Question

Which of the following is correct for acid buffer?

[salt = HA, acid = A]

• A. $pKa=pH-log\frac{\left[HA\right]}{\left[A\right]}$
• B. $pH=pKa+log\frac{\left[HA\right]}{\left[A\right]}$
• C. $pKa=pH-log\frac{\left[A\right]}{\left[HA\right]}$
• D. $pH=pKa+log\frac{\left[A\right]}{\left[HA\right]}$

Answer 1

Answer: (B)

$pH=pKa+log\frac{\left[HA\right]}{\left[A\right]}$

The $pH$ of a buffer is determined by two factors:

1) The equilibrium constant $Ka$ of the weak acid and

2) the ratio of the weak base $\left[A^{-}\right]$ to weak acid $\left[HA\right]$ in solution.

Different weak acids have different equilibrium constants $\left(Ka\right)$. $Ka$ tells us what proportion of $HA$ will be dissociated into $H^{+}$ and $A^{-}$ in solution. The more $H^{+}$ ions that are created, the more acidic and lower the $pH$ of the resulting solution.

This is formalized in the Henderson-Hasselbalch equation: $pH=p{K}_a+\text{log}\left(\frac{\left[\text{salt}\right]}{\left[\text{acid}\right]}\right)$.

Hence, the correct option is $\text{B}$