Question

Consider the following gas phase reaction:
$H_2\left(g\right)+{Br}_2\left(g\right)\rightleftharpoons 2HBr\left(g\right)$
The concentrations of $H_2,{Br}_2$ and $HBr$ are 0.05 M, 0.03 M, and 500.0 M, respectively. The concentration of equilibrium constant for this reaction at ${400}^{\circ }C$ is $2.5\times {10}^3$. Is this system at equilibrium? Choose the right option.

• A. Yes, the system is at equilibrium.
• B. No, the reaction must shift to the right in order to reach equilibrium.
• C. No, the reaction must shift to the left in order to reach equilibrium.
• D. It cannot be determines.
• E. This system will never be at equilibrium.

Answer 1

Answer: (C)

No, the reaction must shift to the left in order to reach equilibrium.

$Q_c=\frac{[HBr{]}^2}{\left[H_2\right]\left[B{r}_2\right]}$

$Q_c=\frac{{500}^2}{0.05\times 0.03}=1.66\times {10}^8$

Since, $Q_c>K_c$, the reaction will proceed in reverse direction, converting product to reactants