Allotropy 3

The heat of the reaction is analogous to the heat of evaporation of any solid or liquid, and if P is the "vapour" pressure, which is called the dissociation pressure in this and similar cases, where Q is the heat of reaction, i.e., the heat absorbed on dis sociation of i gram-molecule, and includes the work done against atmospheric pressure by the evolution of i gram-molecule of stant being the difference between the num ber of the reacting molecules and the total number formed by the reaction: d = Qp/R V. It can also be easily shown that where Q, is now not necessarily the diminution in total energy, but corresponds to the heat evolved when the reaction takes place at constant pressure, without the performance of any work other than that due to the change of volume, if any, associated with the reaction. This equation is similar to that found to hold good for the change of vapour pressures of liquids (and solids) with the temperature, viz., d p/ dT =X/ RI'', where X is the latent heat of evaporation of one gram-molecule of the liquid or solid.

Some Thermochemical Reactions.

The application of these equations to chemistry inaugurated a new era, and led not only to a rapid extension of our knowledge of chemical reaction, but also to the recognition of the connection between many apparently diverse phenomena. Some illustrations of the use of the equations will now be given.

Hydrogen and oxygen combine to form steam with evolution of heat according to the equation Now at ordinary temperatures, the amount of hydrogen and oxygen left uncombined is very small and escapes direct measure ment. As x is small, K is also small. But van't Hoff's equation shows that, as U is large and positive, log K, and therefore K, in creases with the temperature. Hence x increases with the tempera ture. In other words, water vapour will dissociate when the temperature is raised.

This result may be generalized. All compounds which are formed with evolution of heat tend to decompose when the tern perature is raised. The same applies to molecules such as H2, 02, N2, C12, 12, etc., which are formed from their atoms with evolution of heat. Conversely, compounds which are formed with absorption of heat become more stable at high temperatures. Nitric oxide cannot be made from nitrogen and oxygen at low temperatures, but is formed when air is blown through an electric arc. Hydrogen peroxide may be detected in an oxy-hydrogen flame. Calcium carbide is formed in an electric furnace at a very high temperature. If v in the equation given above is made smaller, then x must also diminish, because K is constant. High carbon dioxide. The dissociation pressure, like the vapour pres sure of liquids, increases with the temperature. At a high tempera ture, about goo° C, the dissociation pressure is equal to ordinary atmospheric pressure. If the carbonate is heated above this temperature it is converted into lime and carbon dioxide is rapidly evolved.

If other carbonates are examined it is found that the tempera ture at which they are converted into their oxides varies directly with their heat of dissociation as the following table shows: On similar reasoning, it may be shown generally that in all reactions in which gases and solids take part, the equilibrium conditions only take account of the gases. For instance, when carbon dioxide is passed over red-hot coke it is converted into carbon monoxide with absorption of heat, C+CO2= 2C0-38,500 calories. The equilibrium condition is from which it can be deduced (a) that the ratio CO/CO2 in the gas mixture at equilibrium is constant if the temperature and pressure are constant; (b) that if the temperature is kept con stant, and the pressure is increased, the proportion of carbon monoxide in the mixture goes down; (c) that if the temperature is increased, the proportion of carbon monoxide goes up.

In the blast furnace, air is blown through a white-hot mixture of oxides of iron and coke, and the carbon monoxide formed by the combustion of coke reduces the oxides to metallic iron. The exit gases must contain a large proportion of carbon monoxide to carbon dioxide, corresponding to the equilibrium conditions which obtain near the cooler top of the furnace. It was at one time thought that by building furnaces higher, the period of contact between the gases and the ore would result in a greater measure of reduction, a diminished proportion of carbon mon oxide in the exit gases, and consequently a smaller consumption of coke per ton of ore reduced. Attempts to achieve these re sults, by building furnaces as high as ioo ft., failed. As Le Cha telier pointed out, if there had been a better acquaintance with the laws of chemical equilibrium, these experiments would have been unnecessary. It was an expensive, although no doubt an instructive, method of proving the correctness of one of the de ductions from the second law.

Heat and Velocity of Reaction.

These examples will suffice to indicate how the quantitative application of simple fundamental thermodynamic principles has been successful in reducing to order an enormous range of observations on chemical reactions apparently very diverse in character. It has also given the chemical engineer the power to calculate the conditions for achiev ing the maximum technical success in the conduct of any chemical manufacture on the large scale. Unfortunately, though the ap plication of thermodynamics has been so successful in the study of chemical equilibrium, it has as yet taught us nothing about the rate at which equilibrium is attained, or, in other words, about the rate of chemical reaction. There seems to be no direct connection between the rate of a reaction and either the heat evolved (change in total energy), or the maximum work of the reaction (change in free energy). Hydrogen has a great affinity for oxygen, and much heat is evolved when they combine ; yet a mixture of them can be kept for an indefinite period of time at the ordinary temperature without a trace of combination. Nitric oxide is an extremely unstable substance at low tempera ture; on thermodynamical grounds we should expect that only an infinitesimal amount of it would be in equilibrium with nitrogen and oxygen at the ordinary temperature ; yet it is so difficult to decompose by heat that even at 900° C it only de composes slowly. Ammonia becomes progressively less stable as the temperature is raised. Thermodynamics indicates that in order to obtain high yields, its synthesis from nitrogen and hydrogen should be conducted.at as low a temperature as pos sible. The reaction is, however, so slow that technical manu facture only becomes possible at temperatures so high that the yield of ammonia at atmospheric pressure is extremely small. The ill effects of the high temperature have to be counter balanced by the use of high pressures. Nearly all gas reactions take place extremely slowly unless the temperature is very high. Nearly all, too, take place more quickly in the presence of certain solids called catalysts. The nature of the catalyst to produce the best results varies with the nature of the reaction, and though many theories have been put forward to account for their action, there is no general theory which has been suc cessful in predicting anything—which is the ultimate test of the real value of a theory.