Carbon

Carbon (symbol C, atomic weight 12), one of the chemical non-metallic elements. It is found native as the diamond (q.v.), graphite (q.v.), as a constituent of all animal and vegetable tissues and of coal and petroleum. It also enters (as carbonates) into the composition of many minerals, such as chalk, dolomite, calcite, witherite, calamine and spathic iron ore. In combination with oxygen (as carbon dioxide) it is also found to a small extent in the atmosphere. It is a solid substance which occurs in several modifications, differing very much in their physical properties. Amorphous carbon is obtained by the destructive distillation of many carbon compounds, the various kinds differing very greatly as regards physical characters and purity, according to the substance used for their preparation. The most common varieties met with are lampblack, gas carbon, wood charcoal, animal charcoal and coke. Lampblack is prepared by burning tar, resin, turpentine and other substances rich in carbon, with a limited supply of air; the products of combustion being conducted into condensing chambers in which cloths are suspended, on which the carbon collects. It is further purified by heating in closed vessels, but even then it still contains a certain amount of mineral matter and more or less hydrocarbons. It is used in the manufacture of printer’s ink, in the preparation of black paint and in calico printing. Gas carbon is produced by the destructive distillation of coal in the manufacture of illuminating gas (see Gas: Manufacture), being probably formed by the decomposition of gaseous hydrocarbons. It is a very dense form of carbon, and is a good conductor of heat and electricity. It is used in the manufacture of carbon rods for arc lights, and for the negative element in the Bunsen battery.

Charcoal is a porous form of carbon; several varieties exist. Sugar charcoal is obtained by the carbonization of sugar. It is purified by boiling with acids, to remove any mineral matter, and is then ignited for a long time in a current of chlorine in order to remove the last traces of hydrogen. Animal charcoal (bone black) is prepared by charring bones in iron retorts. It is a very impure form of carbon, containing on the average about 80% of calcium phosphate. It possesses a much greater decolorizing and absorbing power than wood charcoal. A variety of animal charcoal is sometimes prepared by calcining fresh blood with potassium carbonate in large cylinders, the mass being purified by boiling out with dilute hydrochloric acid and subsequent reheating. Wood charcoal is a hard and brittle black substance, which retains the external structure of the wood from which it is made. It is prepared (where wood is plentiful) by stacking the wood in heaps, which are covered with earth or with brushwood and turf, and then burning the heap slowly in a limited supply of air. The combustion of the wood is conducted from the top downwards, and from the exterior towards the centre; great care has to be taken that the process is carried out slowly. The disadvantage in this process is that the by-products, such as pyroligneous acid, acetone, wood spirit, &c., are lost; as an alternative method, wood is frequently carbonized in ovens or retorts and the volatile products are condensed and utilized.

Charcoal varies considerably in its properties, depending upon the particular variety of wood from which it is prepared, and also upon the process used in its manufacture. It can be made at a temperature as low as 300° C., and is then a soft, very friable material possessing a low ignition point. When made at higher temperatures it is much more dense, and its ignition point is considerably higher. Charcoal burns when heated in air, usually without the formation of flame, although a flame is apparent if the temperature be raised. It is characterized by its power of absorbing gases; thus, according to J. Hunter [Phil. Mag., 1863 (4), 25, p. 363], one volume of charcoal absorbs (at 0° C. and 760 mm. pressure) 171.7 ccs. of ammonia, 86.3 ccs. of nitrous oxide, 67.7 ccs. of carbon monoxide, 21.2 ccs. of carbon dioxide, 17.9 ccs. of oxygen, 15.2 ccs. of nitrogen, and 4.4 ccs. of hydrogen [see also J. Dewar, Ann. Chim. Phys., 1904 (8), 3, p. 5]. It also has the power of absorbing colouring matters from solution. Charcoal is used as a fuel and as a reducing agent in metallurgical processes.

The element carbon unites directly with hydrogen to form acetylene when an electric arc is passed between carbon poles in an atmosphere of hydrogen (M. Berthelot); it also unites directly with fluorine, producing, chiefly, carbon tetrafluoride CF4. It burns when heated in an atmosphere of oxygen, forming carbon dioxide, and when heated in sulphur vapour it forms carbon bisulphide (q.v.). When heated with nitrogenous substances, in the presence of carbonated or caustic alkali, it forms cyanides. It combines directly with silicon, at the temperature of the electric furnace, yielding carborundum, SiC; and H. Moissan has also shown that it will combine with many metals at the temperature of the electric furnace, to form carbides (q.v.).

The specific heat of carbon varies with the temperature the following values having been obtained by H.F. Weber (Jahresberichte, 1874, p. 63):—

Diamond. | Graphite. | Porous wood carbon.  
---|---|---  
t°. | Sp. Ht. | t°. | Sp. Ht. | t°. | Sp. Ht.  
−50.5 | 0.0635 | −50.3 | 0.1138 | 0−23 | 0.1653  
−10.6 | 0.0955 | −10.7 | 0.1437 | 0−99 | 0.1935  
+10.7 | 0.1128 | +10.8 | 0.1604 | 0−223 | 0.2385  
85.5 | 0.1765 | 61.3 | 0.1990 |   |    
206.1 | 0.2733 | 201.6 | 0.2966 |   |    
606.7 | 0.4408 | 641.9 | 0.4454 |   |    
985.0 | 0.4589 | 977.0 | 0.4670 |   |    
  
The atomic weight of carbon has been determined by J.B.A. Dumas and by J.S. Stas [Ann. Chim. Phys., 1841 (3), 1, p. 1: Jahresb., 1849, 223] by estimating the amount of carbon dioxide formed on burning graphite or diamond in a current of oxygen, the value obtained being 12.0 (O = 16). Confirmatory evidence has also been obtained by O.L. Erdmann and R.F. Marchand (Jour. Prak. Chem., 1841, 23, p. 159; see also F.W. Clarke, Jahresb., 1881, p. 7).

Compounds.—Three oxides of carbon are known, namely, carbon suboxide, C3O2, carbon monoxide, CO, and carbon dioxide, CO2. Carbon suboxide, C3O2, is formed by the action of phosphorus pentoxide on ethyl malonate (O. Diels and B. Wolf, Ber., 1906, 39, p. 689), CH2(COOC2H5)2 = 2C2H4 \+ 2H2O + C3O2. At ordinary temperatures it is a colourless gas, possessing a penetrating and suffocating smell. It liquefies at 7° C. It is an exceedingly reactive compound, combining with water to form malonic acid, with hydrogen chloride to form malonyl chloride, and with ammonia to form malonamide. When kept for some time in sealed tubes it changes to a yellowish liquid, from which a yellow flocculent substance gradually separates, and finally it suddenly solidifies to a dark red mass, which appears to be a polymeric form. Its vapour density agrees with the molecular formula C3O2, and this formula is also confirmed by exploding the gas with oxygen and measuring the amount of carbon dioxide produced (see Ketenes).

Carbon monoxide, CO, is found to some extent in volcanic gases. It was first prepared in 1776 by J.M.F. Lassone (Mem. Acad. Paris) by heating zinc oxide with carbon, and was for some time considered to be identical with hydrogen. Cruikshank concluded that it was an oxide of carbon, a fact which was confirmed by Clement and J.B. Désormes (Ann. Chim. Phys., 1801, 38, p. 285). It may be prepared by passing carbon dioxide over red-hot carbon, or red-hot iron; by heating carbonates (magnesite, chalk, &c.) with zinc dust or iron; or by heating many metallic oxides with carbon. It may also be prepared by heating formic and oxalic acids (or their salts) with concentrated sulphuric acid (in the case of oxalic acid, an equal volume of carbon dioxide is produced); and by heating potassium ferrocyanide with a large excess of concentrated sulphuric acid, K4Fe(CN)6 \+ 6H2SO4 \+ 6H2O = 2K2SO4 \+ FeSO4 \+ 3(NH4)2SO4 \+ 6CO. It is a colourless, odourless gas of specific gravity 0.967 (air = 1). It is one of the most difficultly liquefiable gases, its critical temperature being −139.5° C., and its critical pressure 35.5 atmos. The liquid boils at −190° C., and solidifies at −211°C. (L.P. Cailletet, Comptes rendus, 1884, 99, p. 706). It is only very slightly soluble in water. It burns with a characteristic pale blue flame to form carbon dioxide. It is very poisonous, uniting with the haemoglobin of the blood to form carbonyl-haemoglobin. It is a powerful reducing agent, especially at high temperatures. It is rapidly absorbed by an ammoniacal or acid (hydrochloric acid) solution of cuprous chloride. It unites directly with chlorine, forming carbonyl chloride or phosgene (see below), and with nickel and iron to form nickel and iron carbonyls (see Nickel and Iron). It also combines directly with potassium hydride to form potassium formate (see Formic Acid). The volume composition of carbon monoxide is established by exploding a mixture of the gas with oxygen, two volumes of the gas combining with one volume of oxygen to form two volumes of carbon dioxide. This fact, coupled with the determination of the vapour density of the gas, establishes the molecular formula CO.

Carbon dioxide, CO2, is a gas first distinguished from air by van Helmont (1577-1644), who observed that it was formed in fermentation processes and during combustion, and gave to it the name gas sylvestre. J. Black (Edin. Phys. and Lit. Essays, 1755) showed that it was a constituent of the carbonated alkalis and called it “fixed air.” T.O. Bergman, in 1774, pointed out its acid character, and A.L. Lavoisier (1781-1788) first proved it to be an oxide of carbon by burning carbon in the oxygen obtained from the decomposition of mercuric oxide. It is a regular constituent of the atmosphere, and is found in many spring waters and in volcanic gases; it also occurs in the uncombined condition at the Grotto del Cane (Naples) and in the Poison Valley (Java). It is a constituent of the minerals cerussite, malachite, azurite, spathic iron ore, calamine, strontianite, witherite, calcite aragonite, limestone, &c. It may be prepared by burning carbon in excess of air or oxygen, by the direct decomposition of many carbonates by heat, and by the decomposition of carbonates with mineral acids, M2CO8 \+ 2HCl = 2MCl + H2O + CO2. It is also formed in ordinary fermentation processes, in the combustion of all carbon compounds (oil, gas, candles, coal, &c.), and in the process of respiration.

It is a colourless gas, possessing a faint pungent smell and a slightly acid taste. It does not burn, and does not support ordinary combustion, but the alkali metals and magnesium, if strongly heated, will continue to burn in the gas with formation of oxides and liberation of carbon. Its specific gravity is 1.529 (air = 1). It is readily condensed, passing into the liquid condition at 0° C. under a pressure of 35 atmospheres. Its critical temperature is 31.35° C., and its critical pressure is 72.9 atmos. The liquid boils at −78.2° C. (l atmo.), and by rapid evaporation can be made to solidify to a snow-white solid which melts at −65° C.(see Liquid Gases). Carbon dioxide is moderately soluble in water, its coefficient of solubility at 0° C. being 1.7977 (R. Bunsen). It is still more soluble in alcohol. The solution of the gas in water shows a faintly acid reaction and is supposed to contain carbonic acid, H2CO3. The gas is rapidly absorbed by solutions of the caustic alkalis, with the production of alkaline carbonates (q.v.), and it combines readily with potassium hydride to form potassium formate. It unites directly with ammonia gas to form ammonium carbamate, NH2COONH4. It may be readily recognized by the white precipitate which it forms when passed through lime or baryta water. Carbon dioxide dissociates, when strongly heated, into carbon monoxide and oxygen, the reaction being a balanced action; the extent of dissociation for varying temperatures and pressures has been calculated by H. Le Chateller (Zeit. Phys. Chem., 1888, 2, p. 782; see H. Sainte-Claire Deville, Comptes rendus, 1863, 56, p. 195 et seq.). The volume composition of carbon dioxide is determined by burning carbon in oxygen, when it is found that the volume of carbon dioxide formed is the same as that of the oxygen required for its production, hence carbon dioxide contains its own volume of oxygen. Carbon dioxide finds industrial application in the preparation of soda by the Solvay process, in the sugar industry, in the manufacture of mineral waters, and in the artificial production of ice.

Carbonyl chloride (phosgene), COCl2, was first obtained by John Davy (Phil. Trans., 1812, 40, p. 220). It may be prepared by the direct union of carbon monoxide and chlorine in sunlight (Th. Wilm and G. Wischin, Ann., 1868, 14, p. 150); by the action of phosphorus pentoxide on carbon tetrachloride at 200-210° C. (G. Gustavson, Ber., 1872, 5, p. 30), 4CCl4 \+ P4O10 = 2CO2 \+ 4POCl3 \+ 2COCl2; by the oxidation of chloroform with chromic acid mixture (A. Emmerling and B. Lengyel, Ber., 1869, 2, p. 54), 4CHCl3 \+ 3O2 = 4COCl2 \+ 2H2O + 2Cl2; or most conveniently by heating carbon tetrachloride with fuming sulphuric acid (H. Erdmann, Ber., 1893, 26, p. 1993), 2SO3 \+ CCl4 = S2O5Cl2 \+ COCl2.

It is a colourless gas, possessing an unpleasant pungent smell. Its vapour density is 3.46 (air = 1). It may be condensed to a liquid, which boils at 8° C. It is readily soluble in benzene, glacial acetic acid, and in many hydrocarbons. Water decomposes it violently, with formation of carbon dioxide and hydrochloric acid. It reacts with alcohol to form chlorcarbonic ester and ultimately diethyl carbonate (see Carbonates), and with ammonia it yields urea (q.v.). It is employed commercially in the production of colouring matters (see Benzophenone), and for various synthetic processes.

Carbon oxysulphide, COS, was first prepared by C. Than in 1867 (Ann. Suppl., 5, p. 236) by passing carbon monoxide and sulphur vapour through a tube at a moderate heat. It is also formed by the action of sulphuretted hydrogen on the isocyanic esters, 2CONC2H5 \+ H2S = COS + CO(NHC2H5)2, by the action of concentrated sulphuric acid on the isothiocyanic esters, RNCS + H2O = COS + RNH2, or of dilute sulphuric acid on the thiocyanates. In the latter reaction various other compounds, such as carbon dioxide, carbon bisulphide and hydrocyanic acid, are produced. They are removed by passing the vapours in succession through concentrated solutions of the caustic alkalis, concentrated sulphuric acid, and triethyl phosphine; the residual gas is then purified by liquefaction (W. Hempel, Zeit. angew. Chemie, 1901, 14, p. 865). It is also formed when sulphur trioxide reacts with carbon bisulphide at 100° C., CS2 \+ 3SO3 = COS + 4SO2, and by the decomposition of ethyl potassium thiocarbonate with hydrochloric acid, CO(OC2H5)SK + HCl = COS + KCl + C2H5OH. It is a colourless, odourless gas, which burns with a blue flame and is decomposed by heat. Its vapour density is 2.1046 (air = 1). The liquefied gas boils at −47° C. under atmospheric pressure. It is soluble in water; the aqueous solution gradually decomposes on standing, forming carbon dioxide and sulphuretted hydrogen. It is easily soluble in solutions of the caustic alkalis, provided they are not too concentrated, forming solutions of alkaline carbonates and sulphides, COS + 4KHO = K2CO3 \+ K2S + 2H2O.

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