What happens if an acid is added to this buffer?

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?