Strong vs Weak Acids and Bases

A strong acid or base ionizes (or dissociates) completely in water, while a weak acid or base ionizes only partially and reaches an equilibrium between the molecules and their ions. In brief: 0.100 mol/L HCl has [H⁺] = 0.100 mol/L and pH 1.00, but 0.100 mol/L acetic acid has only about 1.3% of its molecules ionized, [H⁺] = 1.3 × 10⁻³ mol/L and pH 2.87.

Strength describes how far an acid or base reacts with water, and it is different from concentration. This article lists the common strong acids and bases, explains partial ionization with equilibrium constants, compares solutions of equal concentration in a table, shows how to calculate the percent ionization, works through examples, discusses the factors that determine strength and gives some history.

Which acids and bases are strong?

The common strong acids are hydrochloric (HCl), hydrobromic (HBr), hydroiodic (HI), nitric (HNO₃), sulfuric (H₂SO₄, first ionization only), perchloric (HClO₄) and chloric (HClO₃) acids. The strong bases are the hydroxides of group 1 (LiOH, NaOH, KOH, RbOH, CsOH) and of calcium, strontium and barium (Ca(OH)₂, Sr(OH)₂, Ba(OH)₂), which dissociate completely for the part that dissolves. All other common acids, such as HF, CH₃COOH, H₂CO₃ and H₃PO₄, and bases, such as NH₃ and amines, are weak.

What happens in a weak acid or base?

A weak acid HA establishes the equilibrium HA + H₂O ⇌ H₃O⁺ + A⁻, with K_a = [H⁺][A⁻] ÷ [HA]. A small K_a means that few molecules ionize. A weak base B has B + H₂O ⇌ BH⁺ + OH⁻, with K_b = [BH⁺][OH⁻] ÷ [B]. Most of the weak acid remains as HA, so the solution contains mainly molecules and a few ions, and conducts electricity poorly.

Strong and weak compared

The last row shows that strength does not change the total amount of base needed: a weak acid supplies more H⁺ as the equilibrium shifts when the base removes them.

How do you calculate percent ionization?

percent ionization = ([H⁺] at equilibrium ÷ initial concentration of HA) × 100%. For a weak acid with concentration c and small ionization, [H⁺] ≈ √(K_ac). Percent ionization increases when the solution is diluted.

Worked examples

1. Acetic acid, 0.100 mol/L. K_a = 1.8 × 10⁻⁵. [H⁺] ≈ √(1.8 × 10⁻⁵ × 0.100) = 1.34 × 10⁻³ mol/L. pH = 2.87. Ionization = 1.34%.
2. Dilution effect. At 0.0100 mol/L: [H⁺] = √(1.8 × 10⁻⁷) = 4.2 × 10⁻⁴ mol/L, pH 3.37, ionization 4.2%. A tenfold dilution changed the ionization from 1.3% to 4.2%.
3. Hydrofluoric acid, 0.0500 mol/L. K_a = 6.8 × 10⁻⁴ is too large for the approximation, so solve x² + K_ax − K_ac = 0: x = 5.5 × 10⁻³ mol/L, pH 2.26, ionization 11%.
4. Ammonia, 0.100 mol/L. K_b = 1.8 × 10⁻⁵. [OH⁻] = 1.34 × 10⁻³ mol/L, pOH 2.87, pH 11.13.
5. Strong base. 0.0010 mol/L NaOH: [OH⁻] = 1.0 × 10⁻³ mol/L, pOH 3.00, pH 11.00.
6. Compare pH. 0.100 mol/L solutions: HCl pH 1.00, HF pH 2.1, acetic acid pH 2.87, HCN pH 5.1. The stronger the acid, the lower the pH at equal concentration.

What determines the strength of an acid?

• Bond strength: down group 17, H–X bonds get weaker, so HI > HBr > HCl >> HF in strength. HF is weak despite the high electronegativity of fluorine because the H–F bond is very strong.
• Polarity of the bond: across a period, acidity increases with electronegativity: CH₄ < NH₃ < H₂O < HF.
• Stability of the conjugate base: more stable anions (charge delocalized by resonance or spread over a large atom) mean stronger acids.
• Oxyacids: more oxygen atoms on the central atom give a stronger acid: HClO < HClO₂ < HClO₃ < HClO₄.

How was strength measured?

Wilhelm Ostwald studied the conductivity of acids in the 1880s, and showed that the conductivity of weak acids varies with dilution in a way that fits the equilibrium law, leading to his dilution law (1888). The acid dissociation constant became the standard measure of strength, with Arrhenius’s ionization theory as a basis.

Which mistakes are common?

Mistake 1: Confusing strong with concentrated

Strength is about ionization. Concentration is about the amount per volume.

Mistake 2: Using the approximation when ionization is large

If x is more than 5% of the initial concentration, use the quadratic equation.

Mistake 3: Writing a weak acid with a single arrow

Use ⇌ for weak acids and bases.

Mistake 4: Assuming that a weak acid needs less base to neutralize it

Equal moles of monoprotic acids need equal moles of base, whatever their strength.