Chemistry
Chemistry
Problem
A technician compares the effect of a strong base on (i) pure water and (ii) an aqueous buffer composed of 0.050 M formic acid and 0.050 M sodium formate (pK_a = 3.75). In each case, 0.0010 mol of NaOH is added to 1.00 L of solution at 25 °C. (a) Calculate the pH of pure water after the NaOH addition. (b) Calculate the pH of the buffer solution after the NaOH addition, assuming the added volume is negligible. (c) Using your numerical results, explain in words why the unbuffered solution exhibits a much larger pH change than the buffered solution. In your explanation, discuss the roles of both the ratio and the presence of a conjugate acid–base pair.Intuitive explanation
When you add strong base to a solution, the key difference is what it reacts with. In pure water, NaOH simply increases and dramatically raises pH. In a buffer, the added immediately neutralizes the weak acid present, converting it to its conjugate base—a process that consumes most of the base before pH changes much. The buffer's resistance comes from this reservoir of weak acid that "absorbs" the added base through acid–base equilibrium, keeping pH nearly constant, while pure water has no such defense.
Intuitive explanation
When you add strong base to a solution, the key difference is what it reacts with. In pure water, NaOH simply increases and dramatically raises pH. In a buffer, the added immediately neutralizes the weak acid present, converting it to its conjugate base—a process that consumes most of the base before pH changes much. The buffer's resistance comes from this reservoir of weak acid that "absorbs" the added base through acid–base equilibrium, keeping pH nearly constant, while pure water has no such defense.