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Standard XII

Chemistry

Introduction

Equilibrium c...

Question

Equilibrium constant, $K_{c}$ for the reaction, $N_{2 (g)} + 3 H_{2 (g)} ⇌ 2 N H_{3 (g)}$ ; at $500 K$ is $0.061 l i t r e^{2} m o l e^{- 2}$ . At a particular time, the analysis shows that composition of the reaction mixture is $3.00 m o l l i t r e^{- 1}$ of $N^{2}$ , $2.00 m o l l i t r e^{- 1}$ of $H_{2}$ , and $0.500 m o l l i t r e^{- 1}$ of $N H_{3}$ . Is the reaction at equilibrium? If not, in which direction does the reaction tend to proceed to reach equilibrium?

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Solution

$Q_{c} = \frac{[N H_{3}]^{2}}{[N_{2}] [H_{2}]^{3}}$
$Q = \frac{(0.50)^{2}}{(3.0) (2.0)^{3}} = 0.0104$
The value of the reaction quotient is not equal to the value of the equilibrium constant.
Hence, the reaction is not at equilibrium.
The value of the reaction quotient is less than the value of the equilibrium constant.
Hence, the reaction will proceed to form more products, i.e., in the forward direction.

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At a particular time, the analysis shows that composition of the reaction mixture is 3.0 mol $L^{- 1} N_{2}, 2.0 m o l L^{- 1} H_{2}$ and 0.5 mol $L^{- 1} N H_{3}$ , Is the reaction at equilibrium? If not in which direction does the reaction tend to proceed to reach equilibrium?

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