Question

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

• 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 determined.

Answer 1

Answer: (C)

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

The given reaction is :-

$H_2\left(g\right)+{Br}_2\left(g\right)\rightleftharpoons 2HBr\left(g\right);K_C=2.5\times {10}^{3}at{400}^{0}c$

Now, $Q_C=\frac{{\left[HBr\right]}^2}{\left[H_2\right]\left[{Br}_2\right]}$

$=\frac{500\times 500}{0.05\times 0.03}=\frac{5\times 5\times {10}^6}{5\times 3}=\frac{5}{3}\times {10}^6$

Now, Here, $Q_C>K_C$

So, to estabish the equation $Q_C$ must decrease to become equal to $K_C$. So, reaction shifts to left.