Activity (chemistry)


Activity (chemistry)

In chemical thermodynamics, activity (symbol a) is a measure of the “effective concentration” of a species in a mixture, meaning that the species' chemical potential depends on the activity of a real solution in the same way that it would depend on concentration for an ideal solution.

By convention, activity is treated as a dimensionless quantity, although its actual value depends on customary choices of standard state for the species. The activity of pure substances in condensed phases (solid or liquids) is normally taken as unity (the number 1). Activity depends on temperature, pressure and composition of the mixture, among other things. For gases, the effective partial pressure is usually referred to as fugacity.

The difference between activity and other measures of composition arises because molecules in non-ideal gases or solutions interact with each other, either to attract or to repel each other. The activity of an ion is particularly influenced by its surroundings.

Activities should be used to define equilibrium constants but, in practice, concentrations are often used instead. The same is often true of equations for reaction rates. However, there are circumstances where the activity and the concentration are significantly different and, as such, it is not valid to approximate with concentrations where activities are required. Two examples serve to illustrate this point:

  • In a solution of potassium hydrogen iodate at 0.02 M the activity is 40% lower than the calculated hydrogen ion concentration, resulting in a much higher pH than expected.
  • When a 0.1 M hydrochloric acid solution containing methyl green indicator is added to a 5 M solution of magnesium chloride, the color of the indicator changes from green to yellow—indicating increasing acidity—when in fact the acid has been diluted. Although at low ionic strength (<0.1 M) the activity coefficient approaches unity, this coefficient can actually increase with ionic strength in a high ionic strength regime. For hydrochloric acid solutions, the minimum is around 0.4 M.[1]

Contents

Definition

The activity of a species i, denoted ai, is defined[2][3] as:

a_i = \exp\left (\frac{\mu_i - \mu^{\ominus}_i}{RT}\right )

where μi is the chemical potential of the species under the conditions of interest, μoi is the chemical potential of that species in the chosen standard state, R is the gas constant and T is the thermodynamic temperature. This definition can also be written in terms of the chemical potential:

\mu_i = \mu_i^{\ominus} + RT\ln{a_i}

Hence the activity will depend on any factor that alters the chemical potential. These include temperature, pressure, chemical environment etc. In specialised cases, other factors may have to be considered, such as the presence of an electric or magnetic field or the position in a gravitational field. However the most common use of activity is to describe the variation in chemical potential with the composition of a mixture.

The activity also depends on the choice of standard state, as it describes the difference between an actual chemical potential and a standard chemical potential. In principle, the choice of standard state is arbitrary, although there are certain conventional standard states which are usually used in different situations.

Activity coefficient

The activity coefficient γ, which is also a dimensionless quantity, relates the activity to a measured amount fraction xi, molality mi or amount concentration ci:

a_i = \gamma_{x,i} x_i\,
a_i = \gamma_{m,i} m_i/m^{\ominus}\,
a_i = \gamma_{c,i} c_i/c^{\ominus}\,

The division by the standard molality mo or the standard amount concentration co is necessary to ensure that both the activity and the activity coefficient are dimensionless, as is conventional.[3]

When the activity coefficient is close to one, the substance shows almost ideal behaviour according to Henry's law. In these cases, the activity can be substituted with the appropriate dimensionless measure of composition xi, mi/mo or ci/co. It is also possible to define an activity coefficient in terms of Raoult's law: the International Union of Pure and Applied Chemistry (IUPAC) recommends the symbol ƒ for this activity coefficient,[3] although this should not be confused with fugacity.

a_i = f_i x_i\,. Solution can also become too diluted when necessary.

Standard states

Gases

In most laboratory situations, the difference in behaviour between a real gas and an ideal gas is dependent only on the pressure and the temperature, not on the presence of any other gases. At a given temperature, the "effective" pressure of a gas i is given by its fugacity ƒi: this may be higher or lower than its mechanical pressure. By historical convention, fugacities have the dimension of pressure, so the dimensionless activity is given by:

a_i = \frac{f_i}{p^{\ominus}} = \phi_i x_i \frac{p}{p^{\ominus}}

where φi is the dimensionless fugacity coefficient of the species, xi is its fraction in the gaseous mixture (x = 1 for a pure gas) and p is the total pressure. The value po is the standard pressure: it may be equal to 1 atm (101.325 kPa) or 1 bar (100 kPa) depending on the source of data, and should always be quoted.

Mixtures in general

The most convenient way of expressing the composition of a generic mixture is by using the amount fractions x of the different components, where

x_i = \frac{n_i}{n}
\sum_i x_i = 1\,

The standard state of each component in the mixture is taken to be the pure substance, i.e. the pure substance has an activity of one. When activity coefficients are used, they are usually defined in terms of Raoult's law,

a_i = f_i x_i\,

where ƒi is the Raoult's law activity coefficient: an activity coefficient of one indicates ideal behaviour according to Raoult's law.

Dilute solutions (non-ionic)

A solute in dilute solution usually follows Henry's law rather than Raoult's law, and it is more usual to express the composition of the solution in terms of the amount concentration c (in mol/L) or the molality m (in mol/kg) of the solute rather than in amount fractions. The standard state of a dilute solution is a hypothetical solution of concentration co = 1 mol/L (or molality mo = 1 mol/kg) which shows ideal behaviour (also referred to as "infinite-dilution" behaviour). The standard state, and hence the activity, depends on which measure of composition is used. Molalities are often preferred as the volumes of non-ideal mixtures are not strictly additive and are also temperature-dependent: molalities do not depend on volume, whereas amount concentrations do.[4]

The activity of the solute is given by:

a_{c,i} = \gamma_{c,i}\, \frac{c_i}{c^{\ominus}}
a_{m,i} = \gamma_{m,i}\, \frac{m_i}{m^{\ominus}}

Ionic solutions

When the solute undergoes ionic dissociation in solution (a salt e.g.), the system becomes decidedly non-ideal and we need to take the dissociation process into consideration. We can define activities for the cations and anions separately (a+ and a).

It should be noted however that in a liquid solution the activity coefficient of a given ion (e.g. Ca2+) isn't measurable because it is experimentally impossible to independently measure the electrochemical potential of an ion in solution. (We cannot add cations without putting in anions at the same time). Therefore one introduces the notions of

mean ionic activity
a±ν = a+ν+aν
mean ionic molality
m±ν = m+ν+mν
mean ionic activity coefficient
γ±ν = γ+ν+γν

where ν = ν+ + ν represent the stoichiometric coefficients involved in the ionic dissociation process

Even though γ+ and γ cannot be determined separately, γ± is a measureable quantity that can also be predicted for sufficiently dilute systems using Debye–Hückel theory. For the activity of a strong ionic solute (complete dissociation) we can write:

a2 = a±ν = γ±νm±ν

Measurement

The most direct way of measuring an activity of a species is to measure its partial vapor pressure in equilibrium with a number of solutions of different strength. For some solutes this is not practical, say sucrose or salt (NaCl) do not have a measurable vapor pressure at ordinary temperatures. However, in such cases it is possible to measure the vapor pressure of the solvent instead. Using the Gibbs–Duhem relation it is possible to translate the change in solvent vapor pressures with concentration into activities for the solute.

Another way to determine the activity of a species is through the manipulation of colligative properties, specifically freezing point depression. Using freezing point depression techniques, it is possible to calculate the activity of a weak acid from the relation,

m^{\prime} = m(1 + a)\,

where m' is the total molal equilibrium concentration of solute determined by any colligative property measurement(in this case ΔTfus, m is the nominal molality obtained from titration and a is the activity of the species.

There are also electrochemical methods that allow the determination of activity and its coefficient.

The value of the mean ionic activity coefficient γ± of ions in solution can also be estimated with the Debye–Hückel equation, the Davies equation or the Pitzer equations.

Use

Chemical activities should be used to define chemical potentials, where the chemical potential depends on the temperature T, pressure p and the activity ai according to the formula:

\mu_i = \mu_i^{\ominus} + RT\ln{a_i}

where R is the gas constant and µio is the value of µi under standard conditions.

Formulae involving activities can be simplified by considering that:

  • For a chemical solution:
    • the solvent has an activity of unity
    • At a low concentration, the activity of a solute can be approximated to the ratio of its concentration over the standard concentration:
a_i = \frac{c_i}{c^{\ominus}}

Therefore, it is approximately equal to its concentration.

  • For a mix of gas at low pressure, the activity is equal to the ratio of the partial pressure of the gas over the standard pressure:
a_i = \frac{p_i}{p^{\ominus}}
Therefore, it is equal to the partial pressure in bars (compared to a standard pressure of 1 bar).
  • For a solid body, a uniform, single species solid at one bar has an activity of unity. The same thing holds for a pure liquid.

The latter follows from any definition based on Raoult's law, because if we let the solute concentration x1 go to zero, the vapor pressure of the solvent p will go to p*. Thus its activity a = p/p* will go to unity. This means that if during a reaction in dilute solution more solvent is generated (the reaction produces water e.g.) we can typically set its activity to unity.

Solid and liquid activities do not depend very strongly on pressure because their molar volumes are typically small. Graphite at 100 bars has an activity of only 1.01 if we choose po = 1 bar as standard state. Only at very high pressures do we need to worry about such changes.Changes can also come as a result of too much dilution of solution.

Example values

Example values of activity coefficients of sodium chloride in aqueous solution are given in the table.[5] In an ideal solution, these values would all be unity. The deviations tend to become larger with increasing molality and temperature, but with some exceptions.

Molality (mol/kg) 25 °C 50 °C 100 °C 200 °C 300 °C 350 °C
0.05 0.820 0.814 0.794 0.725 0.592 0.473
0.50 0.680 0.675 0.644 0.619 0.322 0.182
2.00 0.669 0.675 0.641 0.450 0.212 0.074
5.00 0.873 0.886 0.803 0.466 0.167 0.044

See also

References

  1. ^ McCarty, Christopher G.; Vitz, Ed (2006), "pH Paradoxes: Demonstrating that it is not true that pH ≡ -log[H+]", J. Chem. Ed. 83: 752, doi:10.1021/ed083p752 
  2. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "activity (relative activity), a".
  3. ^ a b c International Union of Pure and Applied Chemistry (1993). Quantities, Units and Symbols in Physical Chemistry, 2nd edition, Oxford: Blackwell Science. ISBN 0-632-03583-8. pp. 49–50. Electronic version.
  4. ^ Kaufman, Myron (2002), Principles of thermodynamics, CRC Press, p. 213, ISBN 0-8247-0692-7 
  5. ^ Cohen, Paul (1988), The ASME Handbook on Water Technology for Thermal Systems, American Society of Mechanical Engineers, p. 567, ISBN 0-7918-0300-7 

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