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The two related definitions of the Law of Mass Action are:

  1. At constant temperature the product of active masses on one side of a chemical equation, when divided by the product of active masses on the other side of the chemical equation, is a constant, regardless of the amounts of each substance present at the beginning of the action.

  2. At constant temperature the rate of reaction is proportional to the concentration of each kind of substance taking part in the reaction [Weast (1982)].

The second definition is often used in stating reaction rate laws for single, homogeneous, unidirectional chemical reactions (e.g., rate = k CA CB, where k is the rate constant and C denotes concentrations). The rates of forward and reverse chemical reactions in a reversible system can be expressed in similar terms, leading to the first definition above.

The Law of Mass Action has been described as a restrictive condition expressing the indestructibility of matter during chemical reactions [Rushbrook (1962)]. There are relatively few references in modern texts to this law, which relates to the generally accepted indestructibility of matter under conditions of changing equilibria. Considering the general homogeneous, reversible chemical reaction

the Law of Mass Action requires the rate of forward reaction to be proportional to the concentrations of A and B. Defining these as CA and CB, respectively, the rate of the forward reaction can be written as

Similarly, the rate of the reverse reaction is expressed as

where CC and CD denote the concentrations of species C and D, respectively. In these equations, kf and kr denote the forward and reverse (temperature dependent, concentration independent) reaction rate constants; the temperature dependence of kf and kr is described by the Arrhenius equation.

At equilibrium, rforward = rreverse and

which can be used to define the equilibrium constant for this particular reaction:

More generally, if the chemical reaction has stoichiometric coefficients a, b, c and d such that

at equilibrium

[Moore (1963)]. The concept finds ready application in mass action equilibria of sorption and ion exchange processes [Perry and Chilton (1984)].

REFERENCES

Atkins, P. W. (1986) Physical Chemistry. OUP. Oxford. UK.

Moore, W. J. (1963) Physical Chemistry. Prentice Hall. Englewood Cliffs. New Jersey.

Perry, R. H. and Chilton, C. H. (Eds.). (1984) Chemical Engineers' Handbook. McGraw-Hill. New York. USA.

Rushbrooke, G. S. (1962) OUP. Oxford, UK.

Weast, (1982) (63rd Edn.) Handbook of Chemistry and Physics. CRC Press. Boca Raton. Florida. USA.

References

  1. Atkins, P. W. (1986) Physical Chemistry. OUP. Oxford. UK.
  2. Moore, W. J. (1963) Physical Chemistry. Prentice Hall. Englewood Cliffs. New Jersey.
  3. Perry, R. H. and Chilton, C. H. (Eds.). (1984) Chemical Engineers' Handbook. McGraw-Hill. New York. USA.
  4. Rushbrooke, G. S. (1962) OUP. Oxford, UK.
  5. Weast, (1982) (63rd Edn.) Handbook of Chemistry and Physics. CRC Press. Boca Raton. Florida. USA.
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