Buffer Solutions

Full Notes

What are Buffer Solutions?

Buffer solutions are aqueous systems that resist changes in pH when small amounts of strong acid or strong base are added. They are essential in many biological and chemical processes where maintaining a constant pH is crucial (e.g., blood, enzyme activity, fermentation).

They work by containing a weak acid/base and its conjugate base/acid, which establish an equilibrium that can neutralize added H⁺ or OH⁻ ions.

Note: Buffers are covered in more detail in AP Chemistry Sections (see 8.8 and 8.9). This summary includes just what you need for NCERT Class 11, Chapter 6.12

Types of Buffer Solutions

1. Acidic Buffer
Contains a weak acid and its salt with a strong base.

Example Acetic acid (CH₃COOH) + Sodium acetate (CH₃COONa)

If HCl is added: Acetate ions (CH₃COO⁻) react with H⁺ to form acetic acid — limiting the drop in pH.

If NaOH is added: Acetic acid donates H⁺ to neutralize OH⁻, forming acetate — resisting any pH rise.

2. Basic Buffer
Contains a weak base and its salt with a strong acid.

Example Ammonia (NH₃) + Ammonium chloride (NH₄Cl)

If HCl is added: NH₃ neutralizes it by forming NH₄⁺.

If NaOH is added: NH₄⁺ donates H⁺ to form NH₃ and water — resisting pH increase.

Buffer Action Mechanism

Buffer action is governed by Le Chatelier’s Principle. When acid or base is added:

• The equilibrium shifts to counteract the change, minimizing the shift in [H⁺] or [OH⁻].
• This helps stabilize the pH.

For Example In an acetic acid buffer: CH₃COOH ⇌ CH₃COO⁻ + H⁺

• Add acid: H⁺ increases, equilibrium shifts left and more CH₃COOH forms = pH stabilizes.
• Add base: OH⁻ reacts with H⁺, equilibrium shifts right to replace H⁺ = pH stabilizes.

Designing Buffer Solutions

Preparation of Acidic Buffer
To prepare an acidic buffer, mix a weak acid with its salt formed from a strong base.

For a weak acid HA:

HA + H₂O ⇌ H₃O⁺ + A⁻

This gives us the equilibrium constant: K_a = [H₃O⁺][A⁻] / [HA]

Taking logarithms and rearranging: pK_a = pH − log([A⁻]/[HA])

Which leads to the Henderson–Hasselbalch equation:

• [A⁻] is the salt (conjugate base), and [HA] is the weak acid.
• If [A⁻] = [HA], then pH = pK_a
• To design a buffer with a specific pH, choose a weak acid with a pK_a close to the target pH.

Preparation of Basic Buffer
A similar approach is used for weak bases and their salts. For a base B and its conjugate acid BH⁺:

B + H₂O ⇌ BH⁺ + OH⁻

We use the base dissociation constant K_b to describe this equilibrium.

The equation is: pOH = pK_b + log([BH⁺]/[B])

Using the relation pH + pOH = 14, we get: pH = 14 − pK_b − log([BH⁺]/[B])

Rewritten as: pH = pK_a + log([Conjugate acid] / [Base])

For Example: Using ammonium chloride (NH₄Cl) and ammonium hydroxide (NH₄OH): pK_a (NH₄⁺) = 9.25
so pH = 9.25 + log([NH₄⁺]/[NH₃])

Important Note on Dilution

The pH of a buffer is unaffected by dilution.

Why? Because the ratio of [acid] and [conjugate base] (or base and conjugate acid) remains constant under the logarithm, even though both concentrations decrease equally.

Summary

• Buffers resist pH changes using a weak acid or base and its conjugate partner.
• Le Chatelier’s Principle explains how buffers neutralize small additions of acid or base.
• Use Henderson–Hasselbalch to calculate buffer pH from component ratios.
• Choose an acid or base with pK close to the target pH for effective buffering.