Buffer Solution

Interactive visualization of buffer solutions - Explore pH calculation, titration curves, conjugate pairs, acid/base resistance, buffer capacity, and molecular dynamics animation

pH Calculator - Henderson-Hasselbalch

Current pH: 0.00
pKa: 0.00
[A⁻]/[HA] Ratio: 0.00

Titration Curve: pH vs Volume (Buffer Region Marked)

Buffer Range: pKa ± 1
Effective Buffer: Yes

Molecular View: Conjugate Pair Animation

HA (acid) A⁻ (conjugate base)

Buffer Capacity (β)

Max Capacity at pH = pKa: 0.00
Current β: 0.00

Buffer Solution Controls

Buffer Parameters

Value is log₁₀(Ka), range: 10⁻¹⁰ to 10⁻²

Acid/Base Addition (Test Buffer Capacity)

Original pH: 0.00
New pH: 0.00
ΔpH: 0.00

Common Buffer Systems

Titration Simulation

Buffer Solution Equations

Henderson-Hasselbalch Equation: pH = pKa + log([A⁻]/[HA])
Acid Ionization: HA ⇌ H⁺ + A⁻
Acid Dissociation Constant: Ka = [H⁺][A⁻]/[HA]
Buffer Capacity (β): β = 2.303 × [H⁺] × [A⁻] / ([H⁺] + [A⁻])
Optimal Buffer Condition: pH = pKa when [A⁻] = [HA]
Effective Buffer Range: pKa ± 1 pH unit

What is a Buffer Solution?

A buffer solution is a solution that resists changes in pH when small amounts of acid or base are added. It consists of a weak acid (HA) and its conjugate base (A⁻), or a weak base and its conjugate acid. The buffer capacity is maximum when the concentrations of the acid and conjugate base are equal ([HA] = [A⁻]), which occurs when pH = pKa.

Henderson-Hasselbalch Equation

The Equation: pH = pKa + log([A⁻]/[HA]) relates buffer pH to pKa and concentration ratio.
When [A⁻] = [HA]: Ratio is 1, log(1) = 0, so pH = pKa (maximum buffer capacity).
Buffer Range: Effective buffers work within ±1 pH unit of pKa (ratio 0.1 to 10).
Practical Use: Calculate buffer pH or determine ratio for desired pH.

Buffer Capacity (β)

Definition: Quantifies acid/base a buffer can neutralize before significant pH change.
Formula: β = 2.303 × [H⁺] × [A⁻] / ([H⁺] + [A⁻]).
Maximum Capacity: At pH = pKa when [HA] = [A⁻]. Higher concentration increases capacity.
Practical Implication: Stronger buffers resist larger pH changes.

Buffer Mechanism

Adding Acid (H⁺): A⁻ neutralizes added H⁺: A⁻ + H⁺ → HA.
Adding Base (OH⁻): HA neutralizes added OH⁻: HA + OH⁻ → A⁻ + H₂O.
Le Chatelier's Principle: Equilibrium shifts to counteract disturbances.
Limitation: Buffer capacity is finite - once HA or A⁻ is depleted, pH changes rapidly.

Biological Buffer Systems

Blood Buffering: Multiple systems maintain pH ≈ 7.4. Bicarbonate buffer (H₂CO₃/HCO₃⁻, pKa ≈ 6.1) is most important.
Phosphate Buffer: H₂PO₄⁻/HPO₄²⁻ (pKa ≈ 7.2) buffers intracellular fluid.
Protein Buffers: Amino acid side chains act as buffers in proteins.
Clinical Significance: Acid-base imbalances occur when buffer systems are overwhelmed.

Real-World Applications

Laboratory Use: Calibrate pH meters, maintain enzyme activity, control reaction conditions.
Pharmaceuticals: Buffered formulations for drug stability.
Food Industry: Control pH in processed foods and beverages.
Water Treatment: Maintain optimal pH for treatment processes.
Agriculture: Soil buffers affect nutrient availability.

Buffer Preparation

Choosing pKa: Select weak acid with pKa within ±1 of desired pH.
Calculating Ratio: Use Henderson-Hasselbalch to determine [A⁻]/[HA] ratio.
Concentration: Higher concentrations (0.01-1.0 M) provide greater capacity.
Common Buffers: Acetate (pH 3.6-5.6), phosphate (pH 5.8-8.0), TRIS (pH 7.0-9.0).