Strong Acids and Bases

A strong acid or strong base dissociates completely in water. There is no equilibrium, no Ka, no partial reaction. Every molecule breaks apart into ions. This makes pH calculations for strong acids and bases trivially easy - but first, you need to memorize which ones are strong.

The Seven Strong Acids

There are exactly seven strong acids. Everything else is weak.

Why HF Is Weak

Students are often surprised that HF is not a strong acid. Fluorine is the most electronegative element - shouldn’t it pull the proton away most easily?

The answer lies in bond strength, not electronegativity. The H-F bond is extremely short and strong (bond energy 570 kJ/mol vs. 431 kJ/mol for H-Cl). This bond is so tough to break that HF only partially dissociates in water. Electronegativity determines polarity, but bond strength determines how easily the proton leaves.

The Strong Bases

Strong bases are the hydroxides of Group 1 metals and the heavier Group 2 metals. They dissociate completely in water.

What “Complete Dissociation” Means for Calculations

For a strong acid like HCl at concentration c:

• HCl → H⁺ + Cl⁻ (100% dissociation)
• [H⁺] = c
• No Ka needed, no ICE table needed

For a strong base like NaOH at concentration c:

• NaOH → Na⁺ + OH⁻ (100% dissociation)
• [OH⁻] = c

For Ca(OH)₂ at concentration c:

• Ca(OH)₂ → Ca²⁺ + 2 OH⁻
• [OH⁻] = 2c (two hydroxides per formula unit)

Sulfuric Acid - A Special Case

H₂SO₄ is diprotic. Its first dissociation is strong (complete), but its second dissociation is weak ($K_{a2} = 1.2 \times 10^{-2}$).

• H₂SO₄ → H⁺ + HSO₄⁻ (strong, 100%)
• HSO₄⁻ ⇌ H⁺ + SO₄²⁻ (weak, Ka₂ = 0.012)

For dilute H₂SO₄ solutions on the MCAT, you can usually approximate that both protons dissociate fully, giving [H⁺] ≈ 2c. But for more concentrated solutions or precise calculations, only the first dissociation is truly complete.