To study the effects of a weakly acidic medium on the rate of corrosion of a metal alloy, you have been asked to prepare a 1 L buffer solution containing equilibrium concentrations of 0.739 M CH3COOH and 0.220 M CH3COONa. Given that the Ka of CH3COOH is 1.8 x 10–5, determine the pH of this solution at 25°C.

To determine the pH of the buffer solution, we can use the Henderson-Hasselbalch equation, which is given by:

$\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)$

where:

• $\text{p}K_a$ is the negative logarithm of the acid dissociation constant ($K_a$).
• $[\text{A}^-]$ is the concentration of the conjugate base (CH3COONa).
• $[\text{HA}]$ is the concentration of the weak acid (CH3COOH).

First, we need to calculate $\text{p}K_a$:

$\text{p}K_a = -\log(K_a)$ $\text{p}K_a = -\log(1.8 \times 10^{-5})$ $\text{p}K_a \approx 4.74$

Next, we use the Henderson-Hasselbalch equation to find the pH:

$\text{pH} = 4.74 + \log \left( \frac{0.220}{0.739} \right)$

Calculate the ratio inside the logarithm:

$\frac{0.220}{0.739} \approx 0.298$

Now, find the logarithm of this ratio:

$\log(0.298) \approx -0.526$

Finally, calculate the pH:

$\text{pH} = 4.74 + (-0.526)$ $\text{pH} \approx 4.21$

Therefore, the pH of the buffer solution is approximately 4.21.