The dissociation constant of HA and HB are 6 - 10-4 and 2-0 - 10-4 respectively- What is the pH of a solution which is a mixture of 0-5 M of HA and 0-5 M of HB - Take log 2-0-30 -

Total concentration of H^+ ion, [H^+]=√(K_a1× C_1+K_a2× C_2) Where K_a1 and nbsp;K_a2 are dissociation constant of nbsp;HA and nbsp;HB. So, [H^+]=√((6× 0. ...

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Standard XII

Chemistry

Common Ion Effect

The dissociat...

Question

The dissociation constant of $H A$ and $H B$ are $6 \times 10^{- 4}$ and $2.0 \times 10^{- 4}$ respectively. What is the pH of a solution which is a mixture of 0.5 M of $H A$ and 0.5 M of $H B$ ? (Take log2= 0.30).

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Solution

Total concentration of $H^{+}$ ion,
$[H^{+}] = \sqrt{K_{a 1} \times C_{1} + K_{a 2} \times C_{2}}$
Where $K_{a 1}$ and $K_{a 2}$ are dissociation constant of $H A$ and $H B$ .
So,
$[H^{+}] = \sqrt{(6 \times 0.5 + 2.0 \times 0.5) \times 10^{- 4}} = 2 \times 10^{- 2}$
$p H = - l o g (2 \times 10^{- 2}) = 1.70$

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The dissociation constant of HA and HB are 6×10−4 and 2.0×10−4 respectively. What is the pH of a solution which is a mixture of 0.5 M of HA and 0.5 M of HB? (Take log2= 0.30).

Total concentration of H+ ion, [H+]=√Ka1×C1+Ka2×C2 Where Ka1 and Ka2 are dissociation constant of HA and HB. So, [H+]=√(6×0.5+2.0×0.5)×10−4=2×10−2 pH=−log(2×10−2)=1.70

The dissociation constant of $H A$ and $H B$ are $6 \times 10^{- 4}$ and $2.0 \times 10^{- 4}$ respectively. What is the pH of a solution which is a mixture of 0.5 M of $H A$ and 0.5 M of $H B$ ? (Take log2= 0.30).