7.2 Brønsted–Lowry theory of acids and bases
- Syllabus
- 9701–2028–2029
- Topic
- 7.2
- Level
- AS
Learning objectives
7.2.1Names and formulas of the common acids• State the names and formulas of the common acids, limited to hydrochloric acid, HCl, sulfuric acid, H2SO4, nitric acid, HNO3 and ethanoic acid, CH3COOH7.2.2Names and formulas of the common alkalis• State the names and formulas of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH and ammonia, NH37.2.3Brønsted–Lowry theory of acids and bases• Describe the Brønsted–Lowry theory of acids and bases7.2.4Strong acids and strong bases as fully• Describe strong acids and strong bases as fully dissociated in aq. solution and weak acids and weak bases as partially dissociated in aq. solution7.2.5Appreciate that, under standard conditions,• Appreciate that, under standard conditions, water has pH of 7 , acid solutions pH of below 7 and alkaline solutions pH of above 77.2.6The differences in behaviour between strong• Explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity7.2.7Neutralisation reactions occur when H +(aq)• Understand: neutralisation reactions occur when H +(aq) and OH–(aq) form H2O(l)7.2.8Salts are formed in neutralisation reactions• Understand: salts are formed in neutralisation reactions7.2.9Sketch the pH titration curves of titrations• Sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis7.2.10Select suitable indicators for acid-alkali• Select suitable indicators for acid-alkali titrations, given appropriate data (p Ka values will not be used)
Know the four syllabus acids by name and formula
| Common acid | Formula |
|---|---|
| hydrochloric acid | HCl |
| sulfuric acid | H₂SO₄ |
| nitric acid | HNO₃ |
| ethanoic acid | CH₃COOH |
Match the whole formula to the name. Ethanoic acid is the only listed organic acid and is commonly written CH₃COOH to keep its acidic carboxyl group visible.
This recall objective is limited to these four acids. Do not omit ethanoic acid, change sulfuric acid to sulfurous acid, or infer acid strength solely from the number of H atoms in a formula.
Know the three syllabus alkalis by name and formula
| Common alkali | Formula |
|---|---|
| sodium hydroxide | NaOH |
| potassium hydroxide | KOH |
| ammonia | NH₃ |
An alkali is a base that dissolves in water. Sodium hydroxide and potassium hydroxide contain OH⁻ directly; ammonia is still an alkali in aqueous solution because it reacts with water to generate OH⁻.
NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq)
Do not add OH to ammonia's formula: ammonia is NH₃, not NH₄OH for this syllabus recall. ‘Alkali’ is not a synonym for every base; solubility in water matters.
A Brønsted–Lowry acid donates H⁺ and a base accepts H⁺
In the Brønsted–Lowry model, an acid is a proton, H⁺, donor and a base is a proton acceptor. An acid–base reaction is therefore identified by tracking transfer of one H⁺ between species.
HA+B⇌A−+HB+
HA donates H⁺ and becomes its conjugate base A⁻; B accepts H⁺ and becomes its conjugate acid HB⁺. Each conjugate pair differs by exactly one H⁺.
In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, NH₃ accepts H⁺ and is the base, while H₂O donates H⁺ and is the acid. The pairs are NH₃/NH₄⁺ and H₂O/OH⁻.
A base need not contain OH⁻: NH₃ qualifies through proton acceptance. Acid/base labels can depend on the reaction partner; identify the transferred proton rather than relying only on a memorised substance label.
Strength is the extent of dissociation in aqueous solution
| Type | Behaviour in aqueous solution | Equation convention |
|---|---|---|
| strong acid | fully or essentially fully dissociated | HA(aq) → H⁺(aq) + A⁻(aq) |
| weak acid | partially dissociated | HA(aq) ⇌ H⁺(aq) + A⁻(aq) |
| strong base | fully or essentially fully dissociated into ions | MOH(aq) → M⁺(aq) + OH⁻(aq) |
| weak base | partially reacts with water to form ions | B(aq) + H₂O(l) ⇌ BH⁺(aq) + OH⁻(aq) |
HCl, HNO₃ and H₂SO₄ are strong acids in this context; CH₃COOH is weak. NaOH and KOH are strong bases, while NH₃ is weak.
Strength is not concentration. A strong solution may be dilute and a weak solution may be concentrated: strength describes the fraction dissociated, while concentration describes amount per volume.
At standard conditions, pH 7 is neutral
| Aqueous solution at standard conditions | pH | Relative H⁺ and OH⁻ concentrations |
|---|---|---|
| acidic | below 7 | [H⁺] > [OH⁻] |
| pure water / neutral | 7 | [H⁺] = [OH⁻] |
| alkaline | above 7 | [OH⁻] > [H⁺] |
Moving to a lower pH means a higher hydrogen-ion concentration and greater acidity; moving to a higher pH means greater alkalinity. The classification applies to aqueous solutions under the stated standard conditions.
pH 7 is the standard-condition neutral reference used by this syllabus. Do not confuse ‘neutral’ with ‘contains no ions’: water contains equal small concentrations of H⁺ and OH⁻.
At equal concentration, stronger acids show more ions and faster initial reactions
Compare strong and weak acids at the same concentration and temperature. The strong acid is much more fully dissociated, so it has a higher concentration of mobile H⁺ and counter-ions than the weak acid.
| Observation | Strong acid compared with an equally concentrated weak acid | Explanation |
|---|---|---|
| pH meter | lower pH | higher [H⁺] |
| universal indicator | colour further toward the acidic/red end | responds to lower pH |
| conductivity | higher | more mobile ions carry charge |
| reaction with the same reactive metal | faster initial bubbling / hydrogen production | more frequent effective collisions involving H⁺ |
Mg(s)+2H+(aq)→Mg2+(aq)+H2(g)
The comparison requires controlled concentration, volume, temperature and metal surface area. Faster initial reaction does not necessarily mean more final H₂: if equal stoichiometric amounts of two monoprotic acids react completely with excess metal, both can eventually supply the same total protons.
Neutralisation forms water from aqueous H⁺ and OH⁻
H+(aq)+OH−(aq)→H2O(l)
In acid–alkali neutralisation, H⁺ from the acid and OH⁻ from the alkali combine in a 1:1 ratio to form water. Other unchanged ions are spectators in the net ionic equation.
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). Removing spectator Na⁺ and Cl⁻ leaves the net ionic equation above.
Neutralisation describes the reaction, not a guarantee that the final mixture has pH 7. Excess reagent or a weak-acid/weak-base conjugate ion can make the final solution acidic or alkaline.
The acid supplies the anion and the base supplies the salt cation
| Neutralisation | Salt formed | Ion sources |
|---|---|---|
| HCl + NaOH → NaCl + H₂O | sodium chloride | Na⁺ from alkali; Cl⁻ from acid |
| HNO₃ + KOH → KNO₃ + H₂O | potassium nitrate | K⁺ from alkali; NO₃⁻ from acid |
| CH₃COOH + NaOH → CH₃COONa + H₂O | sodium ethanoate | Na⁺ from alkali; CH₃COO⁻ from acid |
| HCl + NH₃ → NH₄Cl | ammonium chloride | NH₄⁺ formed when NH₃ accepts H⁺; Cl⁻ from acid |
Name the salt from the base-derived cation followed by the acid-derived anion: hydrochloric acid gives chlorides, nitric acid nitrates, sulfuric acid sulfates and ethanoic acid ethanoates.
Balance charges when writing a salt formula. Salt formation does not require the final solution to be pH 7, and neutralisation by ammonia forms ammonium ions without producing OH⁻ as a written reactant.
Sketch titration curves from acid/base strength and addition direction
Put pH on the vertical axis and volume of titrant added on the horizontal axis. Mark the initial pH, a smooth approach to the stoichiometric equivalence volume, the steep or shallow inflection region, and the final pH set by excess titrant.
Assuming alkali is added to acid:
| Acid + alkali | Initial region | Equivalence region | After excess alkali |
|---|---|---|---|
| strong + strong | very low initial pH; gradual rise | large near-vertical jump centred at pH 7 | approaches high pH of strong alkali |
| weak + strong | higher initial pH; broad gradual/buffer-like rise | vertical jump with equivalence pH above 7 | approaches high pH of strong alkali |
| strong + weak | very low initial pH; gradual rise | smaller vertical jump with equivalence pH below 7 | approaches the lower pH of weak alkali |
| weak + weak | moderately acidic start and gradual rise | no sharp vertical section; equivalence pH depends on relative strengths | approaches the pH of excess weak alkali |
If acid is added to alkali, reverse the curve vertically: it starts alkaline and falls as volume is added. The equivalence volume is fixed by stoichiometric moles, not by where pH happens to equal 7.
Do not force every equivalence point to pH 7 or draw a sharp vertical jump for weak acid–weak alkali. This syllabus asks for qualitative sketches; pKa is not needed to place the required curve features.
Choose an indicator range wholly inside the steep pH change
Given a titration curve and indicator transition ranges, choose an indicator whose complete colour-change range lies within the steep section around the equivalence point. A small added volume then carries the indicator through its range, minimising endpoint error.
| Titration | Suitable choice from common indicators | Reason |
|---|---|---|
| strong acid + strong alkali | methyl orange (pH 3.1–4.4) or phenolphthalein (pH 8.3–10.0) | the wide vertical section spans both ranges |
| strong acid + weak alkali | methyl orange | acidic transition range lies within the vertical section |
| weak acid + strong alkali | phenolphthalein | alkaline transition range lies within the vertical section |
| weak acid + weak alkali | neither | no sufficiently steep section gives a sharp indicator endpoint |
The endpoint is the observed indicator colour change; the equivalence point is the stoichiometric point. A suitable range makes them close, but the two terms are not definitions of the same event.
Select from the supplied transition data rather than memorising an indicator name in isolation. No indicator pKa calculation is required, and a range merely containing the equivalence pH is insufficient if it extends outside the steep section.