# Salts and buffer solutions

> IB Chemistry HL · S3: Classification of matter
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u3-salts-and-buffer-solutions/

This sub-topic covers pH behavior of dissolved salts, the composition and mechanism of buffer solutions, and pH calculations for buffer systems. It is a high-frequency core topic for HL IB Chemistry exams.

**Prerequisites:** [Bronsted-Lowry acids and bases](https://www.owlsprep.com/study/ib-chemistry-hl-bronsted-lowry-acids-bases/); [pH and Ka calculations](https://www.owlsprep.com/study/ib-chemistry-hl-ph-ka-calculations/)

## Learning objectives

- Predict the pH of aqueous salt solutions from hydrolysis of ions
- Explain the mechanism of buffer action
- Calculate buffer pH using the Henderson-Hasselbalch equation
- Describe factors affecting buffer capacity

## Salt Hydrolysis and Salt Solution pH

**Salt hydrolysis** — A reaction between a dissolved salt ion and water that produces $H^+$ or $OH^-$ ions, changing the solution pH from neutral.

*Example:* Ammonium ions reacting with water to produce hydronium ions

- Salt of **strong acid + strong base**: Neutral (no hydrolysis, pH ≈ 7)
- Salt of **weak acid + strong base**: Basic (conjugate base of weak acid hydrolyzes to produce $OH^-$, pH > 7)
- Salt of **strong acid + weak base**: Acidic (conjugate acid of weak base hydrolyzes to produce $H^+$, pH < 7)
- Salt of **weak acid + weak base**: pH depends on relative $K_a$ of the acid and $K_b$ of the base

**Worked example:** Predict whether an aqueous solution of sodium ethanoate ($CH_3COONa$) is acidic, basic, or neutral, and justify your answer.

1. Dissociate the salt into ions and identify parent acid/base:
2. $$CH_3COONa(s) \rightarrow CH_3COO^-(aq) + Na^+(aq)$$
3. Parent acid: ethanoic acid ($CH_3COOH$, weak acid), parent base: sodium hydroxide ($NaOH$, strong base)
4. Identify the hydrolyzable ion: $Na^+$ does not react with water, $CH_3COO^-$ is the conjugate base of a weak acid so it hydrolyzes:
5. $$CH_3COO^-(aq) + H_2O(l) \rightleftharpoons CH_3COOH(aq) + OH^-(aq)$$
6. Production of $OH^-$ increases pH above 7, so the solution is basic.

> **Exam tip:** Always check both the cation and anion for hydrolysis before predicting pH — never assume all salts are neutral.

## Buffer Composition and Mechanism of Action

**Buffer solution** — A solution that resists changes in pH when small amounts of strong acid, strong base, or water are added.

*Notation:* Composed of a weak conjugate acid-base pair

Buffers are made in two common ways: mixing a weak acid with its conjugate base (as a soluble salt), or mixing a weak base with its conjugate acid (as a soluble salt). Strong acids and bases cannot form buffers, as they dissociate completely with no equilibrium to absorb added ions.

**Worked example:** Explain how an ammonia / ammonium chloride buffer resists an increase in pH when a small amount of $NaOH$ is added.

1. Main species in buffer: $NH_3$ (weak base, high concentration), $NH_4^+$ (conjugate acid from ammonium chloride, high concentration)
2. Added $NaOH$ dissociates completely to produce extra $OH^-$ ions.
3. Excess $OH^-$ reacts with the conjugate acid $NH_4^+$:
4. $$NH_4^+(aq) + OH^-(aq) \rightleftharpoons NH_3(aq) + H_2O(l)$$
5. Almost all excess $OH^-$ is consumed, so pH increases only slightly.

> **warning**
>
> Buffers cannot resist large pH changes: adding enough strong acid/base to consume all buffer components will cause a dramatic pH shift.

## Buffer pH Calculations: Henderson-Hasselbalch Equation

**Derivation:** Derive the Henderson-Hasselbalch equation for an acidic buffer

*Starting from:* Acid dissociation equilibrium of a weak acid $HA$

1. Write the acid dissociation equilibrium and $K_a$ expression:
2. $$HA(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + A^-(aq)$$
3. $$K_a = \frac{[H_3O^+][A^-]}{[HA]}$$
4. Rearrange to isolate $[H_3O^+]$:
5. $$[H_3O^+] = K_a \times \frac{[HA]}{[A^-]}$$
6. Take $-\log_{10}$ of both sides and simplify using $pH = -\log[H_3O^+]$ and $pK_a = -\log K_a$:
7. $$-\log[H_3O^+] = -\log K_a - \log\left(\frac{[HA]}{[A^-]}\right)$$

The common ion effect suppresses dissociation of the weak acid, so we can approximate $[HA] \approx$ initial weak acid concentration and $[A^-] \approx$ initial salt concentration for calculations.

**Worked example:** Calculate the pH of a buffer containing 0.15 mol dm⁻³ propanoic acid ($pK_a = 4.87$) and 0.10 mol dm⁻³ sodium propanoate.

1. Identify components: weak acid = propanoic acid, $[acid] = 0.15$, conjugate base = propanoate, $[base] = 0.10$
2. Substitute into the Henderson-Hasselbalch equation:
3. $$pH = 4.87 + \log_{10}\left(\frac{0.10}{0.15}\right)$$
4. Calculate the log term: $\frac{0.10}{0.15} \approx 0.667$, $\log_{10}(0.667) \approx -0.18$
5. Compute final pH:
6. $$pH = 4.87 - 0.18 = 4.69$$

*Calculator:* allowed

## Buffer Capacity

**Buffer capacity** — A measure of the amount of strong acid or base that can be added to a buffer before the pH changes significantly.

- **Total concentration of buffer components**: Higher total concentration = higher buffer capacity, as more ions are available to react with added acid/base.
- **Ratio of buffer components**: Buffer capacity is maximum when $\frac{[conjugate\ base]}{[weak\ acid]} = 1$, so when $pH = pK_a$ of the weak acid.
- Buffers are only effective within $pH = pK_a \pm 1$ of the weak acid's $pK_a$.

> **tip**
>
> When asked to choose a buffer for a target pH, always pick a weak acid with a $pK_a$ within 1 unit of your target pH.

## Common pitfalls

- **Wrong:** Assuming all salts form neutral aqueous solutions
  - Why it fails: Most salts made from weak acids or bases have ions that hydrolyze to change pH
  - Correct: Always check both the cation and anion for hydrolysis, using the parent acid/base rules to predict pH
- **Wrong:** Swapping the [acid] and [base] terms in the Henderson-Hasselbalch equation
  - Why it fails: The derivation gives the ratio of conjugate base to weak acid in the log term; swapping gives an incorrect pH
  - Correct: Memorize that $pH = pK_a + \log(\frac{[base]}{[acid]})$, or re-derive quickly from the $K_a$ expression if unsure
- **Wrong:** Using a strong acid and strong base to prepare a buffer solution
  - Why it fails: Strong acids/bases dissociate completely, so no equilibrium exists to absorb added $H^+$ or $OH^-$
  - Correct: Buffers always require a weak conjugate acid-base pair
- **Wrong:** Thinking buffer capacity only depends on the 1:1 ratio of components
  - Why it fails: Total concentration of components also impacts capacity: dilute buffers have lower capacity even with a 1:1 ratio
  - Correct: Remember buffer capacity depends on both the ratio of components (maximum at 1:1) and total concentration (higher = higher capacity)
- **Wrong:** Expecting large pH changes when diluting a buffer
  - Why it fails: Dilution changes both [acid] and [base] equally, so their ratio stays the same
  - Correct: Buffer pH remains almost unchanged upon dilution, only extreme dilution causes significant change

## Cheatsheet

| Property | Rule |
| --- | --- |
| Strong acid + strong base salt | Neutral, pH ≈ 7 |
| Weak acid + strong base salt | Basic, pH > 7 |
| Strong acid + weak base salt | Acidic, pH < 7 |
| Henderson-Hasselbalch (acid buffer) | $pH = pK_a + \log(\frac{[A^-]}{[HA]})$ |
| Maximum buffer capacity | $pH = pK_a$, 1:1 [base]:[acid] |
| Effective buffer pH range | $pK_a \pm 1$ |

## What's next

Salts and buffer solutions are a core foundation for understanding acid-base titrations, a common extended response topic in IB Chemistry HL exams. Buffers also have important biological applications, such as the bicarbonate buffer that maintains human blood pH, which is frequently tested in exam questions. Next, you can explore acid-base titration curves and indicators, where you will apply your understanding of salt pH and buffer action to calculate pH at different points of a titration, including at the equivalence point.

- [AHL: Extended periodic trends](https://www.owlsprep.com/study/ib-chemistry-hl-u3-ahl-extended-periodic-trends/)
- [AHL: Advanced Lewis acid-base concepts](https://www.owlsprep.com/study/ib-chemistry-hl-u3-ahl-advanced-lewis-acid-base/)

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