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An activity coefficient is a factor used in thermodynamics to account for deviations from ideal behaviour in a mixture of chemical substances.1 In an ideal mixture the interactions between each pair of chemical species are the same (or more formally, the enthalpy of mixing is zero) and, as a result, properties of the mixtures can be expressed directly in terms of simple concentrations or partial pressures of the substances present e.g. Raoult's law. Deviations from ideality are accommodated by modifying the concentration by an activity coefficient. Analogously, expressions involved gases can be adjusted for non-ideality by scaling partial pressures by a fugacity coefficient.
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Thermodynamics
The chemical potential, μB, of a substance B in an ideal mixture is given by
where 
is the chemical potential in the standard state and xB is the mole fraction of the substance in the mixture.
This is generalised to include non-ideal behaviour by writing
when aB is the activity of the substance in the mixture with
- aB = xBγB
where γB is the activity coefficient. As xB approaches 1, the substance behaves as if it were ideal, and thus 
, which is known as Raoult's Law. For γB > 1 and γB < 1 substance B shows positive and negative deviation from Raoult's law respectively. A positive deviation implies that the substance becomes more volatile.
For the case where xB goes to zero, the activity coefficient of substance B approaches a constant; this relationship is Henry's Law for the solvent. These relationships are related to each other through the Gibbs-Duhem equation 2. Note that in general activity coefficients are dimensionless.
Modifying mole fractions or concentrations by activity coefficients gives the effective activities of the components, and hence allows expressions such as Raoult's law and equilibrium constants constants to be applied to both ideal and non-ideal mixtures.
Knowledge of activity coefficients is particularly important in the context of electrochemistry since the behaviour of electrolyte solutions is often far from ideal, due the effects of the ionic atmosphere.
Application to chemical equilibrium
At equilibrium, the sum of the chemical potentials of the reactants is equal to the sum of the chemical potentials of the products. The Gibbs free energy change for the reactions, ΔG, is equal to the difference between these sums and therefore, at equilibrium, is equal to zero. Thus, for an equilibrium such as
Substitute in the expressions for the chemical potential of each reactant:
Upon rearrangement this expression becomes
The sum 
is the standard free energy change for the reaction, ΔGO. Therefore
K is the equilibrium constant. Note that activities and equilibrium constants are dimensionless numbers.
This derivation serves two purposes. It shows the relationship between standard free energy change and equilibrium constant. It also shows than an equilibrium constant is defined as a quotient of activities. In practical terms this is inconvenient. When each activity is replaced by the product of a concentration and an activity coefficient, the equilibrium constant is defined as
![K= \frac{[S]^\sigma[T]^\tau}{[A]^\alpha[B]^\beta} \times \frac{\gamma_S^\sigma \gamma_T^\tau}{\gamma_A^\alpha \gamma_B^\beta}](http://upload.wikimedia.org/math/4/8/b/48b1a5305731008ed0b72d58a9184efc.png)
where [S] denotes the concentration of S, etc. In practice equilibrium constants are determined in a medium such that the quotient of activity coefficient is constant and can be ignored, leading to the usual expression
![K= \frac{[S]^\sigma[T]^\tau}{[A]^\alpha[B]^\beta}](http://upload.wikimedia.org/math/e/e/f/eef75b6fdacc80806a69add59e6f4110.png)
which applies under the conditions that the activity quotient has a particular (constant) value.
Measurement and prediction of activity coefficients
Activity coefficients may be measured experimentally or calculated theoretically, using the Debye-Hückel equation or extensions such as Davies equation3 or Pitzer equations4. Specific Ion Theory (SIT) 5 may also be used. Alternatively correlative methods such as UNIQUAC. NRTL or UNIFAC may be employed, provided fitted component-specific or model parameters are available.
A new alternative for activity coefficients prediction, which is less dependent on model parameters, are the COSMO-based methods. In these methods the required information comes from quantum mechanics calculations specific to each molecule (sigma profiles)6.
For uncharged species, the activity coefficient γ0 mostly follows a "salting-out" model7:
- log10(γ0) = bI
This simple model predicts activities of many species (dissolved undissociated gases such as CO2, H2S, NH3, undissociated acids and bases) to high ionic strengths (up to 5 mol/kg). The value of the constant b for CO2 is 0.11 at 10 °C and 0.20 at 330 °C89.
For water (solvent), the activity aw can be calculated using7:
φ
where ν is the number of ions produced from the dissociation of one molecule of the dissolved salt, m is the molal concentration of the salt dissolved in water, φ is the osmotic coefficient of water, and the constant 55.51 represents the molal concentration of water. In the above equation, the activity of a solvent (here water) is represented as inversely proportional to the number of particles of salt versus that of the solvent.
References
- ^ International Union of Pure and Applied Chemistry. "Activity coefficient". Compendium of Chemical Terminology Internet edition.
- ^ R. DeHoff, Thermodynamic in Materials Science, Taylor and Francis, 2006. pp230-1
- ^ C.W. Davies, Ion Association,Butterworths, 1962
- ^ I. Grenthe and H. Wanner, Guidelines for the extrapolation to zero ionic strength, http://www.nea.fr/html/dbtdb/guidelines/tdb2.pdf
- ^ "Project: Ionic Strength Corrections for Stability Constants". IUPAC. http://www.iupac.org/web/ins/2000-003-1-500. Retrieved on 2008-11-15.
- ^ Ryo Kato, Jurgen Gmehling. "Systems with ionic liquids: Measurement of VLE and [gamma][infinity] data and prediction of their thermodynamic behavior using original UNIFAC, mod. UNIFAC(Do) and COSMO-RS(Ol)". The Journal of Chemical Thermodynamics 37: 603-619. doi:.
- ^ a b J.N. Butler, "Ionic Equilibrium", John Wiley and Sons, Inc., 1998.
- ^ A.J. Elis and R.M. Golding, Am. J. Sci, 162, p 47-60, 1963.
- ^ S.D.Malinin, Geokhimiya, 3, p. 235-245, 1959.
External links
Software for solution chemistry Includes SIT theory.
Wikipedia content modification information:
- This page was last modified on 28 December 2008, at 01:45.
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