1. The problem statement, all variables and given/known data In grapes, there is an equilibrium between tartaric acid and hydrogen tartrate and hydrogen ions: (1) H2T (aq) <=> HT- (aq) + H+ NOTE: "T" stands for the tartrate ion C4H4O6. There is also a buffer system in grapes, involving a solubility equilibrium of potassium hydrogen tartrate and hydrogen tartrate and potassium ions: (2) K+ (aq) + HT- (aq) <=> KHT (s) What happens if you add H2T to this system? 2. Relevant equations N/A 3. The attempt at a solution If I added an acid to this system, the buffer will oppose the decrease in pH and so the pH will stay the same, right? Suppose I added H2T to this system: then I have increased the concentration of H2T, and due to Le Chatelier's Principle, equilibrium (1) will shift to produce more HT- and H+. This will increase the concentration of HT- however, so equilibrium (2) will shift so that K+ reacts with HT- to produce KHT, which will precipitate. But hold on, the concentration of H+ has increased, meaning the pH has gone down. The buffer has not opposed a decrease in pH! What's going on?