Question

In the following gaseous phase equilibrium at constant temperature.
$2S{O}_2\left(g\right)+O_2\left(g\right)\rightleftharpoons 2S{O}_3\left(g\right)$
The concentration of $\left[S{O}_2\right]=3.0\times {10}^{-3}M,\left[O_2\right]=3.5\times {10}^{-3}\text{M},\left[S{O}_3\right]=5.0\times {10}^{-2}\text{M}$. Calculate equilibrium constant.

• A. $K_c=7.936\times {10}^4$
• B. $K_c=17.936\times {10}^4$
• C. $K_c=12.936\times {10}^2$
• D. $K_c=2.936\times {10}^4$

Answer 1

The correct option is A $K_c=7.936\times {10}^4$

Given, the concentration of $\left[S{O}_2\right]=3.0\times {10}^{-3}M;\left[O_2\right]=3.5\times {10}^{-3}\text{M};\left[S{O}_3\right]=5.0\times {10}^{-2}\text{M}$
For the reaction,
$2S{O}_2\left(g\right)+O_2\left(g\right)\rightleftharpoons 2S{O}_3\left(g\right)$
$K_c=\frac{[S{O}_3{]}^2}{\left[S{O}_2{]}^2\right[O_2]}$
Putting all the values,
$=\frac{(5\times {10}^{-2}{)}^2}{(3\times {10}^{-3}{)}^2(3.5\times {10}^{-3})}=7.936\times {10}^4$