Cyanogen, (CN)₂

Cyanogen gas is an endothermic compound which is formed by the union of its elements at very high temperature. According to von Wartenberg 44 per cent, of cyanogen should be produced when carbon and nitrogen are in equilibrium at 3500° C., the temperature of the electric arc. The presence of this gas in the arc is shown by its spectrum, but Wallis failed to find it in the gases withdrawn from the arc chamber. This may be due to the catalytic decomposition of cyanogen by carbon particles, or by its union with hydrogen from the arc carbons to form HCN.

Cyanogen is formed in small quantity in the gases from coke-ovens and blast-furnaces; and also when a mixture of coal-gas and ammonia is burnt in a Bunsen burner, the reaction probably being:

4CO + 4NH₃ + O₂ = 2(CN)₂ + 6H₂O.

It may also be obtained by dehydrating ammonium oxalate by heating it with phosphoric oxide:

H₄NOOC-COONH₄ – 4H₂O = NC-CN;

but is most conveniently prepared by heating certain metallic cyanides. Metals whose oxides yield metal and oxygen when heated, i.e. mercury, silver, and gold, are those whose cyanides give in the same way metal and cyanogen. In addition to these, cupric cyanide very readily decomposes like cupric iodide, yielding cuprous cyanide and cyanogen.

Cyanogen is frequently prepared by heating powdered mercuric cyanide in a hard glass tube, and is collected over mercury. Much heat is needed for this reaction, as is shown by the thermal equation:

Hg(CN)₂ = Hg + (CN)₂ - 19,000 calories;

admixture of mercuric chloride, however, enables the reaction to proceed at a lower temperature, owing to the formation of mercurous chloride by an exothermic reaction, thus:

Hg(CN)₂ + HgCl₂ = Hg₂Cl₂ + (CN)₂ + 440 calories.

A mixture of potassium ferrocyanide and mercuric chloride also when heated yields cyanogen, probably mixed with nitrogen, The most convenient way, however, of obtaining cyanogen is by warming together potassium cyanide and cupric sulphate solutions. Cyanogen is evolved, and cuprous cyanide precipitated:

2CuSO₄ + 4KCN = 2K₂SO₄ + Cu₂(CN)₂ + (CN)₂.

If the cuprous cyanide is filtered, washed, and treated with ferric chloride solution the remainder of the cyanogen is evolved, thus:

Cu₂(CN)₂ + 2FeCl₃ = Cu₂Cl₂ + 2FeCl₂ + (CN)₂.

Cyanogen is a colourless, poisonous gas with an odour resembling that of peach kernels. Its density is 1.8064 (Gay Lussac), the theoretical density being 1.7993. The gas shows no sign of dissociation at 800° C.; but is liquefied at -20.7° C. under 1 atm. The vapour pressures at different temperatures are as follow:

Temperature ° C	-20.7°	0°	5°	10°	15°
Vapour pressure (atm.).	1.0	2.37	2.83	3.38	4.04

The density of liquid cyanogen at 17.2° C. is 0.866; according to Dewar its critical temperature is 124° C. and critical pressure 61.7 atm. whence the constants for van der Waals equation are derived: a = 0.01446; b = 0.0029. Cardoso and Baume, however, give: critical temperature 128.3° C., critical pressure 59.6 atm. When cooled below -35° C. liquid cyanogen freezes to a crystalline mass which melts at -34.4° C.

The heat of formation of cyanogen from graphite and nitrogen is -73,000 to -70,000 calories; its molecular heat of combustion to CO₂ and N₂ is 262,500 calories (Berthelot) or 259,600 calories (Thomsen).

Water, at atmospheric temperature, dissolves about four and a half times its volume of cyanogen gas, alcohol about twenty-three times its volume. Water, however, acts chemically on cyanogen, causing the separation of a brown insoluble substance known as azulmic acid, and at the same time hydrolysing the cyanogen (NC-CN), producing from it oxamide (H₂NOC-CONH₂), oxamic acid (HOOC-CONH₂), oxalic acid (HOOC-COOH), and their ammonium salts; together with the following compounds containing only one carbon atom: hydrocyanic acid, urea (H₂N-CO-NH₂), and ammonium carbonate. When cyanogen is passed into water at 0° C., however, the reaction is simple, hydrocyanic and cyanic acids being formed thus:

(CN)₂ + H₂O = HCN + HCNO.

This reaction has been studied by measuring the conductivity due to hydrocyanic acid, cyanic acid being almost a non-conductor.

The chemical constitution of cyanogen, N≡C-C≡N, follows from the reactions which this gas undergoes. Thus it is reduced by tin and hydrochloric acid to ethylene diamine, which is NH₂-CH₂-CH₂-NH₂, and is hydrolysed by water and dilute acids as follows:

→ → →

Cyanogen → Oxamide → Oxamic acid → Oxalic acid.

Cyanogen is thus the nitrile of oxalic acid.

Dixon and Taylor, however, favour a cyclic constitution for cyanogen: , partly on account of its preparation by heating silver and mercuric cyanides, which are probably iso-compounds, the reaction being supposed to proceed thus:

2Ag.N=C → 2Ag + →

and also on account of its reaction with alkalis being similar to that of cyanogen bromide, which suggests that one cyanogen group behaves towards the other like a halogen atom, thus:

:,

the formula for cyanogen bromide being that of Gutmann. Dixon and Taylor suggest that if cyanogen has an iso- constitution oxalic acid might be derived from it, thus:

→ → →

There is a polymeric modification of cyanogen known as paracyanogen (CN)_n. It is a brown powder of unknown molecular weight and constitution, whose formation is observed when mercuric cyanide is heated. It is produced when cyanogen is heated for some time at 440° C., and is formed at the anode during the electrolysis of potassium cyanide solution. Cyanogen is converted into paracyanogen under ordinary pressure at about 310° C., and under 300 atmospheres pressure at 220° C. Paracyanogen yields ordinary cyanogen gas, which condenses to liquid cyanogen. The two forms are thus monotropic like white and red phosphorus. Paracyanogen is insoluble in water and alcohol; it shows, however, many of the reactions of cyanogen, although its reactivity is slow.

Cyanogen is decomposed into its elements by being passed through a red-hot tube, or by the agency of electric sparks or the arc flame. It burns with a characteristic peach-coloured flame, which consists of a red internal and a blue external zone separable by means of Smithells' apparatus. In the inner flame the reaction (CN)₂ + O₂ = 2CO + N₂ takes place; in the outer flame CO burns to CO₂.

Cyanogen can burn or be exploded in dried oxygen, one or other of the following reactions occurring in the latter case according to the volume of oxygen employed:

(CN)₂ + O₂ = 2CO + N₂.
(CN)₂ + 2O₂ = 2CO₂ + N₂.

A mixture of dried oxygen and carbon monoxide is not explosive, but the admixture of 10 per cent, of cyanogen gas secures the explosion of the carbon monoxide in the absence of moisture.

Whilst behaving as the nitrile of oxalic acid in some of its reactions, in others cyanogen resembles free chlorine. Thus it unites with hydrogen to form hydrocyanic acid and with metals such as potassium, zinc, copper, lead, mercury, and silver to produce cyanides. Its reaction with sulphurous acid, though slower than, is similar to, that of the halogens: (CN)₂ + H₂SO₃ + H₂O = 2HCN + H₂SO₄.

Moreover, just as chlorine reacts with dilute alkali to form chloride and hypochlorite, so cyanogen yields cyanide and cyanate, thus:

(CN)₂ + 2KOH = KCN + KCNO + H₂O.

Cyanogen gas is easily detected by its smell, by its decomposition by electric sparks into nitrogen and carbon, and by its characteristic flame, whose spectrum contains prominent bands in the blue and violet. The presence of this gas in a mixture may be detected by its reaction with a mixture of 2 c.c. of cold saturated solution of picric acid, 18 c.c. of alcohol, and 5 c.c. of a 15 per cent, solution of caustic soda. With this solution cyanogen gives a purple-red colour turning brown, due to the formation of isopurpuric acid. If much cyanogen is present the potassium salt of this acid separates as a purple-red oil. Cyanogen may be Estimated by combustion in a stream of oxygen; explosion with oxygen does not give accurate results.