Chemical elements
  Carbon
    Isotopes
    Energy
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    Application
    Physical Properties
    Chemical Properties
    Diamonds
    Graphite
    Amorphous Carbon
      Preparation
      Sugar-Charcoal
      Lampblack and Soot
      Retort Carbon
      Wood-Charcoal
      Animal Charcoal
      Coke
      Physical Properties
      Chemical Properties
      Fusion and Vaporisation
    Coal

Chemical Properties of Amorphous Carbon






In accordance with its less density and the absence of crystalline form, amorphous carbon is more susceptible to oxidation by chemical reagents than graphite or diamond. A mixture of hot concentrated nitric acid and potassium chlorate, which has no action on diamond and converts graphite into graphitic acid, more or less readily dissolves amorphous carbon, converting it into humic and mellitic acids. This reaction, which serves to distinguish between the different allotropic forms of carbon, is known as Brodie's reaction.

Sugar-charcoal, wood-charcoal, lignite, coal, anthracite, and soot are acted upon when heated with concentrated nitric acid alone, giving brown solutions; coke- and gas-carbon do not colour the acid. A mixture of concentrated nitric and sulphuric acids dissolves all forms of amorphous carbon; but coke is the most resistant towards this reagent. Lignite and coal are oxidised to oxalic acid by alkaline permanganate solution; and alkali hypobromite dissolves lignite.

Amorphous carbon is oxidised to carbon dioxide when heated with sulphuric acid and potassium dichromate, or with molten potassium chlorate or nitrate; it combines with hydrogen under certain conditions to form CH4 and C2H2 (q-v.), and with sulphur vapour, forming CS2. Carbon also combines at high temperatures with metals, such as Fe, Ni, Co, to form carbides - as, for instance, in the cementation and case-hardening processes for steel - and reacts with lime in the electric furnace to form calcium carbide (q.v.).


Amorphous Carbon as a Reducing Agent

Besides being used as a fuel, amorphous carbon, in one form or another, finds application as a reducing agent. The "blowpipe reactions on charcoal," familiar to every student of qualitative analysis, exemplify the use of carbon as a reducing agent. In these reactions metallic oxides, such as those of silver, bismuth, copper, lead, tin, are reduced to the metallic state.

The use of coke in the blast-furnace, of anthracite in the reduction of tin and zinc ores, the process of "poling" in copper-smelting, are examples of reduction by carbon in metallurgy. The carbon anodes employed in the electrometallurgy of aluminium are consumed in the process, and no doubt their material aids the reduction of the alumina by combining with its oxygen. The preparation of phosphorus also involves the reduction of phosphate by charcoal or coke.

Combustion of Carbon

When carbon burns in excess of air or oxygen the product consists of carbon dioxide only; when, however, carbon dioxide comes in contact with glowing carbon it is reduced to carbon monoxide. It is an interesting question whether CO is first formed by the direct combustion of carbon or whether the product of combustion of carbon is at once and only CO2.

Lang believed that all the CO formed when oxygen was passed over gas-carbon at 500° owed its origin to the reduction of CO2, but Dixon pointed out that Lang's experimental results did not preclude the formation of CO previous to CO2; and, moreover, that CO2 is not reduced to CO below 600°. H. B. Baker's experiments appear to show that carbon burns in two stages, and that when thoroughly dried carbon and oxygen combine the chief product is CO. C. J. Baker found that dry carbon which had adsorbed dry oxygen at 12° C. gave off a gas at 450° C. which was mainly CO. Dixon has shown that CO is formed before CO2 in the combustion of cyanogen, i.e. of gaseous carbon. Finally Rhead and Wheeler have proved from a consideration of the relative rates of the various reactions between oxygen, CO, CO2, and carbon that in the burning of carbon the two oxides are produced simultaneously; and they 2 are of opinion that the oxygen first enters the carbon molecule, oxygenating it by forming an unstable physico-chemical complex, CO, and that the energy of this combination causes some of the oxygen molecules to detach some of the carbon atoms and depart with them as molecules of carbon dioxide, whilst other oxygen molecules are "torn apart in the process - become atomised - and leave the carbon molecule as carbon monoxide."

Heat of Combustion of Carbon, and Relations between the Different Allotropic Forms

Small differences have been observed between the heats of combustion of the different allotropic forms of carbon. The following values were obtained by Favre and Silbermann for the reaction [C,O2], i.e. for the combustion of 12 grams of carbon to 44 grams of carbon dioxide:

Diamond (1). - 93,240 calories; Diamond (2). - 94.650 calories; Natural graphite - 93,560 calories; Cast-iron graphite - 93,140 calories; Wood-charcoal - 96,960 calories;

The heat of the reaction [C,O] was found to be 29,000 calories.

The following figures were obtained by Berthelot:

Diamond - 94,310 calories; Graphite - 94,810 calories; Amorphous Carbon - 97,650 calories.

whence the following heats of formation are calculated:

Diamond from amorphous carbon - 3340 calories.
Graphite from amorphous carbon - 2840 calories.
Diamond from graphite - 500 calories.

According to Mixter, the carbon obtained by the decomposition of acetylene is to be regarded as a distinct allotropic form, whose heat of combustion is 94,730 calories, whilst that of sugar-charcoal is 96,680 calories, and of graphite 93,970 calories. The density of this form is 1.919, and Moissan has shown that it is not identical with graphite.

From a consideration of the above thermal differences between diamond, graphite, and amorphous carbon it appears that the formation of graphite from diamond is an endothermic reaction, and that diamond contains less internal energy than the other forms of carbon, and is therefore probably the most stable modification at ordinary temperatures.

From a study of equilibrium in the systems C, CO, CO2 and Fe, FeO, CO, CO2 Schenckand Heller conclude that between 400° C. and 800° C. graphite is the most stable and charcoal the least stable form of carbon. These considerations have an important bearing on the reduction of iron in the blast-furnace. Falcke,1 however, throws doubt on Schenck's conclusions.

Physical states of carbon
Physical states of carbon
Roozeboom points out that from 3000° C., at which temperature diamond is Converted into graphite, down to about 1000° C., when graphite is still the most stable form of carbon, diamond is in a metastable condition (Fig.). No transition-point is known, however, between graphite and diamond; therefore it appears uncertain which is the more stable form at atmospheric temperature. Nevertheless, by Nernst's theorem the transition temperature is calculated to be 372° C., below which temperature diamond would be the most stable form of carbon. It has further been calculated that graphite is converted into diamond at the temperature of molten iron, under a pressure of about 10,000 atmospheres.
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