Editing Equilibrium constant (section)

== Unknown activity coefficient values ==
[[File:PK acetic acid.png|thumb|200px|Variation of log&nbsp;''K''<sub>c</sub> of acetic acid with ionic strength]]
It is very rare for activity coefficient values to have been determined experimentally for a system at equilibrium. There are three options for dealing with the situation where activity coefficient values are not known from experimental measurements.
#Use calculated activity coefficients, together with concentrations of reactants. For equilibria in solution estimates of the activity coefficients of charged species can be obtained using [[Debye–Hückel theory]], an extended version, or [[SIT theory]]. For uncharged species, the activity coefficient ''γ''<sub>0</sub> mostly follows a "salting-out" model: {{nowrap begin}}log<sub>10</sub> ''γ''<sub>0</sub> = ''bI''{{nowrap end}} where ''I'' stands for [[ionic strength]].<ref name = Butler>{{cite book|first=J. N. |last=Butler |title=Ionic Equilibrium |publisher=John Wiley and Sons |date=1998}}</ref>
#Assume that the activity coefficients are all equal to 1. This is acceptable when all concentrations are very low.
#For equilibria in solution use a medium of high ionic strength. In effect this redefines the [[standard state]] as referring to the medium. Activity coefficients in the standard state are, by definition, equal to 1. The value of an equilibrium constant determined in this manner is dependent on the ionic strength. When published constants refer to an ionic strength other than the one required for a particular application, they may be adjusted by means of specific ion theory (SIT) and other theories.<ref>
{{cite web |url=http://www.iupac.org/web/ins/2000-003-1-500 |title=Project: Ionic Strength Corrections for Stability Constants |publisher=International Union of Pure and Applied Chemistry |access-date=2008-11-23 |archive-url=https://web.archive.org/web/20081029193538/http://www.iupac.org/web/ins/2000-003-1-500 |archive-date=29 October 2008 |url-status=dead }}</ref>