What are the applications of the Equilibrium Constant?

Applications of the equilibrium constant

1. The magnitude of the equilibrium constant :
The magnitude of the equilibrium constant K means the extent to which a reaction can proceed i.e. measuring the completion of the reversible reaction.

If the value of K is larger, then the equilibrium concentration of the components on the right hand side of the reaction will be greater than the components on the left hand side of the reaction. Hence the reaction proceeds to a greater extent. For example:
Consider the following reaction:
2 NO₂ (g) ⇌ N₂ (g) +2 O₂ (g)
The equilibrium constant for the reaction at 298 K is:
K = [N₂] [O₂] ²/[NO₂] ²
= 6.7 X 10^-16 mol L^-1
The value of K is very small which means that the molar concentrations of N₂ and O₂ in the equilibrium mixture has proceeded to small extent only. Hence, we can conclude that NO₂ is quite stable and decomposes slightly.

2. Predicting the direction of the reactions:
The equilibrium constant also helps in finding the direction in which reaction proceeds. For this, we have to calculate the reaction quotient Q.

The reaction quotient Q is the ratio of the product of concentrations of the products to that of the reactants. For example:
Consider the following reaction:
A + B ⇌ X + Y
The reaction quotient, Q = [X] [Y] / [A] [B]
In other words the reaction quotient is defined in the same way as the equilibrium constant with molar concentrations to give Q_c or partial pressures to give Q_p as:
aA + bB ⇌ cC + dD
Q_c = [C] _c [D] _d/[A] _a [B] _b

Q_p = p^c_C p^d_D/ p^a_A p^b_B

At equilibrium, Q_c = K
Hence,
If Q_c > K_c, the reaction will proceed in the reverse direction i.e. in the direction of reactants.
If Q_c < K_c, the reaction will proceed in the forward direction i.e. in the direction of products.
If Q_c = K_c, the reaction will not proceed as reaction is already in equilibrium.
Hence we concluded that a reaction has the tendency to form products if Q < K and to form reactants if Q > K.