and pressure ranges. The standard Gibbs energy (for each species or for the entire reaction) can be represented (from the basic definitions) as: In the above equation, the effect of temperature on Gibbs energy (and thus on the equilibrium constant) is ascribed entirely to heat capacity. To evaluate the integrals in this equation, the form of the dependence of heat capacity on temperature needs to be known. If the standard molar heat capacity C can be approximated by some analytic function of temperature (e.g. the Shomate equation), then the integrals involved in calculating other parameters may be solved to yield analytic expressions for them. For example, using approximations of the following forms: For pure substances (solids, gas, liquid): For ionic species at : then the integrals can be evaluated and the following final form is obtained: The constants A, B, C, a, b and the absolute entropy, S̆, required for evaluation of C(T), as well as the values of G298 K and S298 K for many species are tabulated in the literature. Pressure dependence The pressure dependence of the equilibrium constant is usually weak in the range of pressures normally encountered in industry, and therefore, it is usually neglected in practice. This is true for condensed reactant/products (i.e., when reactants and products are solids or liquid) as well as gaseous ones. For a gaseous-reaction example, one may consider the well-studied reaction of hydrogen with nitrogen to produce ammonia: N2 + 3 H2 2 NH3 If the pressure is increased by the addition of an inert gas, then neither the composition at equilibrium nor the equilibrium constant are appreciably affected (because the partial pressures remain constant, assuming an ideal-gas behaviour of all gases involved). However, the composition at equilibrium will depend appreciably on pressure when: the pressure is changed by compression or expansion of the gaseous reacting system, and the reaction results in the change of the number of moles of gas in the system. In the example reaction above, the number of moles changes from 4 to 2, and an increase of pressure by system compression will result in appreciably more ammonia in the equilibrium mixture. In the general case of a gaseous reaction: α A + β B σ S + τ T the change of mixture composition with pressure can be quantified using: where p denote the partial pressures and X the mole fractions of the components, P