Common Acids and Alkalis to Memorise

· Hydrochloric acid = HCl.
· Sulfuric acid = H₂SO₄.
· Nitric acid = HNO₃.
· Ethanoic acid = CH₃COOH.
· Sodium hydroxide = NaOH.
· Potassium hydroxide = KOH.
· Ammonia = NH₃.
· Alkali = a soluble base that produces OH⁻(aq) in water.

Brønsted–Lowry Theory

· Brønsted–Lowry acid = proton donor.
· Brønsted–Lowry base = proton acceptor.
· A proton means H⁺.
· Acid–base reactions involve proton transfer from the acid to the base.
· Example: HCl + NH₃ → NH₄⁺ + Cl⁻.
· In this reaction: HCl donates H⁺, so HCl is the acid; NH₃ accepts H⁺, so NH₃ is the base.
· When explaining, always identify which species donates H⁺ and which species accepts H⁺.

This diagram shows that a Brønsted–Lowry acid donates H⁺ to a base. It is useful for visualising proton transfer rather than just memorising the definition. Source

Strong and Weak Acids/Bases

· Strong acid = fully dissociated in aqueous solution.
· Example: HCl(aq) → H⁺(aq) + Cl⁻(aq).
· Weak acid = partially dissociated in aqueous solution.
· Example: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq).
· Strong base = fully dissociated in aqueous solution.
· Example: NaOH(aq) → Na⁺(aq) + OH⁻(aq).
· Weak base = partially dissociated in aqueous solution.
· Example: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq).
· Strong/weak describes degree of dissociation, not concentration.
· Concentrated/dilute describes amount of solute per volume.

pH and Measuring Acidity/Alkalinity

· Water has pH 7.
· Acidic solutions have pH below 7.
· Alkaline solutions have pH above 7.
· Lower pH = greater acidity.
· Higher pH = greater alkalinity.
· A pH meter gives a more precise pH value.
· Universal indicator estimates pH by colour.
· Conductivity is higher when there are more mobile ions in solution.
· At the same concentration, a strong acid usually has a lower pH and higher conductivity than a weak acid because it is more fully dissociated.

This pH scale diagram helps students remember that pH < 7 is acidic, pH = 7 is neutral, and pH > 7 is alkaline. It is useful for quick classification questions. Source

Strong vs Weak Acid Behaviour

· Strong acids react more vigorously with reactive metals than weak acids of the same concentration.
· This is because strong acids produce a greater concentration of H⁺(aq).
· With a reactive metal: acid + metal → salt + hydrogen.
· Example: Mg + 2HCl → MgCl₂ + H₂.
· Strong acids of the same concentration have lower pH values than weak acids.
· Strong acids of the same concentration usually show higher conductivity because they contain more ions.
· Universal indicator and pH meter results can be used to distinguish strong vs weak acids, but concentration must be considered.

Neutralisation and Salt Formation

· Neutralisation occurs when H⁺(aq) reacts with OH⁻(aq) to form water.
· Ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l).
· Salt = ionic compound formed when the H⁺ of an acid is replaced by a metal ion or NH₄⁺.
· General equation: acid + alkali → salt + water.
· Example: HCl + NaOH → NaCl + H₂O.
· For sulfuric acid: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.
· Always balance equations, especially for diprotic acids such as H₂SO₄.

pH Titration Curves

· A pH titration curve shows how pH changes as acid/alkali is added.
· You must be able to sketch curves for combinations of strong/weak acids and strong/weak alkalis.
· Strong acid + strong alkali: steep vertical section; equivalence point around pH 7.
· Weak acid + strong alkali: starts at a higher pH than a strong acid; equivalence point above pH 7.
· Strong acid + weak alkali: starts low; equivalence point below pH 7.
· Weak acid + weak alkali: smaller pH change; endpoint is less sharp, so indicator choice is harder.
· The steep vertical region is where the best indicator must change colour.
· Exam sketches should include labelled axes: pH on the y-axis and volume of acid/alkali added on the x-axis.

This diagram compares the four main titration curve shapes required for this topic. It is useful for spotting how acid/base strength affects starting pH, equivalence pH, and steepness of the endpoint. Source

Choosing Indicators for Titrations

· An indicator is suitable if its colour-change pH range lies within the steep vertical section of the titration curve.
· Strong acid + strong alkali: many indicators work because the vertical section is large.
· Weak acid + strong alkali: use an indicator changing in the alkaline range, e.g. phenolphthalein if data supports it.
· Strong acid + weak alkali: use an indicator changing in the acidic range, e.g. methyl orange if data supports it.
· Weak acid + weak alkali: usually no suitable indicator because there is no sharp pH change.
· In exams, use the given indicator data; pKa values are not required for this syllabus point.
· Do not select an indicator by memorising only; justify using the indicator range and vertical part of the curve.

These curves show why indicator choice depends on where the rapid pH change occurs. They support the exam rule: the indicator’s colour-change range must fall within the steep vertical section of the titration curve. Source

Common Exam Mistakes to Avoid

· Do not write “strong acid = concentrated acid”; strength and concentration are different.
· Do not say weak acids do not react; they partially dissociate and still produce H⁺(aq).
· Do not forget state symbols in ionic equations when required.
· Do not choose indicators randomly; match the indicator range to the steep part of the titration curve.
· Do not assume all neutralisations have equivalence point pH 7; this is only typical for strong acid + strong alkali.
· Do not confuse endpoint with equivalence point: endpoint is the observed indicator colour change; equivalence point is when acid/base have reacted in the exact stoichiometric ratio.