Buffer Solutions: Complete Guide to Buffers
Buffer solutions are essential in chemistry, biology, and medicine. They resist changes in pH when small amounts of acid or base are added. Let’s explore how buffers work and how to calculate their pH.
What is a Buffer Solution?
Definition: A buffer solution is one that resists changes in pH on addition of a small amount of a strong acid or strong base.
A buffer solution consists of:
- A weak acid and its salt with a strong base (acidic buffer)
- OR a weak base and its salt with a strong acid (alkaline buffer)
Types of Buffers
1. Acidic Buffer (Weak Acid + Salt of Weak Acid)
Example: Acetic acid (CH₃COOH) + Sodium acetate (CH₃COONa)
Composition:
CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq) (Weak acid) (Conjugate base from salt)
How it maintains pH:
When a strong acid (HCl) is added:
H⁺(aq) + CH₃COO⁻(aq) → CH₃COOH(aq) The added H⁺ is neutralized by the acetate ions, preventing pH drop
When a strong base (NaOH) is added:
OH⁻(aq) + CH₃COOH(aq) → CH₃COO⁻(aq) + H₂O(l) The acetic acid neutralizes the added OH⁻, preventing pH rise
Ideal Buffering Range: pH 6-8 for acetic acid buffer
2. Alkaline Buffer (Weak Base + Salt of Weak Base)
Example: Ammonia (NH₃) + Ammonium chloride (NH₄Cl)
Composition:
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) (Weak base) (Conjugate acid from salt)
Other Examples:
- H₂CO₃/NaHCO₃ (carbonic acid buffer)
- H₃PO₄/Na₂HPO₄ (phosphate buffer)
- HCO₃⁻/NaHCO₃ (human blood buffer)
pH Calculation for Buffers
Using the Henderson-Hasselbalch Equation
For an acidic buffer:
pH = pKa + log([salt]/[acid]) OR pH = pKa + log([A⁻]/[HA])
For an alkaline buffer:
pOH = pKb + log([salt]/[base]) OR pH = 14 - pOH
Worked Example 1: Acetic Acid Buffer
A buffer solution containing 0.1M acetic acid (CH₃COOH) and 0.2M sodium acetate (CH₃COONa) is prepared. Given Ka for acetic acid = 1.8 × 10⁻⁵, calculate the pH of the buffer.
Solution:
pKa = -log(1.8 × 10⁻⁵) = 4.74 pH = pKa + log([CH₃COO⁻]/[CH₃COOH]) pH = 4.74 + log(0.2/0.1) pH = 4.74 + log(2) pH = 4.74 + 0.30 pH = 5.04
Worked Example 2: Ammonia Buffer
A buffer solution contains 0.06M ammonia (NH₃) and 0.06M ammonium chloride (NH₄Cl). Given Kb for ammonia = 1.8 × 10⁻⁵, calculate the pH of the buffer.
Solution:
pKb = -log(1.8 × 10⁻⁵) = 4.74 pOH = pKb + log([NH₄⁺]/[NH₃]) pOH = 4.74 + log(0.06/0.06) pOH = 4.74 + log(1) pOH = 4.74 + 0 pOH = 4.74 pH = 14 - pOH = 14 - 4.74 = 9.26
Buffer Capacity
Definition: Buffer capacity is the amount of acid or base a buffer can neutralize while maintaining relatively constant pH.
Factors Affecting Buffer Capacity:
- Concentration: Higher concentration = greater buffer capacity
- Ratio of Base to Acid: Buffer works best when [A⁻] ≈ [HA], i.e., ratio is close to 1:1
- Nature of Acid/Base: Different acids have different pKa values
Buffer Range: A buffer is most effective when pH is within ±1 of the pKa value.
Effective pH range = pKa ± 1
Example: For acetic acid (pKa = 4.74), the buffer range is 3.74 – 5.74
Effect of Adding Acid or Base to a Buffer
Adding a Strong Acid to an Acetic Acid Buffer
Example: Add 10 cm³ of 1M HCl to 1 dm³ of buffer containing 0.1M CH₃COOH and 0.2M CH₃COONa
Moles of H⁺ added = 1 × 10/1000 = 0.01 mol CH₃COO⁻ + H⁺ → CH₃COOH New [CH₃COOH] = 0.1 + 0.01 = 0.11 M New [CH₃COO⁻] = 0.2 - 0.01 = 0.19 M New pH = 4.74 + log(0.19/0.11) = 4.74 + 0.24 = 4.98 Original pH was 5.04 Change in pH = 5.04 - 4.98 = 0.06 units (minimal change)
Uses of Buffer Solutions
- Laboratory Analysis: Calibration of pH meters
- Biochemistry: Maintaining pH for enzyme reactions
- Medicine: Human blood pH maintenance (H₂CO₃/HCO₃⁻ buffer system)
- Food Industry: pH control in food processing
- Fermentation: Maintaining pH in fermentation processes
- Pharmaceuticals: Drug formulation and stability
Blood as a Buffer System
Human blood contains a carbonic acid-bicarbonate buffer system:
H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)
Normal blood pH is maintained at 7.4 ± 0.2. This is crucial because:
- pH below 7.0 (acidosis) or above 7.8 (alkalosis) can be fatal
- The buffer system responds immediately to metabolic changes
- The respiratory system helps by regulating CO₂ levels
- The kidneys help by regulating HCO₃⁻ levels
Summary Table: Common Buffers
| Buffer System | Type | pKa/pKb | Effective pH Range | Application |
|---|---|---|---|---|
| Acetic acid/Acetate | Acidic | 4.74 | 3.74 – 5.74 | Laboratory |
| Phosphoric acid/Phosphate | Acidic | 7.21 | 6.21 – 8.21 | Biochemistry |
| Ammonia/Ammonium | Alkaline | 9.26 | 8.26 – 10.26 | Laboratory |
| Carbonic acid/Bicarbonate | Neutral | 6.35 | 5.35 – 7.35 | Blood/Physiology |
Key Points to Remember
- Buffers contain a weak acid and its conjugate base (or weak base and its conjugate acid)
- Buffers resist pH changes by neutralizing added acid or base
- Buffer capacity increases with concentration
- Use Henderson-Hasselbalch equation: pH = pKa + log([salt]/[acid])
- Buffer works best when pH = pKa (ratio of 1:1)
- Effective buffer range is pKa ± 1
