Topic 1 of 3

Models and strength

Identify roles and distinguish concentration from dissociation.

A-Level 8873, revised syllabus (2026-2027)

Follow what the acid and base actually do

One model focuses on water; another follows proton transfer.

In the Arrhenius model, an acid produces H⁺(aq) in water and a base produces OH^-(aq). HCl(aq) provides H⁺ and Cl^-; NaOH(aq) provides Na⁺ and OH^-. Neutralisation removes H⁺ and OH^- as water: H⁺(aq) + OH^-(aq) → H₂O(l).

The Bronsted-Lowry model follows a proton: an acid donates H⁺, and a base accepts it. A conjugate acid-base pair differs by exactly one proton. Losing that proton also reduces the signed charge by one; gaining it increases the signed charge by one.

Track the transferred proton
Reaction	Acid and its conjugate base	Base and its conjugate acid
HCl + H₂O → H₃O⁺ + Cl^-	HCl / Cl^-	H₂O / H₃O⁺
NH₃ + H₂O ⇌ NH₄⁺ + OH^-	H₂O / OH^-	NH₃ / NH₄⁺
CH₃CO₂H + H₂O ⇌ CH₃CO₂^- + H₃O⁺	CH₃CO₂H / CH₃CO₂^-	H₂O / H₃O⁺

Water accepts a proton from HCl but donates one to NH₃; a substance capable of either role is amphiprotic. In aqueous equations H⁺(aq) is convenient shorthand for the hydrated proton. Writing H₃O⁺ makes the transfer to water visible.

Worked example

Find a pair in an unfamiliar reaction

HCO₃^- + H₂O ⇌ CO₃^2- + H₃O⁺. Which species donates the proton?

HCO₃^- loses one H⁺ and becomes CO₃^2-. Its charge changes from -1 to -2.
Water gains that proton and becomes H₃O⁺.
Pair each species with the product differing by one proton.

Answer

HCO₃^- is the acid; CO₃^2- is its conjugate base. H₂O is the base and H₃O⁺ its conjugate acid.

Strength describes dissociation, not concentration

Separate how much acid was dissolved from how much becomes ions.

A strong acid is essentially fully dissociated in dilute aqueous solution. A weak acid establishes an equilibrium in which only a proportion is dissociated. Strong and weak bases differ similarly in the extent to which they generate OH^- in water. "Concentrated" and "dilute" instead describe amount per volume. A concentrated weak acid is still weak.

Equal analytical concentrations, different dissociation

Qualitative sketch: the number of printed symbols is illustrative, not a measured dissociation fraction.

Compare equal analytical concentrations at the same temperature
Comparison	Expected observation	Dissociation explanation
Dilute HCl versus ethanoic acid, both monoprotic	HCl has lower pH and usually greater electrical conductivity.	HCl supplies a much larger concentration of H+ and other mobile ions; most ethanoic acid remains as neutral molecules.
Equal Mg pieces in excess of these two acids	Hydrogen is normally released faster initially in HCl.	The larger initial H+ concentration gives more frequent effective collisions at the same exposed metal surface. Control temperature and surface area.
Fully dissolved NaOH versus aqueous NH3	NaOH has higher pH at the same analytical concentration.	NaOH dissociates essentially completely to give OH-; NH3 reacts only partially with water to form NH4+ and OH-.

A weak acid can still react completely with enough strong base. As H⁺ is removed, more weak acid dissociates. Equal volumes of equal-concentration monoprotic strong and weak acids therefore require the same amount of NaOH for complete neutralisation, even though their initial pH and initial reaction rates differ. Comparing pH alone cannot establish acid strength unless concentration and the number of ionisable protons are also controlled.

For a monoprotic weak acid HA(aq) ⇌ H⁺(aq) + A^-(aq), K_a = [H⁺][A^-]/[HA] at equilibrium. At a fixed temperature, a larger K_a means a greater tendency to donate a proton. For a base B + H₂O ⇌ BH⁺ + OH^-, K_b = [BH⁺][OH^-]/[B]. Pure liquid water is omitted from these expressions.

pH = -log₁₀[H⁺(aq)], with the concentration expressed in mol dm^-3 for these calculations. One pH unit corresponds to a tenfold change in hydrogen-ion concentration. K_w = [H⁺][OH^-] is the ionic product of water. At 25 degrees C, K_w = 1.00 × 10^-14 mol² dm^-6.