NEET Crash Course Module - 25

chemca.in

Home Class XI Grand Test

Home › Class XI › NEET Rapid Revision › Buffer Solutions

NEET Crash Course • Module 25

Buffer Solutions & Common Ion Effect

Master the equilibrium that resists change. Decode the Henderson-Hasselbalch equation, calculate buffer capacity, and avoid the notorious in-situ mixing traps.

By chemca Academic Team • Updated for NEET 2027

Module Focus

Normally, adding a single drop of strong acid to water drastically drops its pH. A Buffer Solution is a special mixture designed to resist this change in pH upon the addition of small amounts of acid or base. The entire mechanism of a buffer relies on the Common Ion Effect. In competitive exams, you must be able to identify buffer pairs, use the Henderson-Hasselbalch equation, and spot "hidden" buffers created during neutralization.

1. The Common Ion Effect

The common ion effect is a direct consequence of Le Chatelier's Principle. It is defined as the suppression of the degree of dissociation ($\alpha$) of a weak electrolyte by the addition of a strong electrolyte containing a common ion.

Example Mechanism

Consider a weak acid, Acetic Acid ($CH_3COOH$), which partially dissociates:

$CH_3COOH(aq) \rightleftharpoons CH_3COO^-(aq) + H^+(aq)$

Now, we add a strong electrolyte, Sodium Acetate ($CH_3COONa$), which completely dissociates:
$CH_3COONa(aq) \rightarrow \mathbf{CH_3COO^-(aq)} + Na^+(aq)$

The Effect: The addition of $CH_3COONa$ floods the solution with the common ion, $CH_3COO^-$. According to Le Chatelier's principle, the equilibrium of the weak acid shifts backward to consume the excess product, drastically reducing the dissociation of $CH_3COOH$ and lowering the $[H^+]$ concentration.

2. Buffer Solutions

A buffer solution resists changes in its pH when small amounts of a strong acid or strong base are added, or when it is diluted.

A. Acidic Buffer

Mixture of a Weak Acid (WA) and its Salt with a Strong Base (WA-SB).

• $CH_3COOH + CH_3COONa$
• $HCN + KCN$
• $H_2CO_3 + NaHCO_3$ (Blood Buffer)

Maintains pH < 7

B. Basic Buffer

Mixture of a Weak Base (WB) and its Salt with a Strong Acid (WB-SA).

• $NH_4OH + NH_4Cl$
• $NH_4OH + (NH_4)_2SO_4$

Maintains pH > 7

Note: Simple Buffer - A salt of a weak acid and a weak base (e.g., $CH_3COONH_4$) acts as a buffer on its own because both its cation and anion undergo hydrolysis.

3. The Henderson-Hasselbalch Equation

This equation is used to calculate the exact pH or pOH of a buffer solution.

For Acidic Buffers

$pH = pK_a + \log_{10} \frac{[\text{Salt}]}{[\text{Acid}]}$

Or more accurately: $pH = pK_a + \log \frac{[\text{Conjugate Base}]}{[\text{Weak Acid}]}$

For Basic Buffers

$pOH = pK_b + \log_{10} \frac{[\text{Salt}]}{[\text{Base}]}$

Don't forget! $pH = 14 - pOH$

Important Trick for Molarity vs. Moles

Since the Salt and the Acid/Base are in the same solution volume (V), the volume term in the concentration ratio cancels out.
You can directly use the number of moles (or millimoles) in the Henderson equation:

$\frac{[\text{Salt}]}{[\text{Acid}]} = \frac{n_{\text{salt}} / V}{n_{\text{acid}} / V} = \mathbf{\frac{n_{\text{salt}}}{n_{\text{acid}}}}$

4. Buffer Capacity & Buffer Range

Buffer Capacity ($\beta$): A quantitative measure of a buffer's resistance to pH change. It is defined as the number of moles of strong acid or strong base required to be added to 1 Litre of buffer to change its pH by exactly 1 unit.

Condition for Maximum Buffer Capacity

A buffer solution is most effective (has maximum capacity) when the concentration of the salt is exactly equal to the concentration of the weak acid/base.

$[\text{Salt}] = [\text{Acid}] \quad \Rightarrow \quad pH = pK_a$

Buffer Range: The useful pH range of an acidic buffer is $pK_a \pm 1$.
(The ratio of $[\text{Salt}]/[\text{Acid}]$ must be between $\frac{1}{10}$ and $\frac{10}{1}$).

5. The Ultimate NEET Trap: In-Situ Buffers

Examiners rarely give you a direct mixture of a weak acid and its salt. Instead, they give you a neutralization reaction that results in a buffer.

How to Spot the Hidden Buffer

If you mix a Weak Acid (WA) and a Strong Base (SB), and the Weak Acid is in excess, the resulting mixture is an Acidic Buffer!

Example: Mix $100 \text{ mL}$ of $0.2 \text{ M } CH_3COOH$ with $50 \text{ mL}$ of $0.2 \text{ M } NaOH$.

• 1. Initial millimoles:
  WA ($CH_3COOH$) = $100 \times 0.2 = \mathbf{20 \text{ mmol}}$
  SB ($NaOH$) = $50 \times 0.2 = \mathbf{10 \text{ mmol}}$
• 2. Reaction: $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$
• 3. The 10 mmol of NaOH will completely consume 10 mmol of the weak acid to produce 10 mmol of Salt ($CH_3COONa$).
• 4. Leftovers in container:
  Unreacted WA = $20 - 10 = \mathbf{10 \text{ mmol}}$
  Formed Salt = $\mathbf{10 \text{ mmol}}$

Result: You have a Weak Acid + its Salt. IT IS A BUFFER!
Apply Henderson eq: $pH = pK_a + \log(\frac{10}{10}) = pK_a$.

Rule of Thumb: Partial Neutralization of a Weak Electrolyte = Buffer.

Target 180/180

NEET Grand Test: Buffer Solutions

15 High-Order Thinking Questions testing Henderson math, in-situ traps, and buffer capacity.

All Quizzes NEET Rapid Revision Series Video Lectures

🎯 NEET 2027 Target 180

Join the Ultimate Chemistry Crash Course

Master Ionic Equilibrium and Physical Chemistry. Get access to our full suite of Rapid Revision modules, formula sheets, and mock tests specifically designed for the NTA NEET pattern.

Explore All NEET Modules →

© 2026 chemca.in. Empowering NEET Aspirants.

🏠 Home 📝 Quiz 📘 Revision Notes ⬇ Downloads 📚 Class Notes ▶ Video Lectures

Get new posts by email

Powered by

📚 Also Read

Lecture Notes