What happens if an acid is added to this buffer?

[FredericChopin]

Homework Statement

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?

Homework Equations

N/A

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?

 

[Borek]

Assuming you managed to prepare solution that behaves the way you describe, you add the acid AND remove the conjugate base at the same time, so it is not as simple as just adding acid to the buffer.

Besides, when you add acid (base) to the buffer pH always goes down (up) - just slower than if the buffer wasn't present.

 

[Teemo]

Besides, when you add acid (base) to the buffer pH always goes down (up) - just slower than if the buffer wasn't present.

The reason for this is because while the acid is neutralized, remember that base was used to neutralize it. You lose hydroxide ions from solution and the volume of acid added will decrease the concentration of all entities involved.