Net charge of citric acid at pHs of: 1) 1 2) 5 3) 7 4) 9
The Correct Answer and Explanation is:
Net Charge of Citric Acid at Different pH Values
Citric acid is a triprotic weak acid, meaning it has three carboxylic acid groups (–COOH) and one hydroxyl group (–OH). The three carboxyl groups each have a dissociation constant (pKa) at which they lose a proton:
- pKa₁ ≈ 3.1
- pKa₂ ≈ 4.8
- pKa₃ ≈ 6.4
These pKa values indicate the pH at which each carboxyl group is 50% dissociated. Using this, we can determine the net charge of citric acid at different pH values by estimating how many protons are lost (deprotonation):
1) pH 1
- pH << pKa₁, pKa₂, and pKa₃
- All three –COOH groups are protonated (neutral)
- The molecule is in its fully protonated form
- Net charge = 0
2) pH 5
- pH > pKa₁ and pKa₂, but < pKa₃
- First and second carboxylic acids are deprotonated (–COO⁻), third is still protonated
- So, two negative charges, one neutral
- Net charge = –2
3) pH 7
- pH > all three pKa values
- All three carboxylic acids are deprotonated
- Three negative charges
- Net charge = –3
4) pH 9
- pH >> all pKa values
- Same as at pH 7 — fully deprotonated
- No further ionizable groups
- Net charge = –3
Summary Table:
| pH | Protonation State | Net Charge |
|---|---|---|
| 1 | All COOH groups protonated | 0 |
| 5 | Two COOH deprotonated | –2 |
| 7 | All COOH deprotonated | –3 |
| 9 | All COOH deprotonated | –3 |
Explanation:
As the pH increases, citric acid loses protons from its acidic groups. This deprotonation is governed by the pKa values of its carboxylic acids. When the pH is well below a group’s pKa, that group remains protonated (neutral). When the pH rises above a group’s pKa, the group tends to lose a proton and becomes negatively charged. This stepwise loss of protons determines the molecule’s net charge at any given
