Lesson 1 of 3
Models and strength
Identify roles and distinguish concentration from dissociation.
A-Level 8873, revised syllabus (2026-2027)
I can apply the Arrhenius model.Syllabus
Syllabus 8873, 3(a). Apply the Arrhenius model.
I can apply proton-transfer theory.Syllabus
Syllabus 8873, 3(b). Apply proton-transfer theory.
I can distinguish acid/base strength.Syllabus
Syllabus 8873, 3(c). Distinguish acid/base strength.
I can explain pH and equilibrium terms.Syllabus
Syllabus 8873, 3(d). Explain pH and equilibrium terms.
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.
| 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 exampleFind a pair in an unfamiliar reaction
HCO3- + H2O ⇌ CO32- + H3O+. Which species donates the proton?
Try it first. Then show the worked stepsHide the worked steps
- HCO3- loses one H+ and becomes CO32-. Its charge changes from -1 to -2.
- Water gains that proton and becomes H3O+.
- Pair each species with the product differing by one proton.
HCO3- is the acid; CO32- is its conjugate base. H2O is the base and H3O+ its conjugate acid.
Check your understanding
In NH3 + H2O <=> NH4+ + OH-, which species acts as the Bronsted-Lowry acid?
Show the answer
H2O, because it donates a proton to NH3
A Bronsted-Lowry acid donates H+. Water gives a proton to ammonia, forming OH-. The conjugate pairs are H2O/OH- and NH4+/NH3.
Strength describes dissociation, not concentration
Separate how much acid was dissolved from how much becomes ions.
Make a guess. You are not marked.
Two acids are both 0.100 mol dm-3: hydrochloric acid and ethanoic acid. Which solution has the higher pH?
Show the answer
Ethanoic acid, because it only partly ionises
Ethanoic acid is weak and only partly ionises, so its [H+] is much lower than the 0.100 mol dm-3 from fully ionised HCl.
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
Strong acid is mostly ions; weak acid is mostly undissociated HA alongside some ions. Water and exact particle counts are omitted.
| 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), Ka = [H+][A-]/[HA] at equilibrium. At a fixed temperature, a larger Ka means a greater tendency to donate a proton. For a base B + H2O ⇌ BH+ + OH-, Kb = [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. Kw = [H+][OH-] is the ionic product of water. At 25 degrees C, Kw = 1.00 × 10-14 mol2 dm-6.
Check your understanding
Which solution has the lowest pH?
Show the answer
0.10 mol dm-3 hydrochloric acid
HCl dissociates fully, so [H+] = 0.10 mol dm-3 and pH = 1.0. Ethanoic acid only partly dissociates, so even 1.0 mol dm-3 has [H+] well below 0.10.
Exam-style question
Separate 25.0 cm3 samples of 0.100 mol dm-3 hydrochloric acid and 0.100 mol dm-3 ethanoic acid are compared at 25 °C.
(a) Explain why the hydrochloric acid has the lower pH. [2]
(b) Equal pieces of magnesium are added. Predict and explain the difference in the initial rate. [1]
(c) Explain why both acids need the same volume of 0.100 mol dm-3 NaOH for neutralisation. [1] [4]
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