What is The Difference Between Reaction Quotient and Equilibrium Constant?

In the previous post we have learnt about equilibrium constant. Let’s revise it quickly.

• Equilibrium constant is a ratio of product and reactant’s molar concentration at equilibrium.
• It is independent of initial concentration of reactants.
• It is specific for a particular reaction.
• It is a temperature dependent constant.
• It is unit less constant because it is the ratio of concentrations.

Equilibrium is a state when reaction seems to be seized because both forward and backward reactions go on at the same rate and no change in the concentration of reactant and product is seen . So what information can we draw by equilibrium constant?

Equilibrium constant is the ratio of molar concentration of product to the molar concentration of reactant and its value gives general information about the extent of reaction. Larger value of equilibrium constant shows that reaction is nearer to the completion when equilibrium is established, as it means product concentration is larger than reactant. Similarly smaller value shows that equilibrium is established at the beginning of the reaction, because reactant concentration is much larger than product concentration. And if its value is equal to 1, it means reaction has finished half way through because product and reactant concentrations are equal at this point.

From above discussion one more concept about equilibrium has been clarified that equilibrium can be established at any point of reaction. It is the situation when rate of forward and backward reaction becomes equal. These are the general applications of equilibrium constant but how do we get information about a particular reaction? For this we need another parameter known as “Reaction Quotient” (Q_c). We can predict the direction of the particular reaction by comparing it with the equilibrium constant.

Reaction Quotient is similar to the equilibrium constant, the only difference is that in reaction quotient, concentrations of product and reactants are at a given time not necessarily at the time of equilibrium. Let’s take an example:

a A + b B ↔ c C + d D
Q_c  = [C]^c [D]^d / [A]^a[B]^b

Here concentration of products and reactants are taken in mole/L at the time t. At the time of equilibrium Q_c becomes equal to the K_c. Let’s see how Qc and Kc help us to predict the direction of the reaction. We will take an example to explain it:

H_2(g) + I_2(g) ↔ 2HI_(g)
K_c = 57.0 at 700K.

At time t molar concentration of H₂ = 0.10M, I₂ = 0.20M and HI = 0.40M so reaction quotient will be:

Q_c = (0.40)2 / (0.10) (0.20)
Q_c = 8.0

At the time t, Q_c is less than K_c and Q_c has to upgrade itself to reach to  K_c. In order to increase the value of Q_c reaction has to produce more product which means reaction will be shifted in forward direction.

If at a certain time t₂ you find Q_cis greater than K_c, then Q_c has to downgrade itself by increasing the concentration of reactants, that means reaction will be shifted in backward direction.

In the above discussion, we have seen that reversible reactions have a tendency to achieve equilibrium state, they adjust themselves either in backward or forward direction to do so. If we want to produce more HI in the above reaction, how can we take advantage of equilibrium. Can we fool it and get more HI? In the next post we will learn it.

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