You continue to use the buffer that you prepared in Example 6G.1 but are worried about how adding sodium hydroxide to the buffer solution will affect the pH, which could upset your experiment. Suppose 1.2 g of sodium hydroxide (0.030 mol NaOH) is dissolved in 500. mL of the buffer solution described in Example 6G.1. Calculate the pH of the resulting solution and the change in pH.

ANTICIPATE Because the buffer solution contains a weak acid that will react with the strong base, you should expect only a small increase in the pH from the 4.44 of the initial solution.

PLAN Solve this problem in two stages. First, calculate the concentrations of acid and conjugate base by noting that the OH^– ions added to the buffer solution react with some of the acid of the buffer system, decreasing the amount of acid and correspondingly increasing the amount of conjugate base. Then rearrange the expression for K_a to obtain the pH of the solution, just as in Example 6G.1.

What should you assume? Assume that the equilibrium concentrations of the acid and its conjugate base remain nearly the same as their initial (formal) concentrations following neutralization and that the neutralization reaction goes to completion. Assume that the volume of the solution remains unchanged.

Example 6G.1

You are working in a microbiology laboratory culturing bacteria that require an acidic environment and you need to prepare a buffer to keep the culture at an appropriate pH. You prepare a buffer solution that is 0.040 m NaCH₃CO₂(aq) and 0.080 m CH₃COOH(aq) at 25°C. What is the pH of the solution?

ANTICIPATE If the weak acid were present alone, you would expect pH < 7. Because some conjugate base has been added, you should expect the pH to be slightly higher than the acid-only value (which is estimated to be about 2.92 in Example 6D.1) but still less than 7.

PLAN First, identify the weak acid and its conjugate base. Then, write the proton transfer equilibrium between them and rearrange the expression for Ka to give [H₃O⁺]. Finally, calculate the pH.

What should you assume? Assume that the protonation of acetate ions and the deprotonation of acetic acid molecules have both proceeded to such a small extent that the concentrations of both species are nearly the same as their initial (formal) values.