$50 \mathrm{~mL}$ of $\mathrm{H}_2 \mathrm{O}$ is added to $50 \mathrm{~mL}$ of $1 imes 10^{-3} \mathrm{M}$ barium hydroxide solution. What is the $\mathrm{pH}$ of the resulting solution?

Question

$50 \mathrm{~mL}$ of $\mathrm{H}_2 \mathrm{O}$ is added to $50 \mathrm{~mL}$ of $1 	imes 10^{-3} \mathrm{M}$ barium hydroxide solution. What is the $\mathrm{pH}$ of the resulting solution?

Aicharya · Accepted Answer

In water, barium hydroxide is hydrolysed as follows. $$ \begin{aligned} \mathrm{Ba}(\mathrm{OH})_2 ightleftharpoons & \mathrm{Ba}^{2+}+2 \mathrm{OH}^{-} \ 	ext {conc. of } \mathrm{Ba}^{2+} & =1 	imes 10^{-3} \mathrm{M} \ 	ext { conc. of }\left[\mathrm{OH}^{-}ight] & =2 	imes 1 	imes 10^{-3} \mathrm{M} \ & =2 	imes 10^{-3} \mathrm{M} \ 	ext { pOH } & =-\log \left[\mathrm{OH}^{-}ight] \ & =-\log \left(2 	imes 10^{-3}ight) \ & =2.69 \end{aligned} $$ $\begin{aligned} \mathrm{pH}+\mathrm{pOH} & =14 \ \mathrm{pH} & =14-\mathrm{pOH} \ & =14-2.69 \ & =11.3 \ & \approx 11.0\end{aligned}$

$50 \mathrm{~mL}$ of $\mathrm{H}_2 \mathrm{O}$ is added to $50 \mathrm{~mL}$ of $1 \times 10^{-3}…

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