Combustion of Carbon Monoxide

From the time of Lavoisier until 1880, the Combustion of Carbon Monoxide was regarded as a simple oxidation process completely represented by the equation

2CO + O₂ = 2CO₂.

In the latter year, however, Dixon, in his address to the Chemical Section of the British Association, made the startling announcement that when an electric spark is passed through a thoroughly dried mixture of two volumes of carbon monoxide and one volume of oxygen, no explosion is caused. The introduction of a small quantity of water- vapour suffices to determine the explosion, which gains in intensity up to a certain point with increasing amounts of water. Gases like H₂S, C₂H₄, NH₃, HCl, and ether-vapour act like water, but SO₂, CO₂, CS₂, C₂N₂, CCl₄ do not, if perfectly dry. It seems clear, therefore, that such substances as will form steam under the conditions of the experiment are capable of determining the explosion.

In order to explain these results, Dixon suggested that "the carbon monoxide is oxidised by the steam in the path of the spark, and that the hydrogen set free unites with oxygen to form steam at a high temperature." The steam thus acts as an oxygen carrier itself, undergoing successive oxidation and reduction as follows:

(i) 2CO + 2OH₂ = 2CO₂ + 2H₂;
(ii) 2H₂ + O₂ = 2H₂O.

Moritz Traube rejected this explanation on the ground that carbon monoxide does not decompose steam at the temperature of the electric spark, for reaction (i) is reversible and under these conditions proceeds in the direction right to left.

This objection, however, is not valid, for, whatever the temperature, the law of mass action requires that definite, even if small, amounts of carbon dioxide and hydrogen shall exist in equilibrium with the other gases in the system. Hence, if for any reason the partial pressure of the carbon dioxide or hydrogen falls below that required for equilibrium, it is always possible for reaction (i) to proceed, even at high temperatures, in the direction of left to right. Traube explained the reaction, however, on the assumption that the function of the steam is to unite with one atom of the oxygen molecule, the second atom being occupied in the oxidation of the carbon monoxide. Thus

CO + O:O + OH₂ = CO₂ + H₂O₂.

The hydrogen peroxide is then reduced to steam by more carbon monoxide

CO + H₂O₂ = CO₂ + H₂O.

This is reminiscent of the Brodie-Schonbein theory of slow oxidation, and in support of it may be mentioned the fact that hydrogen peroxide can be detected when moist carbon monoxide is allowed to burn in air, the flame being made to impinge upon a cold surface of water.

The suggestion has also been made that percarbonic acid, H₂C₂O₆, may be formed intermediately, since by allowing a flame of the moist monoxide to impinge on a cold, dilute solution of cobalt chloride and potassium hydroxide, a precipitate is obtained closely resembling that yielded on adding potassium percarbonate to a solution of cobalt chloride. The Combustion of Carbon Monoxide would thus presumably proceed as follows:

4CO + 3O₂ + 2H₂O = 2H₂C₂O₆;
CO + H₂C₂O₆ = H₂O + 3CO₂.

The evidence is not conclusive, however, as the same precipitate is obtained with hydrogen peroxide in the presence of an alkaline cobalt salt solution. So the above test might well be cited as a further argument in favour of Traube's theory.

The mere fact that hydrogen peroxide can be detected in the above manner is no proof that it plays such an important part in flame kinetics as Traube suggests. Armstrong agrees with Dixon that water is the inter-agent, but whilst Dixon considers that its oxygen becomes affixed to the carbon monoxide molecule, liberating free hydrogen, Armstrong suggests that the oxidation of the carbon monoxide by the oxygen of the water is dependent upon the simultaneous oxidation of the hydrogen of the water by the free oxygen. Thus the states before and after the explosion may be represented by the schemes:



Von Wartenberg and Sieg strongly support Dixon's theory, and suggest the following scheme:

(1) The formation of formic acid -

CO + H₂O = H.COOH.

The production of this acid as an intermediate product during the Combustion of Carbon Monoxide was demonstrated by Wieland in 1912. The flame of the burning gas was allowed to impinge upon ice, and formic acid was found in solution in the water.

(2) The thermal decomposition of formic acid into carbon dioxide and hydrogen, both of which gases can be detected:

H.COOH = CO₂ + H₂.

This is followed by -

(3) The production of hydrogen peroxide:

H₂ + O₂ = H₂O₂,

and (4) Decomposition into water and oxygen:

H₂O₂ = H₂O + O.

Bone has still further developed Dixon's theory. He points out that the flame of hydrogen burning in air is smaller and "sharper " than a flame of carbon monoxide, burning at the same orifice and under the same pressure. The lambent character of the latter flame suggests a slower burning gas than hydrogen. Again, hydrogen-air mixtures have lower ignition temperatures than CO-air mixtures, whilst the maximum flame speed of the former is more than eight times that of the latter. Finally, Bone and Haward have shown that when corresponding H₂-air and CO-air mixtures are exploded in closed vessels, the pressure rapidly rises to a maximum in the case of hydrogen, but much more slowly with carbon monoxide. The addition of only 1 per cent, of hydrogen to the CO-air mixture very greatly accelerates the attainment of maximum pressure. All of these points indicate that carbon monoxide is not capable of being oxidised so readily as hydrogen under these conditions.

Bone therefore suggests that oxygen in flames is capable of functioning in two distinct ways, namely (i) as the undissociated molecule O₂, and (ii) as dissociated atomic O.

An undissociated molecule, on being raised to a sufficiently high temperature, is presumed capable of exerting its latent valencies, and of combining with two hydrogen molecules without itself becoming disrupted. Thus the unstable complex H₄O₂ or



is momentarily formed. This, however, instantly breaks down, yielding in part its constituent elements, in the form of hydrogen molecules and atomic oxygen; and in part as nascent or activated steam molecules. Thus

100H₄O₂ = 2nH₂O: + 2(100 – n)(O: + H₂).

The magnitude of the ratio n/(100 - n) will obviously depend upon the conditions prevailing at the moment. The higher the temperature and the smaller the hydrogen concentration, the lower is the value for n.

Oxygen and carbon monoxide, however, are regarded as incapable of associating in the above manner, their molecules being mutually inert in flames. Before the carbon monoxide can undergo oxidation, the oxygen must either have dissociated, or be in the form of some activated or nascent compound such as the activated steam indicated above. Thus, whilst the reaction

2CO + O₂ = 2CO₂

is impossible, carbon dioxide can readily be obtained in either of the following ways:

CO + :O = CO₂
CO + :OH₂ = CO₂ + H₂.

The foregoing theory receives indirect support from several earlier researches. For example, Russell1 caused nascent carbon monoxide and nascent oxygen to come into contact. This was effected by exploding - either with a spark or by heating in an air bath - a mixture of chlorine peroxide and carbonyl sulphide. It was found that the explosion exerted very considerable influence in bringing about the combination of carbon monoxide and oxygen. Russell was unable to decide whether the effect was direct or "due to the heightening of the action of the 'third substance.' " But if Bone's theory is correct, a very plausible explanation is clearly to hand for these observations. More recently Langmuir has directed attention to the fact that oxygen when brought into contact with carbon monoxide adsorbed on platinum rapidly oxidises it to carbon dioxide. Conversely, free carbon monoxide immediately reacts with adsorbed oxygen. All of these results are suggestive of acceleration of oxidation involving an "activated" condition.