Acids, Bases, and the pH Scale (Without the Confusion)

Acid-base chemistry confuses students because of the multiple definitions. Stick with one — the Brønsted-Lowry definition — and most of the topic gets simpler.

Brønsted-Lowry definition

• An acid is a proton (H⁺) donor
• A base is a proton acceptor

When HCl dissolves in water, it donates H⁺ to water:

HCl + H₂O → H₃O⁺ + Cl⁻

HCl is the acid. Water is the base.

The pH scale

pH = -log₁₀[H⁺]

• pH 7: neutral (pure water)
• pH < 7: acidic (more H⁺)
• pH > 7: basic (more OH⁻)

Each unit of pH is a 10× change in H⁺ concentration. pH 3 is ten times more acidic than pH 4.

Strong vs weak

• Strong acid completely ionises in water (HCl, HNO₃, H₂SO₄)
• Weak acid only partially ionises (CH₃COOH — vinegar)
• Strong base completely ionises (NaOH, KOH)
• Weak base partially ionises (NH₃)

This is why a 1 M solution of HCl has pH ≈ 0 but a 1 M solution of acetic acid has pH ≈ 2.4 — same concentration, different ionisation.

Neutralisation

Acid + base → salt + water.

The H⁺ from the acid combines with the OH⁻ from the base to form water. The other ions form a salt.

HCl + NaOH → NaCl + H₂O

What examiners ask

• Calculate pH from H⁺ concentration
• Identify the conjugate acid/base pair
• Predict the salt from a neutralisation
• Distinguish strong from weak

Common pitfalls

• Confusing strong/weak with concentrated/dilute. They're independent.
• Forgetting to take the negative log
• Treating all bases as alkalis (alkalis are bases that dissolve in water)

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