The following two reactions describe the chemistry of $\text{S(IV)}$ in an aquatic system.
Under ideal conditions, the equilibrium constants for reactions $\text{(i)}$ and $\text{(ii)}$ are $1.3 \times 10^{-2} \mathrm{M}\left(\mathrm{K}_{\mathrm{S} 1}\right)$ and $6.6 \times 10^{-8} \mathrm{M}\left(\mathrm{K}_{\mathrm{S} 2}\right)$, respectively.
If the pH of the system is $4.0$ and the equilibrium concentration of $\mathrm{SO}_{2(\mathrm{aq})}$ is $1.0$ $\mathrm{M}$, the equilibrium concentration of $\mathrm{SO}_{3}^{2-}$ is $\_\_\_\_$ $\mathrm{mM}$ (rounded off to one decimal place).
$$\begin{array}{l}
\mathrm{SO}_{2(a q)} \rightleftharpoons \mathrm{H}^{+}+\mathrm{HSO}_{3}^{-} \quad \text{(i)}\\
\mathrm{HSO}_{3}^{-} \rightleftharpoons \mathrm{H}^{+}+\mathrm{SO}_{3}^{2-} \quad \text{(ii)}
\end{array}$$
