D30.2 Le Châtelier’s principle: Change in Concentration

Using equilibrium constant, reaction quotient, and ICE table, we have a quantitative way to analyze how any system that is out of equilibrium can establish (or re-establish) equilibrium, such as the case for when an acid or a base is added to a buffer system. Alternatively, we can qualitatively evaluate changes in a chemical equilibrium by applying Le Châtelier’s principle.

Le Châtelier’s principle states that when a chemical system is at equilibrium and conditions are changed so that the reaction quotient, Q, changes, the chemical system will react to achieve new equilibrium concentrations or partial pressures; reaction occurs in a way that partially counteracts the change in conditions. To establish the new equilibrium, the reaction proceeds in the forward direction if Q < K or in the reverse direction if Q > K, until Q is again equal to K.

For a chemical system at equilibrium at constant temperature, if the concentration of a reactant or a product is changed, therefore changing Q, the system is no longer at equilibrium because QK. The concentrations of all reaction species will then undergo additional changes until the system reaches a new equilibrium with a different set of equilibrium concentrations. We say that the equilibrium shifts in a direction (forward or reverse) that partially counteracts the change.

For example, consider the chemical reaction:

H2(g) + I2(g) ⇌ 2HI(g)          Kc = 50.0 at 400 °C

A mixture of gases at 400 °C with [H2] = 0.221 M, [I2] = 0.290 M, and [HI] = 1.790 M is at equilibrium in a closed container. For this mixture, Qc = Kc = 50.0.

If additional H2(g) is introduced into the container quickly such that [H2] doubles before it begins to react (that is, the new [H2]t = 0.442 M), Qc is now ½ of Kc:

 Q_c = \dfrac{[\text{HI}]_t^{\;2}}{[\text{H}_2]_t[\text{I}_2]_t} = \dfrac{(1.790 M)^2}{(0.442 M)(0.290 M)} = 25.0 = \frac{1}{2}K_c

The reaction will proceed in the forward direction (towards products) to reach a new equilibrium. You can use an ICE table and calculate that the new equilibrium concentrations are [H2] = 0.362 M, [I2] = 0.210 M, and [HI] = 1.950 M.

Notice that [H2]new equilibrium (0.362 M) is less than the doubled concentration (0.442 M) but more than [H2]first equilibrium (0.221 M). The equilibrium has shifted to partially counteract the change in H2 concentration. Because of the shift, the concentration of the other reactant decreases and the concentration of the product increases. To verify that these new concentrations are equilibrium concentrations, calculate Q:

 Q_c = \dfrac{(1.950 M)^2}{(0.362 M)(0.210 M)} = 50.0 = K_c

The figure below illustrates graphically the effect of adding H2 to the reaction that was at equilibrium.

Figure: Change in concentration. A mixture of H2(g) and I2(g) reacts to reach the equilibrium of H2(g) + I2(g) ⇌ 2HI(g). Then H2(g) is added (at dashed vertical line). The reaction proceeds to form more products until a new equilibrium is achieved.


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Chemistry 109 Fall 2021 by John Moore, Jia Zhou, and Etienne Garand is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.