(c) Heats of Solution.—When substances readily combine with water to form hydrates, the heat of solution in water is usually positive; when, on the other hand, they do not readily form hydrates, or when they are already hydrated, the heat of solution is usually negative. The following examples show the effect of hydration on heat of solution in a large quantity of water : First and Second Laws of Thermodynamics.—It should be clearly realized that the law of conservation of energy, like the law of conservation of mass, is a statement of experimental observa tions and in chemistry admits of no exception. All recorded observations support the law within the unavoidable errors of experiment ; all deductions from the law have been amply verified. Apparent limitations of the law have been brought to light by re cent physical research, and in particular through Einstein's famous theory of relativity (q.v.), but in ordinary chemistry we do not deal with atomic or ultra-atomic changes, but with molecular rearrangements, and the laws of conservation of energy and mass, as originally conceived, still hold good.
Reversible Chemical Reactions.—As the result of the large number of experiments made by them, Thomsen and, later, Berthelot expressed the opinion that heat of reaction must be regarded as a measure of chemical affinity, and that every chem ical change tended to take a course which evolved the maximum amount of heat. Berthelot in particular regarded this as a general guiding principle in chemistry, and impressed its importance so much on others that it continued to be upheld long after Berthelot himself had recognized its error. It is not easy now to realize why this "principle" commanded such universal acceptance; sim ple physical transformations which took place spontaneously either with absorption or evolution of heat, according to the cir cumstances, must have been familiar. For instance, water will freeze spontaneously and in doing so will liberate thermal energy if the temperature is below o°. Above o° ice melts spontaneously and in doing so absorbs heat from its surroundings. Further, even at the time Berthelot put forward his theory, there was a growing recognition of the fact that many chemical reactions are revers ible, i.e., that they take place either in one direction (with evolu tion of heat) or in the opposite direction (with absorption of heat.), according to the conditions of temperature and pressure and the relative masses of the reacting substances. So long ago as 1801, the great French chemist Berthollet had introduced the idea of "chemical equilibrium," and in the same year that Berthelot put forward his erroneous theory (1867) Guldberg and Waage pub lished their famous book on chemical affinity, which dealt with reversible chemical reactions and in which they put forward the law of chemical mass action.
Chemical Equilibrium.—This law, in so far as chemical equilib rium is concerned, may be formulated as follows. Suppose chemi cal substances, etc., react to form B2, etc. The reaction
tends to proceed until a condition of equilibrium is set up, gov erned by the equation where [A1], etc., represents the concentration of the compound A1, etc., in the equilibrium mixture. On the other hand, if the compounds and B2 are originally brought together, then the reaction will proceed in the reverse direction with production of and A2 until the same condition of equilibrium is set up. The value of the equilibrium constant K is itself independent of con centration—f or example, in the case of gaseous reactions it is independent of pressure—but it varies with the temperature. Clearly such constants are of great significance in the study of chemical affinities. If K is very small, then the affinity of for etc., may be said to be very great, and the affinity of for B2, etc., small, although not zero.
The important thing to notice is that every single instance of a reversible chemical reaction really disproves the Berthelot prin ciple. Nevertheless, the principle has this much to be said for it, that most chemical reactions which take place spontaneously at low temperatures do so with evolution of heat, and that it is only when the effects of high temperatures are studied that we become really familiar with important reactions that take place sponta neously with absorption of heat. There is one such reaction that is made use of daily on an enormous scale for the production of heat, namely, the reaction between steam and coke at a bright red heat to yield "water gas," a mixture of approximately equal volumes of hydrogen and carbon monoxide. As the reaction pro ceeds the coke cools, and at regular intervals it is necessary to shut off the supply of steam, and to pass air over the coke in order to raise it again by combustion to a temperature sufficiently high to allow the reaction with steam to continue rapidly. It is true that the cooling of the coke during the passage of steam is not entirely due to the chemical reaction, for it is partly accounted for by the cooling effect of the steam itself, which is introduced to the retort at a lower temperature than the coke. But even if the steam were initially at the same temperature as the coke, it would be found that the mass gradually cooled. The spontaneous formation of nitric oxide by the passage of air through an electric arc, a process by which nitric acid is manufactured on a large scale, is also accompanied by absorption of heat. Pure nitric oxide, on the other hand, decomposes into its elements, nitrogen and oxygen, at a somewhat lower temperature with evolution of heat. Hydrogen and oxygen will combine explosively and completely at low temperatures ; but steam is found to decompose spontaneously though not completely into hydrogen and oxygen if it is heated to a temperature above 2,000° C, and heat is absorbed during the decomposition. Clearly, therefore, the heat evolved during a chem ical reaction cannot give a direct measure of the affinity of the reacting substances.