. The gas is generated by gently heating 7 or 8 g. of oxalic acid with 25 cc. of concentrated sulphuric acid in a 200 cc. flask A. The bottle B contains a solution of sodium hydroxide, which removes the carbon dioxide formed along with the monoxide. C contains a solution of calcium hydroxide to show that the carbon dioxide is completely removed. E is a hard-glass tube containing 1 or 2 g. of copper oxide, which is heated by a burner. The black copper oxide is reduced to reddish metallic copper by the carbon monoxide, which is thereby changed to carbon dioxide. The presence of the carbon dioxide is shown by the precipitate in the calcium hydroxide solution in D. Any unchanged carbon monoxide is collected over water in F.

Carbon disulphide (CS{2}). Just as carbon combines with oxygen to form carbon dioxide, so it combines with sulphur to form carbon disulphide (CS{2}). This compound has been described in the chapter on sulphur.

Hydrocyanic acid (prussic acid)(HCN). Under the proper conditions carbon unites with nitrogen and hydrogen to form the acid HCN, called hydrocyanic acid. It is a weak, volatile acid, and is therefore easily prepared by treating its salts with sulphuric acid:

KCN + H_{2}SO_{4} = KHSO_{4} + HCN.

It is most familiar as a gas, though it condenses to a colorless liquid boiling at 26 deg.. It has a peculiar odor, suggesting bitter almonds, and is extremely poisonous either when inhaled or when taken into the stomach. A single drop may cause death. It dissolves readily in water, its solution being commonly called prussic acid.

The salts of hydrocyanic acid are called cyanides, the cyanides of sodium and potassium being the best known. These are white solids and are extremely poisonous.

Solutions of potassium cyanide are alkaline. A solution of potassium cyanide turns red litmus blue, and must therefore contain hydroxyl ions. The presence of these ions is accounted for in the following way.

Although water is so little dissociated into its ions H^{} and OH^{-} that for most purposes we may neglect the dissociation, it is nevertheless measurably dissociated. Hydrocyanic acid is one of the weakest of acids, and dissociates to an extremely slight extent. When a cyanide such as potassium cyanide dissolves it freely dissociates, and the CN^{-} ions must come to an equilibrium with the H^{} ions derived from the water:

H^{+} + CN^{-} HCN.

The result of this equilibrium is that quite a number of H^{} ions from the water are converted into undissociated HCN molecules. But for every H^{} ion so removed an OH^{-} ion remains free, and this will give the solution alkaline properties.

EXERCISES

1. How can you prove that the composition of the different allotropic forms of carbon is the same?

2. Are lampblack and bone black allotropic forms of carbon? Will equal amounts of heat be liberated in the combustion of 1 g. of each?

3. How could you judge of the relative purity of different forms of carbon?

4. Apart from its color, why should carbon be useful in the preparation of inks and paints?

5. Could asbestos fibers be used to replace the wire in a safety lamp?

6. Why do most acids decompose carbonates?

7. What effect would doubling the pressure have upon the solubility of carbon dioxide in water?

8. What compound would be formed by passing carbon dioxide into a solution of ammonium hydroxide? Write the equation.

9. Write equations for the preparation of K{2}CO{3}; of BaCO{3}; of MgCO{3}.

10. In what respects are carbonic and sulphurous acids similar?

11. Give three reasons why the reaction which takes place when a solution of calcium acid carbonate is heated, completes itself.

12. How could you distinguish between carbonates and sulphites?

13. How could you distinguish between oxygen, hydrogen, nitrogen, nitrous oxide, and carbon dioxide?

14. Could a solution of sodium hydroxide be substituted for the solution of calcium hydroxide in testing for carbon dioxide?

15. What weight of sodium hydroxide is necessary to neutralize the carbonic acid formed by the action of hydrochloric acid on 100 g. of calcium carbonate?

16. What weight of calcium carbonate would be necessary to prepare sufficient carbon dioxide to saturate 10 l. of water at 15 deg. and under ordinary pressure?

17. On the supposition that calcium carbide costs 12 cents a kilogram, what would be the cost of an amount sufficient to generate 100 l. of acetylene measured at 20 deg. and 740 mm.?

18. How would the volume of a definite amount of carbon monoxide compare with the volume of carbon dioxide formed by its combustion, the measurements being made under the same conditions?



CHAPTER XVIII

FLAMES,—ILLUMINANTS

Conditions necessary for flames. It has been seen that when two substances unite chemically, with the production of light and heat, the act of union is called combustion. When one of the substances undergoing combustion remains solid at the temperature occasioned by the combustion, light may be given off, but there is no flame. Thus iron wire burning in oxygen throws off a shower of sparks and is brilliantly incandescent, but no flame is seen. When, however, both of the substances are gases or vapors at the temperature reached in the combustion, the act of union is accompanied by a flame.

Flames from burning liquids or solids. Many substances which are liquids or solids at ordinary temperatures burn with a flame because the heat of combustion vaporizes them slowly, and the flame is due to the union of this vapor with the gas supporting the combustion.

Supporter of combustion. That gas which surrounds the flame and constitutes the atmosphere in which the combustion occurs is said to support the combustion. The other gas which issues into this atmosphere is said to be the combustible gas. Thus, in the ordinary combustion of coal gas in the air the coal gas is said to be combustible, while the air is regarded as the supporter of combustion. These terms are entirely relative, however, for a jet of air issuing into an atmosphere of coal gas will burn when ignited, the coal gas supporting the combustion. Ordinarily, when we say that a gas is combustible we mean that it is combustible in an atmosphere of air.



Either gas may be the supporter of combustion. That the terms combustible and supporter of combustion are merely relative may be shown in the following way: A lamp chimney A is fitted with a cork and glass tubes, as shown in Fig. 62. The tube C should have a diameter of from 12 to 15 mm. A thin sheet of asbestos in which is cut a circular opening about 2 cm. in diameter is placed over the top of the chimney. The opening in the asbestos is closed with the palm of the hand, and gas is admitted to the chimney through the tube B. The air in the chimney is soon expelled through the tube C, and the gas itself is then lighted at the lower end of this tube. The hand is now removed from the opening in the asbestos, when the flame at the end of the tube at once rises and appears at the end within the chimney, as shown in the figure. The excess of coal gas now escapes from the opening in the asbestos and may be lighted. The flame at the top of the asbestos board is due to the combustion of coal gas in air, while the flame within the chimney is due to the combustion of air in coal gas, the air being drawn up through the tube by the escaping gas.