The equilibrium constant does not depend on the concentrations, but does depend on the temperature. The temperature dependence of the equilibrium constant is described by the Van‘t Hoff reaction isobar derived from thermodynamic laws:
It can be seen that with increasing temperature (dT> 0) with negative ΔH, d (InK) also becomes negative. This means:
- for exothermic reactions, the equilibrium shifts to the left to the starting materials when the temperature rises.
- Correspondingly, the equilibrium of endothermic reactions shifts to the right towards the products with increasing temperature.
Both the concentration as well as the pressure and temperature dependency of equilibria are qualitatively described jointly by the aforementioned Le Chatelier principle. If the compulsion is to increase a concentration, for example, the system reacts to this by attempting to lower the increased concentration. Increasing the pressure in a gas equilibrium leads to a shift towards the lower volume. The addition of temperature leads to an endothermic reaction; the system then tries to use the supplied heat.
The principle of Le Chatelier may be of considerable importance for fundamental, more natural-philosophical considerations. For quantitative considerations, the law of mass action (formulated twenty years earlier) for the description of concentration and pressure dependence and the Van‘t Hoff reaction isobars for the temperature dependence must be used. For practical application, the principle has more of the benefit of an orientation rule.