The ionization constant of HA is 1 × 10−6
.
What must be the ratio of [A−] to [HA] for
the preparation of a buffer solution with a pH
of 6.49?
To answer this question, we need to use the Henderson-Hasselbalch equation, which relates the pH of a buffer solution to the ratio of the concentrations of the conjugate base (A-) and weak acid (HA). The Henderson-Hasselbalch equation is given by:
pH = pKa + log([A-] / [HA])
Where:
- pH is the desired pH of the buffer solution (in this case, 6.49)
- pKa is the logarithmic value of the ionization constant (in this case, the ionization constant of HA is 1 × 10^-6)
- [A-] and [HA] are the concentrations of the conjugate base and weak acid, respectively.
In order to find the ratio of [A-] to [HA], we need to rearrange the Henderson-Hasselbalch equation:
log([A-] / [HA]) = pH - pKa
Now, substitute the given values into the equation:
log([A-] / [HA]) = 6.49 - (-6) = 12.49
Now we can get rid of the logarithm by taking the antilog:
10^(log([A-] / [HA])) = 10^12.49
[A-] / [HA] = 10^12.49
Therefore, in order to prepare a buffer solution with a pH of 6.49, the ratio of [A-] to [HA] should be equal to 10^12.49.