The equilibrium constant helps in predicting the direction in which a given reaction will proceed at any stage. For this purpose, we calculate the reaction quotient Q.
The reaction quotient, Q (QC with molar concentrations and QP with partial pressures) is defined in the same way as the equilibrium constant Kc except that the concentrations in QC are not necessarily equilibrium values.
For a general reaction:
a A + b B ⇔ c C + d D
Qc = [C]c[D]d/ [A]a[B]b
If Qc > Kc, the reaction will proceed in the direction of reactants (reverse reaction).
If Qc < Kc, the reaction will proceed in the direction of the products (forward reaction).
If Qc = Kc, the reaction mixture is already at equilibrium.
Consider the gaseous reaction of H2 with I2,
H2(g) + I2(g) ⇔ 2HI(g) Kc = 57.0 at 700 K.
Suppose we have molar concentrations [H2]t=0.10M, [I2]t = 0.20 M and [HI]t = 0.40 M. (the subscript t on the concentration symbols means that the concentrations were measured at some arbitrary time t, not necessarily at equilibrium).
Thus, the reaction quotient, Qc at this stage of the reaction is given by,
Qc = [HI]t2 / [H2]t [I2]t = (0.40)2/ (0.10) × (0.20) = 8.0
Now, in this case, Qc (8.0) does not equal Kc (57.0), so the mixture of H2(g), I2(g) and HI(g) is not at equilibrium; that is, more H2(g) and I2(g) will react to form more HI(g) and their concentrations will decrease till Qc = Kc.
The reaction quotient, Qc is useful in predicting the direction of reaction by comparing the values of Qc and Kc. Thus, we can make the following generalisations concerning the direction of the reaction.
- If Qc < Kc, net reaction goes from left to right
- If Qc > Kc, net reaction goes from right to left.
- If Qc = Kc, no net reaction occurs