Assertion (A): An aqueous solution of ammonium acetate can act as a buffer. Reason (R): Acetic acid is a weak

NCERT Exemplar · Q49

Q. Assertion (A): An aqueous solution of ammonium acetate can act as a buffer. Reason (R): Acetic acid is a weak acid and NH4OH is a weak base.

(i) Both A and R are true and R is correct explanation of A.

(ii) Both A and R are true but R is not correct explanation of A.

(iii) A is false but R is true.

(iv) Both A and R are false.

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Concept understanding — Buffer Solution pH

What is a Buffer Solution?

Imagine you're making lemonade. If you add a few drops of lemon juice to a glass of water, the pH drops sharply — it becomes very acidic. But if you add the same few drops to a glass of already acidic lemonade, the pH barely changes. Why? Because lemonade contains a buffer — a mixture that resists pH change when small amounts of acid or base are added.

A buffer solution is a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid). It "soaks up" added H⁺ or OH⁻ ions without letting the pH swing wildly.

Note

The key is that both components must be present in significant amounts. A weak acid alone won't buffer — you need its conjugate base partner too.

The Intuition: A Chemical Sponge

Think of a buffer as a two-way sponge:

If you add acid (H⁺): The conjugate base in the buffer grabs the extra H⁺, turning into the weak acid. The H⁺ is "absorbed" — pH barely drops.
If you add base (OH⁻): The weak acid donates an H⁺ to neutralise the OH⁻, turning into the conjugate base. The OH⁻ is "absorbed" — pH barely rises.

The buffer works best when the amounts of weak acid and conjugate base are roughly equal. That's when the sponge is most "spongy" — it can absorb shocks in either direction.

The Precise Statement: The Henderson–Hasselbalch Equation

For a buffer made from a weak acid $HA$ and its conjugate base $A^-$ , the pH is given by:

\text{pH} = \text{p}K_a + \log_{10} \left( \frac{[A^-]}{[HA]} \right)

Where:

$\text{p}K_a = -\log_{10} K_a$ (a measure of the weak acid's strength — lower pKa = stronger acid)
$[A^-]$ = concentration of the conjugate base
$[HA]$ = concentration of the weak acid

This equation tells you exactly how the pH depends on the ratio of base to acid, not their absolute amounts.

Tip

When $[A^-] = [HA]$ , the ratio is 1, $\log(1) = 0$ , so $\text{pH} = \text{p}K_a$ . This is the buffer's optimal pH — it resists change most strongly here.

Why This Works: A Quick Derivation

Start from the weak acid equilibrium:

HA \rightleftharpoons H^+ + A^-

The acid dissociation constant is:

K_a = \frac{[H^+][A^-]}{[HA]}

Take negative logs of both sides:

-\log K_a = -\log [H^+] - \log \frac{[A^-]}{[HA]}

Which gives:

\text{p}K_a = \text{pH} - \log \frac{[A^-]}{[HA]}

Rearrange:

\text{pH} = \text{p}K_a + \log \frac{[A^-]}{[HA]}

That's it. The derivation is just algebra on the definition of $K_a$ .

Watch out

The Henderson–Hasselbalch equation assumes that the concentrations $[HA]$ and $[A^-]$ are the initial concentrations you mixed. It works well when both are much larger than $[H^+]$ or $[OH^-]$ from dissociation — which is true for a properly made buffer.

Example: Making an Acetate Buffer

You mix 0.1 M acetic acid ( $\text{p}K_a = 4.76$ ) with 0.1 M sodium acetate. What's the pH?

$\text{pH} = 4.76 + \log \frac{0.1}{0.1} = 4.76 + \log 1 = 4.76$ …

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Ammonium acetate forms a buffer because it contains a weak acid ( $\text{CH}_3\text{COOH}$ ) and its conjugate base ( $\text{CH}_3\text{COO}^-$ ) from the salt, plus a weak base ( $\text{NH}_4\text{OH}$ ) and its conjugate acid ( $\text{NH}_4^+$ ). The assertion is true, but the reason given is incomplete — it only states that both parent species are weak, not that the salt provides the conjugate pair. So both statements are true, but R does not correctly explain A.

What makes a buffer? A buffer resists pH change when small amounts of acid or base are added. The classic recipe is a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) in roughly comparable amounts. For example, acetic acid ( $\text{CH}_3\text{COOH}$ ) plus sodium acetate ( $\text{CH}_3\text{COONa}$ ) works because the solution contains both the weak acid and its conjugate base ( $\text{CH}_3\text{COO}^-$ ).
Now look at ammonium acetate. This salt is formed from acetic acid (weak acid) and ammonium hydroxide (weak base). When dissolved in water, it dissociates completely:

\text{CH}_3\text{COONH}_4 \rightarrow \text{CH}_3\text{COO}^- + \text{NH}_4^+

The acetate ion is the conjugate base of acetic acid, and the ammonium ion is the conjugate acid of ammonia ( $\text{NH}_3$ ). So the solution contains both a weak acid ( $\text{NH}_4^+$ ) and its conjugate base ( $\text{NH}_3$ ), as well as a weak base ( $\text{CH}_3\text{COO}^-$ ) and its conjugate acid ( $\text{CH}_3\text{COOH}$ ). In effect, it is a salt of a weak acid and a weak base — and such a solution can indeed act as a buffer.

Why does it buffer? The key is that both the cation and the anion can react with added $H^+$ or $OH^-$ :
- Added acid ( $H^+$ ) is consumed by the acetate ion: $\text{CH}_3\text{COO}^- + H^+ \rightarrow \text{CH}_3\text{COOH}$
- Added base ( $OH^-$ ) is consumed by the ammonium ion: $\text{NH}_4^+ + OH^- \rightarrow \text{NH}_3 + H_2O$ So the solution has two independent buffering mechanisms. The pH of such a solution is given by $\text{pH} = \frac{1}{2}(\text{p}K_a + \text{p}K_b)$ — it is near neutral for ammonium acetate because $pK_a$ of acetic acid and $pK_b$ of ammonia are both about 4.75.
Evaluate the Assertion (A). "An aqueous solution of ammonium acetate can act as a buffer." This is true, as explained above. …

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Q.Assertion (A): An aqueous solution of ammonium acetate can act as a buffer. Reason (R): Acetic acid is a weak acid and NH4OH is a weak base.

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Q.Assertion (A): An aqueous solution of ammonium acetate can act as a buffer.
Reason (R): Acetic acid is a weak acid and NH4OH is a weak base.

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