Editing Acid dissociation constant (section)

=== Temperature dependence ===
All equilibrium constants vary with [[temperature]] according to the [[van 't Hoff equation]]<ref>
{{cite book
| title = Physical Chemistry
| last1 = Atkins
| first1 = P.W.
| last2 = de Paula
| first2 = J.
| year = 2006
| publisher = Oxford University Press
| isbn = 0-19-870072-5
}} Section 7.4: The Response of Equilibria to Temperature</ref>
:<math alt="The derivative of the natural logarithm of any equilibrium constant K with respect to the [[absolute temperature]] T equals the standard enthalpy change for the reaction divided by the product R times T squared. Here R represents the gas constant, which equals the thermal energy per mole per kelvin. The standard enthalpy is written as Delta H with a superscript plimsoll mark represented by the image strikeO. This equation follows from the definition of the Gibbs energy Delta G equals R times T times the natural logarithm of K.">
 \frac{\mathrm{d} \ln\left(K\right)}{\mathrm{d}T} = \frac{\Delta H^\ominus}{RT^2}
</math>

{{tmath|R}} is the [[gas constant]] and {{tmath|T}} is the [[kelvin|absolute temperature]]. Thus, for [[exothermic]] reactions, the standard [[enthalpy change]], {{tmath|\Delta H^\ominus}}, is negative and ''K'' decreases with temperature. For [[endothermic]] reactions, {{tmath|\Delta H^\ominus}} is positive and ''K'' increases with temperature.

The standard enthalpy change for a reaction is itself a function of temperature, according to [[Gustav Kirchhoff#Kirchhoff's law of thermochemistry|Kirchhoff's law of thermochemistry]]:
:<math>\left(\frac{\partial\Delta H}{\partial T}\right)_p = \Delta C_p</math>
where {{tmath|\Delta C_p}} is the [[Specific heat capacity|heat capacity]] change at constant pressure. In practice {{tmath|\Delta H^\ominus}} may be taken to be constant over a small temperature range.