8.8.A.1  Buffer Action: Conjugate Pair Stabilizes pH by Neutralizing Added Acid or Base:

1. Acids, Bases, and Conjugate Pairs:

a. Arrhenius Definition (simplest):

* Acid: It raises the concentration of hydrogen ions (H⁺) in aqueous solution.
* Base: It raises the concentration of hydroxide ions (OH⁻) in aqueous solution.

Example:

* HCl → H⁺ + Cl⁻ (acid)
* NaOH → Na⁺ + OH⁻ (base)

b. Brønsted-Lowry Definition (more general and commonly used):

* Acid: Donor of a proton (H⁺).
* Base: Acceptor of a proton (H⁺).

This definition is broader than aqueous solutions and is central to the concept of conjugate pairs.

ii. Conjugate Acid-Base Pairs:

If an acid donates a proton (H⁺), it becomes its conjugate base.
If a base accepts a proton, it becomes its conjugate acid.

iii. General Reaction:

Acid ⇌ Conjugate Base + H⁺
Base + H⁺ ⇌ Conjugate Acid

iv. Identifying Conjugate Pairs:

For any acid-base reaction, determine the species that differ by a single proton:

v. Important Notes:

* The stronger the acid, the weaker its conjugate base, and vice versa.
* Water can act as both an acid and a base (amphoteric).
* Every acid-base reaction involves two conjugate pairs.

2. Buffer Composition:

A buffer is a solution which resists pH change when small quantities of acid or base are added to it. Buffers are used in most chemical and biological systems to keep the pH stable.

i. What Constitutes a Buffer?

A buffer consists of:

* A weak acid and its conjugate base, or
* A weak base and its conjugate acid

These elements collaborate to neutralize added bases or acids.

ii. Common Buffer Types

a. Weak Acid + Conjugate Base:

* Example: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa)
* Reaction:

* CH₃COOH ⇌ CH₃COO⁻ + H⁺
* When base (OH⁻) is added, CH₃COOH donates H⁺ to neutralize it.
* When acid (H⁺) is added, CH₃COO⁻ absorbs H⁺ to make more CH₃COOH.

b. Weak Base + Conjugate Acid:

* Example: Ammonia (NH₃) and ammonium chloride (NH₄Cl)
* Reaction:

* NH₃ + H⁺ ⇌ NH₄⁺
* When acid is added, NH₃ accepts H⁺.
* When base is added, NH₄⁺ donates H⁺.

iii. Why Buffers Work:

* The weak acid neutralizes added base (by donating H⁺).
* The conjugate base neutralizes added acid (by accepting H⁺).

Since the components are in equilibrium, the system can adjust and have a relatively constant pH.

iv. Ideal Buffer Conditions:

* The acid and conjugate base (or base and conjugate acid) must be present in comparable concentrations.
* Buffers work best when the pH is near the pKa of the weak acid (±1 pH unit).

v. Buffer Example (Acetic Acid Buffer):

* Components: CH₃COOH (weak acid) and CH₃COONa (salt of conjugate base)
* Buffer Equation:  CH₃COOH ⇌ CH₃COO⁻ + H⁺

3. Buffer Mechanism:

A buffer accomplishes this by neutralizing a small quantity of added acid or base, thereby reducing the degree of change in pH. This is achieved through the fact that a buffer has both:

* A weak acid (which can release H⁺), and
* Its conjugate base (which can absorb H⁺),
or the reverse.

i. When an Acid (H⁺) Is Added:

The conjugate base of the buffer reacts with added H⁺ to neutralize it.

Example (Acetic Acid Buffer):

Buffer components:

* Weak acid: CH₃COOH
* Conjugate base: CH₃COO⁻

Reaction with added acid:H⁺ + CH₃COO⁻ → CH₃COOH

What happens?

* The added H⁺ gets consumed.
* The concentration of the weak acid (CH₃COOH) increases slightly.
* pH doesn’t change much.

ii. When a Base (OH⁻) Is Added:

The weak acid in the buffer gives up a proton (H⁺) to counteract the added OH⁻.

Reaction with added base:

OH⁻ + CH₃COOH → CH₃COO⁻ + H₂O

What happens?

* OH⁻ is neutralized as water is formed.
* Conjugate base (CH₃COO⁻) rises slightly.
* pH remains virtually the same.

iii. Why It Works:

* Buffers work due to the equilibrium between the weak acid and its conjugate base:HA ⇌ H⁺ + A⁻
* Adding H⁺ pushes equilibrium left → fewer free H⁺ is left → pH does not decrease much.
* Adding OH⁻ takes away H⁺ (to form water) → equilibrium moves right to replace H⁺ → pH does not increase much.

iv. Buffer Capacity:

* A buffer can neutralize only certain amounts of acid or base.
* After one component (acid or base) is depleted, the buffer becomes ineffective.
* Ideal buffer capacity is achieved when
[weak acid] ≈ [conjugate base] (i.e., when pH ≈ pKa).

v. Summary Table

4. Buffer Capacity and Limitations:

i. What is Buffer Capacity?

Buffer capacity is the quantity of strong acid or base that can be neutralized without a large change in pH by a buffer solution.

It is a function of:

* Concentration of the buffer components (weak acid and its conjugate base, or vice versa).
* The ratio of acid to base (ideally close to 1:1).

ii. Key Factors Affecting Buffer Capacity:

iii. Buffer’s Effective pH Range:

* Buffers work best when the pH is ±1 unit from the pKa of the weak acid.
* Outside this range, one of the buffer components is in low concentration and can’t resist pH changes well.

iv. What Happens When Buffer Capacity is Exceeded?

If you add more strong acid or base more than the buffer can neutralize, then:

* The weak acid or conjugate base is used up.
* The pH will change rapidly.
* The buffer effectively fails.

v. Example:

* Buffer: 0.1 M CH₃COOH and 0.1 M CH₃COO⁻
* Is able to manage around **0.01–0.02 moles of strong acid or base per liter** before the pH becomes significantly different.

iv. Summary: