01 · Acids and bases — definitions
The most widely used model is Brønsted–Lowry (1923):
- An acid is a proton donor — it releases H⁺ ions in solution.
- A base is a proton acceptor — it takes up H⁺ ions.
An older model, Arrhenius, defines acids as substances that increase [H⁺] in water and bases as substances that increase [OH⁻]. The broader Lewis model defines acids as electron-pair acceptors — useful in organic and inorganic chemistry beyond aqueous solutions.
See our atomic structure guide to understand why protons (H⁺) behave differently to other ions.
02 · The pH scale
pH measures the concentration of H⁺ ions in a solution:
pH = −log₁₀[H⁺]
The scale runs from 0 to 14 at standard conditions. Each unit is a tenfold change in [H⁺]. A solution at pH 3 has ten times more H⁺ than one at pH 4, and a hundred times more than one at pH 5.
Explore the full interactive scale and data sources in our pH examples chart.
03 · Strong vs weak acids
Strong acids dissociate completely in water — every molecule releases its H⁺. Weak acids only partially dissociate; an equilibrium forms between the acid and its ions.
Strong acids
- Fully dissociate in water
- HCl, HNO₃, H₂SO₄
- Low pH at same concentration
Weak acids
- Partially dissociate — equilibrium
- CH₃COOH (ethanoic acid)
- Higher pH at same concentration
Concentration and strength are different. A concentrated weak acid can have a lower pH than a dilute strong acid — never confuse the two.
04 · Neutralisation and salts
When an acid reacts with a base, they neutralise each other to form a salt and water:
HCl + NaOH → NaCl + H₂O
The H⁺ from the acid and the OH⁻ from the base combine to form water. The remaining ions (Na⁺ and Cl⁻) form the salt sodium chloride. This is the basis of titration — a lab technique for finding an unknown concentration.
Buffers exploit this equilibrium: a weak acid and its salt resist pH change when small quantities of strong acid or base are added. Blood uses carbonate buffers to stay at pH 7.35–7.45.