For a reaction at equilibrium, it is useful to be able to write an equilibrium constant expression which relates the concentrations of the species when the system reaches equilibrium

For a generic reaction

aA + bB < - > cC + dD
the equilibrium constant expression has the form
K = [C]c[D]d/[A]a[B]b
where K is the equilibrium constant for the reaction. K is independent of the original concentrations. For gaseous reactions, we can also use the partial pressures of the gases in the reaction instead of the molar concentrations- this equilibrium constant is denoted Kp, and is related to but not the same as the K shown above. Solids and liquids do not figure into the equilibrium constant expression as a rule: see the page on heterogeneous equilibrium.

Although this expression may look similar to the rate constant expression, especially the use of captial K for the equilibrium constant and little k for the rate constant, do not confuse the two- they are not related.

The equilibrium constant expression has a number of special properties.

Example: What is the equilibrium constant expression for the gas phase reaction

N2O4(g) < - > 2NO2(g)

Solution: This reaction only has two species. If we check the generic reaction above, we will see that the expression should be

K = [NO2]2/[N2O4]
We can also write a gas phase version using the pressures
Kp = (PNO2)2/PN2O4

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