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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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