Module 7

Lesson 3.1   Concentration Changes in Equilibrium Systems



Key Concepts


According to the Le Chatelier's Principle, increasing the concentration of one or more reactants in a system at equilibrium favours the forward reaction, causing the system to undergo an equilibrium shift to the right. By shifting to the right, the system attempts to reduce the concentration of reactants by favouring the formation of products. This period of change ends with the establishment of a new equilibrium. Kc does not change.


Alternatively, increasing the concentration of one or more products in a system at equilibrium favours the reverse reaction, causing the system to undergo an equilibrium shift to the left. By shifting to the left, the system attempts to reduce the concentration of products by favouring the formation of reactants. Again, this period of change ends with the establishment of a new equilibrium. Kc does not change.


Example

Consider the following reaction at equilibrium:

\( \mathrm { CCl_4(l) + 2HF (g) \leftrightharpoons CCl_2F_2(g) + 2HCl(g) } \)

If more HF(g) is added to the system, the equilibrium will be disturbed. To consume some of the added HF(g), the system will shift to the right. As a result, more product will be formed and a new equilibrium will be established. However, the imposed change will be only partially counteracted. For this reason, the new equilibrium concentrations of the reactants and products will differ from the original equilibrium concentrations.

See the figure below for a graphic representation of this equilibrium shift.


Fig. 1


The dotted line represents the time that the original equilibrium shift is disturbed. Notice that the concentration of the HF(g) jumps sharply at this point. The concentration of the HF(g) then begins to fall as it reacts to form products. The concentration of the products increases in response. As time passes, the concentrations of the HF(g), CCl2F2(g), and 2HCl(g) become constant again, as evidenced by the flattening of the curves. At this point, a new equilibrium has been established. Note that concentration changes do not affect the value of Kc.


Read pages 691-692 in the textbook.


The collision-reaction theory can be used to explain how concentration changes cause equilibrium shifts.

If the number of reactant entities is increased in a given volume, the number of collisions becomes much more frequent in the container. This increases the rate of the forward reaction.



Fig. 2  Reaction rates and concentration
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Initially, the rate of the reverse reaction remains the same so there is an increase in products. Of course, as the concentration of products increases, the rate of the reverse reaction will begin to accelerate. Meanwhile, the rate of the forward reaction starts to slow because the extra reactant is being consumed. Over time, the rates will again become equal. However, the rates of the reactions at the new equilibrium will be faster than they were at the original equilibrium because the system contains an increased number of particles.