Question

3.2 moles of $HI\left(g\right)$ were heated in a sealed bulb at ${444}^{o}C$ till the equilibrium reached. Its degree of dissociation was found to be 20%. Calculate the number of moles of hydrogen iodide, hydrogen and iodine present at the equilibrium point and determine the value of the equilibrium constant for the reaction $2HI\left(g\right)\rightleftharpoons H_2\left(g\right)+I_2\left(g\right)$. Considering the volume of the container 1 L.

Answer 1

moles $2HI\rightleftharpoons H_2+I_2$

at $t=0$ $3.2$ - -

at $eq{b}^m$ $3.2-0.2\times 3.2$ $\frac{0.2\times 3.2}{2}$ $\frac{0.2\times 3.2}{2}$

$K_{eq}=\frac{{[}^n{H}_2\left]{[}^n{I}_2\right]}{{[}^{n}HI{]}^2}{\left[\frac{1}{V}\right]}^0$

$K_{eq}=\frac{0.32\times 0.32}{[2.56{]}^2}=1.56\times {10}^{-2}$

Number of moles at equilibrium,

$H_2⟶0.32$ moles

$I_2⟶0.32$ moles

$H_I⟶2.56$ moles