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If a reaction occurs in more than one phase, the reaction takes place in a
heterogeneous system. In general, we write equilibrium
constant expressions in terms of molar concentration or gas phase pressures.
For pure liquids or solids, experiments have shown that:
- The position of the equilibrium state of a system does not depend on the
amount of liquid or solid in the reaction, provided that some exists.
- Pure liquids and solids do not appear in the equilibrium expression.
To understand why this is true, consider a simple reaction like the
vaporization of water:
- H2O(l) < - > H2O(g)
At a given
temperature, the vapor
pressure of water is a constant no matter how much liquid water there is.
Thus, the position of the equilibrium does not change when you add or subtract
water, and so the equilibrium expression will have no term for liquid water.
This example helps explain why the term for the solvent never shows up in
equilibrium constant expressions for reactions that are done in that solvent.
Only the solutes appear.
Example: What is the equilibrium constant expression for the following
reaction?
- CO2(g) + H2(g) < - > CO(g) +
H2O(l)
Solution: Write the equilibrium constant expression the normal way for
gases, in terms of the partial pressures of each gas. The liquid water on the
right hand side of the equation will not have a term in the expression, so the
answer is
- K =
PCO/PCO2PH2
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