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What are the mechanisms of chemical change?

24 SL · +21 HL (sum of sub-topics)

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Guiding question

Official titleWhat are the mechanisms of chemical change?

The one big idea

Chemical change happens by three electron mechanisms: transfer (redox and acid–base), sharing (radicals) and pair-sharing (polar). All four sub-topics are electron bookkeeping — the same physics in different wrappers.

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.

Prerequisites are recorded at the sub-topic level only. The list below aggregates the prerequisites of this topic’s sub-topics.

What rests on this

Sub-topics in this topic

Core concepts that must be mastered

Reactivity 3 is the largest topic in the course — 24 SL hours ballooning to 45 at HL — because it covers every reaction mechanism. Four sub-topics partition reactions by what happens to electrons: proton transfer (3.1), electron transfer (3.2), electron sharing (3.3) and electron-pair sharing (3.4). The deep unification is that proton transfer and electron-pair sharing are the same chemistry viewed through different definitions — Brønsted–Lowry and Lewis — as the linking questions in §8 make explicit.

  • Proton transfer reactions (3.1.1–3.1.8, HL 3.1.9–3.1.17): A Brønsted–Lowry acid is a proton donor; a base is a proton acceptor (3.1.1). A proton in aqueous solution may be written as H⁺(aq) or H₃O⁺(aq). A conjugate acid–base pair differs by a single proton (3.1.2), and amphiprotic species — like water and HCO₃⁻ — can act as either acid or base (3.1.3). The logarithmic pH scale — pH = −log₁₀[H⁺], with [H⁺] = 10⁻ᵖᴴ — quantifies acidity (3.1.4). Kw = [H⁺][OH⁻] relates hydrogen and hydroxide ion concentrations; at 298 K, Kw = 1.0 × 10⁻¹⁴, and you can recognise acidic, neutral and basic solutions from the relative [H⁺] and [OH⁻] (3.1.5). Strong and weak acids/bases differ in extent of ionisation, not concentration — "strong" is not the same as "concentrated", and acid–base equilibria lie toward the weaker conjugate (3.1.6). Neutralisation reactions produce salts and water; you must formulate equations for acids with metal oxides, hydroxides, hydrogencarbonates and carbonates, and identify the parent acid and base of different salts (3.1.7). pH curves for strong acid–strong base titrations have a characteristic S-shape with a steep vertical section at the equivalence point (3.1.8); only monoprotic neutralisations are assessed. At HL, the pOH scale enters: pOH = −log₁₀[OH⁻], with full interconversion between [H⁺], [OH⁻], pH and pOH (3.1.9). Weak acid/base strengths are quantified by Ka, Kb, pKa, pKb, with Ka × Kb = Kw for a conjugate pair (3.1.10, 3.1.11) — quadratic equations are not expected. Salt hydrolysis determines the pH of a salt solution from the relative strengths of the parent acid and base (3.1.12). pH curves extend to all four strong/weak combinations, revealing buffer regions and points where pH = pKa (3.1.13). Indicators are weak acids whose conjugate pair members have different colours, with the end-point pH ≈ pKa; a suitable indicator has an end-point range coinciding with the equivalence-point pH (3.1.14, 3.1.15). Buffers resist pH change on addition of small amounts of acid or alkali, and their pH depends on pKa and the acid:conjugate-base concentration ratio (3.1.16, 3.1.17) — dilution of a buffer does not change its pH, because the ratio is unchanged.

  • Electron transfer reactions (3.2.1–3.2.11, HL 3.2.12–3.2.16): Oxidation is loss of electrons (OIL); reduction is gain (RIG), described by electron transfer, oxidation-state change, oxygen gain/loss or hydrogen loss/gain (3.2.1). Half-equations separate the two processes, balanced in acidic or neutral solution (3.2.2). Relative ease of oxidation or reduction is predictable from periodic position (3.2.3); the reactivity series need not be learned — data will be supplied. Acids react with reactive metals to release hydrogen (3.2.4). In any cell, oxidation is at the anode, reduction at the cathode (3.2.5). Primary (voltaic) cells convert energy from spontaneous redox reactions into electricity, with electrons flowing anode → cathode in the external circuit and ions moving through the salt bridge (3.2.6). Secondary (rechargeable) cells involve reversible redox reactions (3.2.7). Electrolytic cells drive non-spontaneous reactions using electrical energy; you deduce the products of electrolysis of a molten salt (3.2.8). Functional groups undergo oxidation (primary alcohols → aldehydes → carboxylic acids via a two-step process using distillation then reflux) and reduction (carboxylic acids → primary alcohols via the aldehyde; ketones → secondary alcohols; the role of hydride ions) (3.2.9, 3.2.10). Reduction of unsaturated compounds by H₂ addition lowers the degree of unsaturation (3.2.11). At HL, the standard hydrogen electrode (H⁺(aq) + e⁻ ⇌ ½H₂(g), E⦵ = 0 by convention) lets you measure standard electrode potentials (3.2.12). E⦵cell is calculated from standard electrode potentials — positive E⦵cell means spontaneous forward (3.2.13) — and ΔG⦵ = −nFE⦵cell connects to thermodynamics (3.2.14). You do not multiply E⦵ by the number of electrons when balancing half-equations, because E⦵ is an intensive property. Electrolysis of aqueous solutions involves competing reactions at each electrode, including oxidation and reduction of water, with effects of concentration and electrode nature (3.2.15). Electroplating coats an object with a thin metallic layer (3.2.16).

  • Electron sharing reactions(3.3.1–3.3.3)A radical is a species with an unpaired electron, making it highly reactive (3.3.1). Radicals form by homolytic fission — each fragment keeps one electron, shown with single-barbed "fish hook" arrows — under UV light or heat, producing the initiation step of a chain reaction (3.3.2). Radical substitution with alkanes proceeds through propagation and termination steps, giving a mixture of products (3.3.3). The stability of alkanes comes from strong, essentially non-polar C–C and C–H bonds. This sub-topic has no HL extension.

  • Electron-pair sharing reactions (3.4.1–3.4.5, HL 3.4.6–3.4.13): A nucleophile donates both bonding electrons; an electrophile accepts them (3.4.1, 3.4.4). Nucleophilic substitution forms a new bond as a leaving group departs (3.4.2). Heterolytic fission — both electrons stay with one fragment — is the electron bookkeeping behind these reactions, shown with curly arrows for electron-pair movement (3.4.3). Alkenes undergo electrophilic addition because the C=C bond has high electron density (3.4.5). At SL, mechanisms are not assessed for addition — only the products. At HL, Lewis acid–base theory unifies nucleophiles (Lewis bases) and electrophiles (Lewis acids) — the same chemistry as proton transfer (3.1), one level more general (3.4.6, 3.4.7). Coordination bonds form when a Lewis base donates an electron pair to a Lewis acid; complex ions form when ligands donate electron pairs to transition-element cations, and you can deduce the charge on a complex from its formula and ligands (3.4.8). SN1 vs SN2 mechanisms — the concerted one-step SN2 of primary halogenoalkanes versus the two-step SN1 of tertiary — differ in stereospecificity (SN2 inverts configuration) and in what affects the rate (3.4.9, 3.4.10). Markovnikov addition to unsymmetrical alkenes follows from carbocation stability: the H adds to the carbon with more hydrogens, producing the more stable carbocation (3.4.12). Electrophilic substitution of benzene with a charged electrophile E⁺ extends the mechanism toolkit to aromatic chemistry, where addition is disfavoured because it would break the delocalised π system (3.4.13). Formation of the electrophile itself is not assessed.

  • One mechanism, three namesThe linking questions in §8 reveal the deep unification: proton transfer (3.1) and electron-pair sharing (3.4) are the same chemistry viewed through different definitions — Brønsted–Lowry and Lewis. Redox (3.2) connects to electrochemistry and thermodynamics via ΔG⦵ = −nFE⦵cell. Every mechanism in Reactivity 3 traces back to electron movement, and every link traces back to the Spine Map.

Quantitative non-negotiables

  • Deduce oxidation states; identify oxidised/reduced species and oxidising/reducing agents.
  • Balance redox half-equations in acidic or neutral solution.
  • Calculate pH = −log₁₀[H⁺] and [H⁺] = 10⁻ᵖᴴ; use Kw = [H⁺][OH⁻].
  • Formulate neutralisation equations for acids with metal oxides, hydroxides, carbonates and hydrogencarbonates.
  • Interconvert [H⁺], [OH⁻], pH and pOH; solve Ka, Kb, pKa, pKb and buffer pH problems (HL).
  • Calculate E⦵cell from standard electrode potentials; use ΔG⦵ = −nFE⦵cell (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.

  • M-61 — The oxidising agent is the thing that gets oxidised.

    Confidence: verified

    Why it’s wrong: The oxidising agent oxidises something else and is therefore itself reduced.

    Correction: Use the 'travel agent' analogy: the agent arranges the trip but doesn't go.

  • M-62 — Oxidation state is the same as ionic charge.

    Confidence: verified

    Why it’s wrong: They coincide for monatomic ions but not otherwise.

    Correction: Oxidation state is a formal bookkeeping value assigned by electronegativity even in purely covalent species (S in SO₄²⁻ is +6, though there's no S⁶⁺ ion).

  • M-63 — Fluorine is the best oxidising agent, so fluoride is the best reducing agent.

    Confidence: verified

    Why it’s wrong: Inverted. Down group 17 the halogens become weaker oxidising agents while the halide ions become stronger reducing agents.

    Correction: F₂ is the strongest oxidiser; I⁻ is the strongest reducer of the halides.

  • M-64 — The anode is always positive.

    Confidence: verified

    Why it’s wrong: Oxidation always occurs at the anode, but the anode is negative in a voltaic cell and positive in an electrolytic cell.

    Correction: Anchor on the process (oxidation at the anode), not the sign.

  • M-65 — You multiply E⦵ by the number of electrons when balancing half-equations.

    Confidence: verified

    Why it’s wrong: E⦵ is intensive — it does not scale with stoichiometry.

    Correction: ΔG⦵ does scale; that's why ΔG⦵ = −nFE⦵ has the n in it.

  • M-66 — A positive E⦵cell means the reaction is fast.

    Confidence: verified

    Why it’s wrong: Positive E⦵cell means spontaneous (ΔG⦵ negative).

    Correction: It says nothing about rate.

  • M-67 — Electrons flow through the salt bridge.

    Confidence: verified

    Why it’s wrong: Electrons flow through the external wire.

    Correction: The salt bridge carries ions to maintain electrical neutrality.

  • M-68 — Electrolysis of aqueous NaCl gives sodium at the cathode.

    Confidence: verified

    Why it’s wrong: Water is reduced in preference (hydrogen is evolved), because reducing water is easier than reducing Na⁺.

    Correction: Reactivity 3.2.15 frames this as competing electrode reactions.

  • M-75 — Radical substitution of methane gives a clean single product.

    Confidence: verified

    Why it’s wrong: It gives a mixture (CH₃Cl, CH₂Cl₂, CHCl₃, CCl₄, plus ethane from termination).

    Correction: This is why it's synthetically poor.

  • M-69 — Curly arrows show where atoms move.

    Confidence: verified

    Why it’s wrong: Curly arrows show the movement of electron pairs (double-headed) or single electrons (fishhook, single-headed).

    Correction: The tail must sit on a lone pair or a bond, never on an atom or a positive charge.

  • M-70 — A nucleophile is attracted to the positive charge, so it attacks the nucleus.

    Confidence: verified

    Why it’s wrong: A nucleophile donates an electron pair to an electron-poor centre.

    Correction: It attacks a δ+ carbon or a carbocation. It is a Lewis base.

  • M-73 — Benzene reacts by addition like an alkene.

    Confidence: verified

    Why it’s wrong: Benzene undergoes electrophilic substitution, preserving the delocalised π system.

    Correction: Addition would destroy the resonance stabilisation.

  • M-74 — In HBr addition to propene, the H goes on whichever carbon.

    Confidence: verified

    Why it’s wrong: Markovnikov — the electrophile adds to give the more stable carbocation (tertiary > secondary > primary).

    Correction: H adds to the carbon with more hydrogens. The explanation is carbocation stability, not a rule to memorise.

What "HL standard" actually looks like

HL adds pOH and Ka/Kb interconversions (Ka × Kb = Kw), salt hydrolysis, pH curves for all four strong/weak combinations, indicator theory, quantitative buffer pH, standard electrode potentials E⦵ with E⦵cell and spontaneity, electrolysis of aqueous solutions with competing reactions, electroplating, Lewis acid–base theory, SN1 vs SN2 mechanisms, Markovnikov addition, and electrophilic substitution of benzene.

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.

Linking questions are recorded at the sub-topic level only. The list below aggregates the linking questions of this topic’s sub-topics.

  1. Official IB

    How can isotope tracers provide evidence for a reaction mechanism?

    Structure 1.2.2 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  2. Official IB

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

    Structure 1.3.6 → R3.1 Proton transfer reactions

    (via R3.1)

  3. Official IB

    Why is the formation of an ionic compound from its elements a redox reaction?

    Structure 2.1.2 → R3.2 Electron transfer reactions

    (via R3.2)

  4. 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

    (via R3.1)

  5. Official IB

    Why do Lewis acid–base reactions lead to the formation of coordination bonds?

    Structure 2.2.3 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  6. Official IB

    What are the structural features of benzene that favour it undergoing electrophilic substitution reactions?

    Structure 2.2.12 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  7. Official IB

    What trends in reactivity of metals can be predicted from the periodic table?

    Structure 2.3.1 → R3.2 Electron transfer reactions

    (via R3.2)

  8. Official IB

    How can oxidation states be used to analyse redox reactions?

    Structure 3.1.6 → R3.2 Electron transfer reactions

    (via R3.2)

  9. Official IB

    What is the nature of the reaction between transition element ions and ligands in forming complex ions?

    Structure 3.1.10 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  10. Official IB

    How can functional group reactivity be used to determine a reaction pathway between compounds, e.g. converting ethene into ethanoic acid?

    Structure 3.2.2 → R3.2 Electron transfer reactions

    (via R3.2)

  11. Official IB

    How does the strength of a carbon–halogen bond affect the rate of a nucleophilic substitution reaction?

    Reactivity 1.2.1 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  12. Official IB

    Which species are the oxidizing and reducing agents in a combustion reaction?

    Reactivity 1.3.1 → R3.2 Electron transfer reactions

    (via R3.2)

  13. Official IB

    What are some of the environmental, economic, ethical and social implications of burning fossil fuels?

    Reactivity 1.3.3 → R3.2 Electron transfer reactions

    (via R3.2)

  14. Official IB

    What are the main differences between a fuel cell and a primary (voltaic) cell?

    Reactivity 1.3.5 → R3.2 Electron transfer reactions

    (via R3.2)

  15. Official IB

    How can electrochemical data also be used to predict the spontaneity of a reaction?

    Reactivity 1.4.3 → R3.2 Electron transfer reactions

    (via R3.2)

  16. Official IB

    When is it useful to use half-equations?

    Reactivity 2.1.1 → R3.2 Electron transfer reactions

    (via R3.2)

  17. Official IB

    Which mechanism in the hydrolysis of halogenoalkanes involves an intermediate?

    Reactivity 2.2.6 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  18. Official IB

    What are the rate equations and units of k for the reactions of primary and tertiary halogenoalkanes with aqueous alkali?

    Reactivity 2.2.11 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  19. 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

    (via R3.1)

  20. 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

    (via R3.1)

  21. Official IB

    Why has the definition of acid evolved over time?

    Reactivity 3.1.1 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  22. 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

    (via R3.1)

  23. 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

    (via R3.1)

  24. 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

    (via R3.1)

  25. 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

    (via R3.1)

  26. 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

    (via R3.1)

  27. Official IB

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

    Reactivity 3.1.6 → R1.1 Measuring enthalpy changes

    (via R3.1)

  28. 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

    (via R3.1)

  29. 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

    (via R3.1)

  30. 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

    (via R3.2)

  31. 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

    (via R3.1)

  32. 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

    (via R3.1)

  33. 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

    (via R3.1)

  34. 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

    (via R3.2)

  35. 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

    (via R3.1)

  36. 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

    (via R3.1)

  37. Official IB

    What are the advantages and limitations of using oxidation states to track redox changes?

    Reactivity 3.2.1 → S3.1 The periodic table: Classification of elements

    (via R3.2)

  38. Official IB

    The surface oxidation of metals is often known as corrosion. What are some of the consequences of this process?

    Reactivity 3.2.1 → S2.3 The metallic model

    (via R3.2)

  39. Official IB

    Why does metal reactivity increase, and non-metal reactivity decrease, down the main groups of the periodic table?

    Reactivity 3.2.3 → S3.1 The periodic table: Classification of elements

    (via R3.2)

  40. Official IB

    Electrical energy can be derived from the combustion of fossil fuels or from electrochemical reactions. What are the similarities and differences in these reactions?

    Reactivity 3.2.6 → R1.3 Energy from fuels

    (via R3.2)

  41. Official IB

    Secondary cells rely on electrode reactions that are reversible. What are the common features of these reactions?

    Reactivity 3.2.7 → R2.3 How far? The extent of chemical change

    (via R3.2)

  42. Official IB

    Under what conditions can ionic compounds act as electrolytes?

    Reactivity 3.2.8 → S2.1 The ionic model

    (via R3.2)

  43. Official IB

    How does the nature of the functional group in a molecule affect its physical properties, such as boiling point?

    Reactivity 3.2.9 → S3.2 Functional groups: Classification of organic compounds

    (via R3.2)

  44. Official IB

    What is the difference between combustion and oxidation of an alcohol?

    Reactivity 3.2.9 → R1.3 Energy from fuels

    (via R3.2)

  45. Official IB

    Why is there a colour change when an alcohol is oxidized by a transition element compound?

    Reactivity 3.2.9 → S3.1 The periodic table: Classification of elements

    (via R3.2)

  46. Official IBwording unverified against the guide

    How can oxidation states be used to show that the following molecules are given in increasing order of oxidation: CH₄, CH₃OH, HCHO, HCOOH, CO₂?

    Reactivity 3.2.10 → S3.1 The periodic table: Classification of elements

    (via R3.2)

  47. Official IB

    Why are some reactions of alkenes classified as reduction reactions while others are classified as electrophilic addition reactions?

    Reactivity 3.2.11 → R3.4 Electron-pair sharing reactions

    (via R3.4)

  48. Official IB

    How can thermodynamic data also be used to predict the spontaneity of a reaction?

    Reactivity 3.2.14 → R1.4 Entropy and spontaneity

    (via R3.2)

  49. Official IB

    How is it possible for a radical to be an atom, a molecule, a cation or an anion? Consider examples of each type.

    Reactivity 3.3.1 → S2.1 The ionic model

    (via R3.3)

  50. Official IB

    Why do chlorofluorocarbons (CFCs) in the atmosphere break down to release chlorine radicals but typically not fluorine radicals?

    Reactivity 3.3.2 → R1.2 Energy cycles in reactions

    (via R3.3)

  51. Official IB

    What is the reverse process of homolytic fission?

    Reactivity 3.3.2 → S2.2 The covalent model

    (via R3.3)

  52. Official IB

    Chlorine radicals released from CFCs are able to break down ozone, O₃, but not oxygen, O₂, in the stratosphere. What does this suggest about the relative strengths of bonds in the two allotropes?

    Reactivity 3.3.2 → S2.2 The covalent model

    (via R3.3)

  53. Official IB

    Why are alkanes described as kinetically stable but thermodynamically unstable?

    Reactivity 3.3.3 → R2.2 How fast? The rate of chemical change

    (via R3.3)

  54. Official IB

    What is the difference between the bond-breaking that forms a radical and the bond-breaking that occurs in nucleophilic substitution reactions?

    Reactivity 3.4.3 → R3.3 Electron sharing reactions

    (via R3.3)

  55. Official IB

    Why is bromine water decolourized in the dark by alkenes but not by alkanes?

    Reactivity 3.4.5 → R3.3 Electron sharing reactions

    (via R3.3)

  56. Official IB

    Why are alkenes sometimes known as "starting molecules" in industry?

    Reactivity 3.4.5 → S2.4 From models to materials

    (via R3.4)

  57. 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

    (via R3.1)

  58. Official IB

    Do coordination bonds have any different properties from other covalent bonds?

    Reactivity 3.4.7 → S2.2 The covalent model

    (via R3.4)

  59. Official IB

    What differences would be expected between the energy profiles for SN1 and SN2 reactions?

    Reactivity 3.4.9 → R2.2 How fast? The rate of chemical change

    (via R3.4)

  60. Official IB

    What are the rate equations for these SN1 and SN2 reactions?

    Reactivity 3.4.9 → R2.2 How fast? The rate of chemical change

    (via R3.4)

  61. Official IB

    How useful are mechanistic models such as SN1 and SN2?

    Reactivity 3.4.9 → R2.2 How fast? The rate of chemical change

    (via R3.4)

  62. Official IB

    Why is the iodide ion a better leaving group than the chloride ion?

    Reactivity 3.4.10 → S3.1 The periodic table: Classification of elements

    (via R3.4)

  63. Official IB

    What are the features of benzene, C₆H₆, that make it not prone to undergo addition reactions, despite being highly unsaturated?

    Reactivity 3.4.13 → S2.2 The covalent model

    (via R3.4)

  64. 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

    (via R3.1)

7 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

    (via R3.1)

  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

    (via R3.1)

  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

    (via R3.1)

  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

    (via R3.1)

  5. Official IB

    Why are some redox titrations described as "self-indicating"?

    Reactivity 3.2.2 → Tools 1

    (via R3.2)

  6. Official IB

    What observations can be made when metals are mixed with aqueous metal ions, and solutions of halogens are mixed with aqueous halide ions?

    Reactivity 3.2.3 → Tools 1

    (via R3.2)

  7. Official IB

    How is an electrolytic cell used for electroplating?

    Reactivity 3.2.16 → Tools 1

    (via R3.2)

Bridge: GCSE → IB HL

Derived. These are the bridge items tagged to this 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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