Acid-Base Reactions Explained

Proton Transfer: The Simplest Explanation of Acid-Base Chemistry

Most of chemistry's most important reactions — from cellular respiration to industrial synthesis — involve proton transfer. Acids donate protons (H⁺ ions); bases accept them. When they meet, they neutralize each other. This Brønsted-Lowry definition, developed independently by Johannes Brønsted and Thomas Lowry in 1923, describes most of the acid-base chemistry you'll encounter in courses and real-world applications.

Here's how it works, with the numbers that make it concrete.

The Brønsted-Lowry Model: Proton Donors and Acceptors

Acid: A proton donor. Releases H⁺ in solution.

• HCl → H⁺ + Cl⁻ (strong acid, dissociates 100% in water)
• CH₃COOH ⇌ H⁺ + CH₃COO⁻ (weak acid, partial dissociation)

Base: A proton acceptor. Accepts H⁺ in solution.

• NaOH → Na⁺ + OH⁻ (strong base, dissociates 100%)
• NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (weak base, accepts a proton from water)

The neutralization reaction: When HCl and NaOH meet: H⁺ + Cl⁻ + Na⁺ + OH⁻ → Na⁺ + Cl⁻ + H₂O

The net ionic equation: H⁺ + OH⁻ → H₂O

Salt and water form. The pH of the resulting solution depends on the strength of the original acid and base.

The pH Scale: What the Numbers Mean

pH = -log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in moles per liter.

• pH 0–6.9: Acidic ([H⁺] > [OH⁻])
• pH 7: Neutral ([H⁺] = [OH⁻] = 1×10⁻⁷ M) — pure water at 25°C
• pH 7.1–14: Basic ([OH⁻] > [H⁺])

Key examples:

• Battery acid (H₂SO₄, concentrated): pH ≈ 0–1
• Lemon juice (citric acid): pH ≈ 2.2
• Black coffee: pH ≈ 5
• Blood (bicarbonate buffer): pH 7.35–7.45 (narrow range critical for enzyme function)
• Seawater: pH 8.1
• Baking soda solution: pH 8.3
• Bleach (sodium hypochlorite): pH 12–13
• Drain cleaner (NaOH solution): pH ≈ 14

The logarithmic scale means each unit represents a 10× change. pH 4 is 10× more acidic than pH 5 and 100× more acidic than pH 6. A blood pH of 7.0 (instead of the normal 7.4) represents a 2.5× increase in [H⁺] — enough to be life-threatening.

Strong vs. Weak Acids and Bases

Strong acids dissociate completely:

• HCl (hydrochloric acid): 100% → H⁺ + Cl⁻
• HNO₃ (nitric acid): 100% → H⁺ + NO₃⁻
• H₂SO₄ (sulfuric acid, first dissociation): 100% → H⁺ + HSO₄⁻

For a 0.1 M HCl solution: [H⁺] = 0.1 M, pH = -log(0.1) = 1.0

Weak acids dissociate partially: The equilibrium constant for dissociation is Ka (acid dissociation constant).

Acetic acid (vinegar): CH₃COOH ⇌ H⁺ + CH₃COO⁻, Ka = 1.8 × 10⁻⁵

For a 0.1 M acetic acid solution: [H⁺] = √(Ka × C) = √(1.8×10⁻⁵ × 0.1) = √(1.8×10⁻⁶) = 1.34×10⁻³ M pH = -log(1.34×10⁻³) = 2.87

Compare: 0.1 M HCl = pH 1.0; 0.1 M acetic acid = pH 2.87. Same concentration, but strong acid is ~75× more acidic.

Titration: Determining Unknown Concentration

Titration uses a solution of known concentration (titrant) to determine the concentration of an unknown solution (analyte). The neutralization point — where moles of acid equal moles of base — is called the equivalence point.

Example: Titrating 25.0 mL of 0.100 M HCl with 0.100 M NaOH

At equivalence point: moles HCl = moles NaOH n(HCl) = 0.100 M × 0.0250 L = 0.00250 mol Volume NaOH needed: 0.00250 mol ÷ 0.100 M = 0.0250 L = 25.0 mL

The equivalence point occurs when exactly 25.0 mL of NaOH has been added. At this point, the solution contains only NaCl and water — pH = 7.00.

The dramatic pH change near equivalence: This is the crucial feature of the titration curve. Consider what happens between 24.9 mL and 25.1 mL:

• At 24.9 mL NaOH added: 0.1 mL of HCl remains unreacted
  • [H⁺] = (0.1 mL × 0.100 M) ÷ 50.1 mL total = 2.0×10⁻⁴ M → pH = 3.70
• At 25.0 mL: equivalence, pH = 7.00
• At 25.1 mL NaOH added: 0.1 mL of NaOH in excess
  • [OH⁻] = (0.1 mL × 0.100 M) ÷ 50.1 mL total = 2.0×10⁻⁴ M
  • pOH = 3.70, pH = 14 - 3.70 = 10.30

Adding just 0.2 mL of NaOH (0.2% of the total volume) causes the pH to jump from 3.70 to 10.30 — a change of 6.6 pH units. This is the equivalence point jump visible in all strong acid-strong base titration curves.

Indicators: pH indicators are weak acids that change color based on pH. Phenolphthalein changes from colorless (below pH 8.2) to pink (above pH 10). For a strong acid-strong base titration, this works well — the equivalence point at pH 7 is within the indicator's transition range during the steep jump.

Buffers: Resisting pH Change

A buffer is a solution that resists pH change when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (or weak base and its conjugate acid) in roughly equal concentrations.

Why blood is buffered: Human blood must stay at pH 7.35–7.45. Deviation to pH 7.0 (acidosis) or pH 7.8 (alkalosis) is potentially fatal. The bicarbonate buffer system (H₂CO₃/HCO₃⁻) maintains this range despite the acids produced by cellular metabolism.

The Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA])

where pKa = -log(Ka), [A⁻] = concentration of conjugate base, [HA] = concentration of weak acid.

Example: Acetate buffer Acetic acid (Ka = 1.8×10⁻⁵, pKa = 4.74) and sodium acetate:

• Equal concentrations (0.1 M each): pH = 4.74 + log(1) = 4.74 + 0 = 4.74
• 2× more acetate than acetic acid: pH = 4.74 + log(2) = 4.74 + 0.30 = 5.04

Buffers work best when pH is within ±1 of pKa. Outside this range, the buffer capacity is insufficient to resist pH change.

Buffer capacity: A buffer with 0.1 M each of weak acid and conjugate base can neutralize approximately 0.05 mol/L of added strong acid or base before the pH changes significantly. Higher concentrations provide more capacity.

Common Real-World Examples

Stomach acid: HCl solution, pH 1.5–3.5. Strong enough to denature proteins and activate pepsin. Antacids (CaCO₃, MgOH₂) neutralize excess stomach acid: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂.

Baking: Baking soda (NaHCO₃, a base) reacts with acidic ingredients (buttermilk, vinegar, brown sugar) to produce CO₂ bubbles that leaven bread. If the recipe lacks acid, baking powder (NaHCO₃ + cream of tartar) provides its own acid component.

Acid rain: SO₂ and NOₓ from combustion dissolve in rainwater to form H₂SO₃ and HNO₃. Normal rain is pH 5.6 (slightly acidic due to dissolved CO₂); acid rain is pH 4–5, acidic enough to leach metals from soil and damage aquatic ecosystems.