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R3.1

Proton transfer reactions

8 SL · +9 HL (SL 3.1.1-3.1.8) (HL 3.1.9-3.1.17)

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Parent topic: What are the mechanisms of chemical change?

Guiding question

Guiding questionWhat happens when protons are transferred?

The one big idea

Acids donate protons, bases accept them. The strength of an acid is how far it ionises; pH measures the result. At HL, Ka, Kb, pKa, pKb, buffer chemistry and indicator theory turn the qualitative model into a quantitative one.

What this rests on

Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.

What rests on this

No dependents recorded.

Core concepts that must be mastered

  • Brønsted–Lowry acids and bases(3.1.1–3.1.3)A Brønsted–Lowry acid is a proton donor; a base is a proton acceptor. A proton in aqueous solution may be written as H⁺(aq) or H₃O⁺(aq). Note the distinction between base and alkali: an alkali is a base that dissolves in water to produce OH⁻(aq). Ammonia is a base but not an alkali until it dissolves. A conjugate acid–base pair differs by a single proton: NH₃/NH₄⁺, H₂O/OH⁻, CH₃COOH/CH₃COO⁻. You must deduce the conjugate acid or base of any given species. Some species are amphiprotic — they can donate or accept a proton: HCO₃⁻, H₂O, HSO₄⁻, H₂PO₄⁻.
  • pH, Kw and strong/weak acids(3.1.4–3.1.6)pH = −log₁₀[H⁺] and [H⁺] = 10⁻ᵖᴴ. You must perform calculations with this logarithmic relationship. A pH change of 1 means a tenfold change in [H⁺]; a change of 2 means a hundredfold. pH can be negative (concentrated strong acids) — it is not bounded between 0 and 14. Kw = [H⁺][OH⁻]; at 298 K, Kw = 1.0 × 10⁻¹⁴. In any aqueous solution: acidic means [H⁺] > [OH⁻]; neutral means [H⁺] = [OH⁻]; basic means [H⁺] < [OH⁻]. Water is always neutral ([H⁺] = [OH⁻]), but Kw increases with temperature because water's ionisation is endothermic — so at higher temperatures neutral water has pH below 7 (still neutral, because [H⁺] still equals [OH⁻]). Strong and weak acids/bases differ in the extent of ionisation, not in concentration. Strong acids are fully ionised (HCl, HNO₃, H₂SO₄, HClO₄, HBr, HI); weak acids are partially ionised (ethanoic acid, carbonic acid). Strong does not mean concentrated, and weak does not mean dilute: 0.001 mol dm⁻³ HCl is a dilute strong acid; 10 mol dm⁻³ ethanoic acid is a concentrated weak acid. Acid–base equilibria lie toward the weaker conjugate: a strong acid has a weak conjugate base, and vice versa.
  • Neutralization and pH curves(3.1.7, 3.1.8)Acids react with bases in neutralization. You must formulate equations for reactions of acids with metal oxides, metal hydroxides, hydrogencarbonates and carbonates. Examples: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l); 2HCl(aq) + Na₂CO₃(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g). You must identify the parent acid and base of different salts. pH curves for strong acid–strong base neutralisation have characteristic shapes. You must sketch and interpret the general shape, including the pH-axis intercept and the equivalence point (the stoichiometric point where moles of acid = moles of base). At the equivalence point of a strong acid–strong base titration, pH = 7. Only monoprotic neutralisations are assessed at SL.
  • HL: pOH, Ka, Kb, pKa, pKb and salt hydrolysis(3.1.9–3.1.12)pOH = −log₁₀[OH⁻] and [OH⁻] = 10⁻ᵖᴼᴴ. You must interconvert [H⁺], [OH⁻], pH and pOH using pH + pOH = 14 (at 298 K). Weak acid and base strengths are described by Ka, Kb, pKa, pKb. Lower pKa = larger Ka = stronger acid. For a conjugate pair, Ka × Kb = Kw — this lets you find Kb from Ka and vice versa. Quadratic equations are not expected; use approximations for weak acids where dissociation is small. The pH of a salt solution depends on the relative strengths of the parent acid and base. Only salts of a strong acid and strong base are neutral (NaCl, KNO₃). Salts of weak acids are basic (sodium ethanoate ≈ pH 9) because the anion hydrolyses: CH₃COO⁻(aq) + H₂O(l) ⇌ CH₃COOH(aq) + OH⁻(aq). Salts of weak bases are acidic (NH₄Cl ≈ pH 4.6) because the cation hydrolyses: NH₄⁺(aq) ⇌ NH₃(aq) + H⁺(aq). You must construct hydrolysis equations for NH₄⁺, RCOO⁻, CO₃²⁻, HCO₃⁻ and predict each ion's effect on pH. Acidity of hydrated transition element ions is not required.
  • HL: pH curves, indicators and buffers(3.1.13–3.1.17)At HL, pH curves extend to all four strong/weak monoprotic combinations: strong acid–strong base (steep, equivalence at pH 7), strong acid–weak base (equivalence below 7), weak acid–strong base (equivalence above 7), weak acid–weak base (no steep rise). You must interpret their general shapes, including the pH-axis intercept, the equivalence point, the buffer region (the shallow part where the solution resists pH change), and points where pH = pKa or pOH = pKb. Indicators are weak acids whose conjugate pair members have different colours. The end-point pH ≈ pKa of the indicator. A suitable indicator has an end-point range coinciding with the equivalence-point pH — distinguish the end point (where the indicator changes) from the equivalence point (the stoichiometric point). A buffer resists pH change on addition of small amounts of acid or alkali. An acidic buffer is a weak acid plus its conjugate base (e.g. CH₃COOH + CH₃COO⁻); a basic buffer is a weak base plus its conjugate acid (e.g. NH₃ + NH₄⁺). Buffer pH depends on pKa (or pKb) and the acid:conjugate-base concentration ratio: pH = pKa + log₁₀([A⁻]/[HA]). You must solve problems involving buffer composition and pH, including the effect of dilution: diluting a buffer changes both concentrations equally, so the ratio is unchanged and pH is essentially unchanged — but buffer capacity falls. A buffer resists change only for small additions; exceed the capacity of either component and pH moves sharply. Buffers must be weak acid/base systems, not strong ones.
  • How this sub-topic connectsProton transfer is the first of the three electron-transfer/sharing/sharing-pair mechanism sub-topics. The equilibrium constant Ka is a direct application of the equilibrium law (R2.3), and the small-K approximation used for weak acids is the same one from R2.3.6. Buffer action is Le Châtelier's principle in action (R2.3). Neutralization enthalpy connects to R1.1 (exothermic bond formation). At HL, the Lewis acid–base definition (R3.4) generalises Brønsted–Lowry: every Brønsted base is a Lewis base, but not every Lewis base is a Brønsted base (it need not accept a proton). The titration technique connects to Tool 1 and to the mole concept (S1.4, R2.1).

Quantitative non-negotiables

  • Deduce the Brønsted–Lowry acid and base in a reaction, and the conjugate acid–base pair.
  • Perform pH = −log₁₀[H⁺] and [H⁺] = 10⁻ᵖᴴ calculations.
  • Recognise acidic, neutral and basic solutions from relative [H⁺] and [OH⁻] using Kw = [H⁺][OH⁻].
  • Formulate neutralization equations for acids with metal oxides, hydroxides, carbonates and hydrogencarbonates.
  • Sketch and interpret pH curves for strong acid–strong base neutralisation (SL, monoprotic).
  • At HL: interconvert [H⁺], [OH⁻], pH and pOH.HL
  • At HL: use Ka, Kb, pKa, pKb and Ka × Kb = Kw to compare and calculate.HL
  • At HL: predict salt-solution pH from hydrolysis equations.HL
  • At HL: calculate buffer pH from composition and pKa.HL

Common failure modes

  • M-51 — Strong acid means concentrated acid.

    Confidence: verified

    Why it’s wrong: Strong = fully ionised (a property of the substance); concentrated = many moles per dm³ (a property of the solution).

    Correction: 0.001 mol dm⁻³ HCl is a dilute strong acid; 10 mol dm⁻³ ethanoic acid is a concentrated weak acid.

  • M-52 — pH must be between 0 and 14.

    Confidence: verified

    Why it’s wrong: pH = −log₁₀[H⁺] with no bounds.

    Correction: Concentrated strong acids give negative pH.

  • M-53 — All salt solutions are neutral.

    Confidence: verified

    Why it’s wrong: Only salts of a strong acid and a strong base are neutral.

    Correction: Salts of weak acids are basic (sodium ethanoate ≈ pH 9); salts of weak bases are acidic (NH₄Cl, CuSO₄ ≈ pH 4.6).

  • M-54 — A higher pKa means a stronger acid.

    Confidence: verified

    Why it’s wrong: Inverted. pKa = −log₁₀Ka.

    Correction: Lower pKa = larger Ka = stronger acid.

  • M-55 — Neutralisation always gives pH 7 at the equivalence point.

    Confidence: verified

    Why it’s wrong: The equivalence point is where stoichiometrically equivalent amounts have reacted; the pH there depends on the salt formed.

    Correction: Weak acid + strong base gives an equivalence pH above 7; strong acid + weak base below 7. Choose an indicator whose range spans the equivalence pH, not one that changes at 7.

  • M-56 — The equivalence point and the end point are the same thing.

    Confidence: verified

    Why it’s wrong: Equivalence is the chemical stoichiometric point; the end point is where the indicator changes.

    Correction: A good indicator makes them nearly coincide.

  • M-57 — Diluting a buffer changes its pH.

    Confidence: verified

    Why it’s wrong: pH depends on pKa and the ratio of conjugate base to acid.

    Correction: Dilution changes both concentrations equally, so pH is essentially unchanged (buffer capacity falls, though).

  • M-58 — A buffer keeps pH constant no matter what you add.

    Confidence: verified

    Why it’s wrong: It resists change for small additions.

    Correction: Exceed the capacity of either component and pH moves sharply.

  • M-59 — Water is neutral so its pH is always 7.

    Confidence: verified

    Why it’s wrong: Water is always neutral ([H⁺] = [OH⁻]) but Kw increases with temperature.

    Correction: At 50 °C neutral water has pH below 7.

  • M-60 — Amphoteric and amphiprotic mean the same thing.

    Confidence: likely

    Why it’s wrong: Amphiprotic species can donate or accept a proton (HCO₃⁻, H₂O). Amphoteric is broader.

    Correction: Amphoteric can act as acid or base, which includes Lewis behaviour (Al₂O₃ is amphoteric but not amphiprotic).

    • Caveat (uncertain) How strictly IB distinguishes amphiprotic from amphoteric: Teacher should confirm how strictly IB distinguishes these.

What "HL standard" actually looks like

HLHL adds pOH and the Ka/Kb/pKa/pKb system (with Ka × Kb = Kw), salt hydrolysis, pH curves for all four strong/weak combinations (including buffer regions), indicator theory (end point ≈ pKa), quantitative buffer pH including the effect of dilution, and the choice of appropriate indicators. +9 HL hours — the heaviest HL block in Reactivity.

See the command terms reference →

Linking questions

Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.

  1. Official IB

    Why are log scales useful when discussing [H⁺] and IEs?

    Structure 1.3.6 → R3.1 Proton transfer reactions

  2. Official IB

    Polyatomic anions are conjugate bases of common acids. What is the relationship between their stability and the conjugate acid's dissociation constant, Ka?

    Structure 2.1.2 → R3.1 Proton transfer reactions

  3. Official IB

    How does the value of K for the dissociation of an acid convey information about its strength?

    Reactivity 2.3.3 → R3.1 Proton transfer reactions

  4. Official IB

    How does the equilibrium law help us to determine the pH of a weak acid, weak base or a buffer solution?

    Reactivity 2.3.6 → R3.1 Proton transfer reactions

  5. Official IB

    Why has the definition of acid evolved over time?

    Reactivity 3.1.1 → R3.4 Electron-pair sharing reactions

  6. Official IB

    What are the conjugate acids of the polyatomic anions listed in Structure 2.1?

    Reactivity 3.1.2 → S2.1 The ionic model

  7. Official IB

    What is the periodic trend in the acid–base properties of metal and non-metal oxides?

    Reactivity 3.1.3 → S3.1 The periodic table: Classification of elements

  8. Official IB

    Why does the release of oxides of nitrogen and sulfur into the atmosphere cause acid rain?

    Reactivity 3.1.3 → S3.1 The periodic table: Classification of elements

  9. Official IB

    Why does the extent of ionization of water increase as temperature increases?

    Reactivity 3.1.5 → R2.3 How far? The extent of chemical change

  10. Official IB

    How would you expect the equilibrium constants of strong and weak acids to compare?

    Reactivity 3.1.6 → R2.3 How far? The extent of chemical change

  11. Official IB

    Why does the acid strength of the hydrogen halides increase down group 17?

    Reactivity 3.1.6 → R1.1 Measuring enthalpy changes

  12. Official IB

    How can the salts formed in neutralization reactions be separated?

    Reactivity 3.1.7 → S1.1 Introduction to the particulate nature of matter

  13. Official IB

    Neutralization reactions are exothermic. How can this be explained in terms of bond enthalpies?

    Reactivity 3.1.7 → R1.1 Measuring enthalpy changes

  14. Official IB

    How could we classify the reaction that occurs when hydrogen gas is released from the reaction between an acid and a metal?

    Reactivity 3.1.7 → R3.2 Electron transfer reactions

  15. Official IB

    Why is the equivalence point sometimes referred to as the stoichiometric point?

    Reactivity 3.1.8 → S1.4 Counting particles by mass: The mole

  16. Official IB

    How can titration be used to calculate the concentration of an acid or base in solution?

    Reactivity 3.1.8 → S1.3 Electron configurations

  17. Official IB

    How can we simplify calculations when equilibrium constants Ka and Kb are very small?

    Reactivity 3.1.11 → R2.3 How far? The extent of chemical change

  18. Official IB

    What are some of the similarities and differences between indicators used in acid–base titrations and in redox titrations?

    Reactivity 3.1.14 → R3.2 Electron transfer reactions

  19. Official IB

    Why must buffer solutions be composed of weak acid or base conjugate systems, not of strong acids or bases?

    Reactivity 3.1.16 → R2.3 How far? The extent of chemical change

  20. Official IB

    How does Le Châtelier's principle enable us to interpret the behaviour of indicators and buffer solutions?

    Reactivity 3.1.17 → R2.3 How far? The extent of chemical change

  21. Official IB

    What is the relationship between Brønsted–Lowry acids and bases and Lewis acids and bases?

    Reactivity 3.4.6 → R3.1 Proton transfer reactions

  22. Official IB

    Nitration of benzene uses a mixture of concentrated nitric and sulfuric acids to generate a strong electrophile, NO₂⁺. How can the acid/base behaviour of HNO₃ in this mixture be described?

    Reactivity 3.4.13 → R3.1 Proton transfer reactions

4 further official linking questions target Tool, Inquiry or Nature of Science strands and are not drawn as edges on the map:

Editorial The decision to surface these off-graph questions here is our editorial judgement — the IB does not prescribe where they should appear.

  1. Official IB

    What is the shape of a sketch graph of pH against [H⁺]?

    Reactivity 3.1.4 → Tools 1

  2. Official IB

    When are digital sensors (e.g. pH probes) more suitable than analogue methods (e.g. pH paper/solution)?

    Reactivity 3.1.4 → Nature of Science

  3. Official IB

    What physical and chemical properties can be observed to distinguish between weak and strong acids or bases of the same concentration?

    Reactivity 3.1.6 → Tools 1

  4. Official IB

    When collecting data to generate a pH curve, when should smaller volumes of titrant be added between each measurement?

    Reactivity 3.1.13 → Tools 1

Bridge: GCSE → IB HL

Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.

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