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What drives chemical reactions?

12 SL · +10 HL (sum of sub-topics)

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Guiding question

Official titleWhat drives chemical reactions?

The one big idea

Chemical change is driven by energy. Enthalpy tracks heat, entropy tracks dispersal, and Gibbs energy unites them: ΔG⦵ = ΔH⦵ − TΔS⦵. A reaction is spontaneous when ΔG < 0.

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 1 answers a single question — what drives chemical reactions? — in four sub-topics that build from the laboratory bench to the conceptual capstone of the course. Enthalpy (1.1, 1.2) measures heat. Fuels (1.3) ground the energy conversation in real-world chemistry. Entropy and Gibbs energy (1.4, HL only) measure dispersal and spontaneity, and the final equation ΔG⦵ = −RT lnK links thermodynamics to equilibrium in Reactivity 2.3.

  • Measuring enthalpy changes(1.1.1–1.1.4)Reactions transfer energy between system and surroundings, and total energy is conserved (1.1.1). Heat and temperature are not the same: heat is energy in transit, temperature measures average kinetic energy. Endothermic reactions absorb energy from the surroundings (temperature falls); exothermic reactions release it (temperature rises) (1.1.2). The relative stability of reactants and products determines which direction applies, visualised on an energy profile with reaction coordinate on the x-axis and potential energy on the y-axis (1.1.3). The standard enthalpy change ΔH⦵ is the heat transferred at constant pressure under standard conditions (298 K, 100 kPa, all species in their standard states, denoted by the plimsoll symbol ⦵), calculated from a measured temperature change using Q = mcΔT and ΔH = −Q/n (1.1.4). The minus sign is the trap: a temperature rise in the water (surroundings) means heat left the system, so ΔH is negative. Units are kJ mol⁻¹. Use the mass of the solution, not the solid reactant, and divide by the moles of the limiting reactant.
  • Energy cycles in reactions(1.2.1–1.2.5)Bond-breaking absorbs energy and bond-forming releases it (1.2.1). Average bond enthalpies let you estimate ΔH by summing bonds broken (endothermic) minus bonds formed (exothermic), though they are averages and may differ from measured values because the same bond type has different energies in different molecules. Hess's law — ΔH is independent of pathway — is the deeper principle: if you can reach a product by two routes, the enthalpy change is the same either way (1.2.2). This is a direct consequence of enthalpy being a state function. At HL, ΔHc⦵ (standard enthalpy of combustion) and ΔHf⦵ (standard enthalpy of formation) data are used in thermodynamic calculations (1.2.3), with the two summation equations ΔH⦵ = ΣΔHf⦵(products) − ΣΔHf⦵(reactants) and ΔH⦵ = ΣΔHc⦵(reactants) − ΣΔHc⦵(products) (1.2.4). The Born–Haber cycle applies Hess's law to ionic compounds, weaving together ionisation energies, atomisation enthalpies (sublimation and/or bond enthalpies), electron affinities, lattice enthalpy and enthalpy of formation (1.2.5). You interpret and determine values from a Born–Haber cycle for univalent and divalent ions, though construction of a complete cycle is not assessed. Remember that ΔHf⦵ of an element in its standard state is zero by definition.
  • Energy from fuels(1.3.1–1.3.5)Combustion of reactive metals, non-metals and organic compounds in oxygen releases energy (1.3.1). You must be able to write balanced combustion equations for hydrocarbons and alcohols. Incomplete combustion produces CO and carbon (C), not just CO₂ — this is why poorly ventilated gas heaters are dangerous (1.3.2). Fossil fuels — coal, crude oil, natural gas — differ in energy per unit mass and CO₂ output, and the link between CO₂ levels and the greenhouse effect is part of this sub-topic (1.3.3). Biofuels fix carbon over a short period via photosynthesis, making them renewable in principle; you should know the reactants and products of photosynthesis (1.3.4). Fuel cells convert chemical energy directly to electrical energy without combustion; you deduce half-equations for hydrogen and methanol fuel cells (1.3.5). Proton exchange membrane details are not assessed.
  • Entropy and spontaneity (HL only)(1.4.1–1.4.4)The capstone sub-topic is entirely HL — 5 hours with no SL content. Entropy S measures the dispersal of matter and energy: more ways of distributing energy means higher entropy, and S(gas) > S(liquid) > S(solid) under the same conditions (1.4.1). You can predict whether a physical or chemical change increases or decreases entropy, and calculate ΔS⦵ from standard entropy values S⦵ in the data booklet. Gibbs energy unites enthalpy and entropy: ΔG⦵ = ΔH⦵ − TΔS⦵ (1.4.2) — watch the units carefully, because ΔH is in kJ mol⁻¹ and ΔS is in J K⁻¹ mol⁻¹, so you must convert one before combining. At constant pressure a change is spontaneous when ΔG < 0 (1.4.3), and you can determine the temperature at which a reaction becomes spontaneous by setting ΔG = 0 and solving for T. As a reaction approaches equilibrium, ΔG becomes less negative and finally reaches zero (1.4.4); the equations ΔG = ΔG⦵ + RT lnQ and ΔG⦵ = −RT lnK connect thermodynamics to equilibrium (Reactivity 2.3).
  • Why this topic is the root of reactivityReactivity 1.4 is the conceptual capstone of the entire course: ΔG unifies enthalpy (1.1, 1.2), equilibrium (R2.3) and electrochemistry (R3.2) in a single equation. It is taught last in dependency order because you need enthalpy and equilibrium before it makes sense. The linking questions in §8 connect enthalpy to bonding (Structure 2), entropy to equilibrium (Reactivity 2.3), and Gibbs energy to electrochemistry (Reactivity 3.2) — these are the edges that turn the syllabus from a list into a network.

Quantitative non-negotiables

  • Apply Q = mcΔT and ΔH = −Q/n for calorimetric enthalpy changes.
  • Calculate ΔH from average bond enthalpies.
  • Apply Hess's law to multistep and multi-pathway reactions.
  • Use ΔH⦵ = ΣΔHf⦵(products) − ΣΔHf⦵(reactants) and ΔH⦵ = ΣΔHc⦵(reactants) − ΣΔHc⦵(products) (HL).
  • Calculate ΔS⦵ from standard entropy values, and apply ΔG⦵ = ΔH⦵ − TΔS⦵ (HL).
  • Use ΔG⦵ = −RT lnK and ΔG = ΔG⦵ + RT lnQ (HL).

Common failure modes

  • M-29 — The temperature went up so the reaction absorbed energy.

    Confidence: verified

    Why it’s wrong: Sign confusion via system/surroundings.

    Correction: Temperature of the surroundings rose because the system released heat; ΔH is negative; exothermic.

  • M-30 — Bond breaking releases energy.

    Confidence: verified

    Why it’s wrong: Breaking bonds always requires energy input (endothermic); forming bonds always releases it.

    Correction: Exothermic overall means more energy released in forming than absorbed in breaking.

  • M-32 — A catalyst changes ΔH.

    Confidence: verified

    Why it’s wrong: A catalyst lowers Ea by providing an alternative pathway; ΔH (a state function) is unchanged.

    Correction: A catalyst does not change ΔH or the equilibrium position.

  • M-31 — Bond enthalpy calculations give the exact enthalpy change.

    Confidence: verified

    Why it’s wrong: Tabulated values are averages across many compounds and apply to gaseous species.

    Correction: Results are approximate and diverge when liquids or solids are involved.

  • M-33 — Exothermic reactions are the ones that happen.

    Confidence: verified

    Why it’s wrong: Spontaneity is governed by ΔG = ΔH − TΔS.

    Correction: Endothermic reactions with a large positive ΔS (e.g. dissolving NH₄NO₃, thermal decomposition of CaCO₃ at high T) are spontaneous.

  • M-34 — Spontaneous means fast.

    Confidence: verified

    Why it’s wrong: Spontaneous is a thermodynamic statement about direction, not rate.

    Correction: Diamond → graphite has ΔG < 0 and takes geological time.

  • M-35 — Entropy is disorder / messiness.

    Confidence: likely

    Why it’s wrong: The tidy-bedroom analogy misleads.

    Correction: Entropy is a measure of the dispersal of matter and energy — the number of ways the system's energy can be distributed. Use the state-change ranking (gas ≫ liquid > solid) and the change in number of gaseous moles, not 'mess'.

  • M-36 — ΔS and ΔH can be added directly.

    Confidence: verified

    Why it’s wrong: ΔH is in kJ mol⁻¹, S⦵ in J K⁻¹ mol⁻¹ — different units.

    Correction: Divide ΔS by 1000 before using ΔG = ΔH − TΔS, and use T in kelvin.

  • M-37 — ΔG⦵ = 0 means nothing is happening.

    Confidence: verified

    Why it’s wrong: ΔG = 0 is the condition for equilibrium.

    Correction: Both directions proceed at equal rate; the system is dynamic, not static.

What "HL standard" actually looks like

HL adds Born–Haber cycles for univalent and divalent ions, thermodynamic calculations using ΔHf⦵ and ΔHc⦵ data, and the entire entropy–spontaneity framework: ΔS⦵, ΔG⦵ = ΔH⦵ − TΔS⦵, the temperature at which a reaction becomes spontaneous, and the relationship ΔG⦵ = −RT lnK linking thermodynamics to equilibrium.

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

    What properties of a greenhouse gas determine its "global warming potential"?

    Structure 3.2.9 → R1.3 Energy from fuels

    (via R1.3)

  2. Official IB

    What is the relationship between temperature and kinetic energy of particles?

    Reactivity 1.1.1 → S1.1 Introduction to the particulate nature of matter

    (via R1.1)

  3. Official IB

    Most combustion reactions are exothermic; how does the bonding in N₂ explain the fact that its combustion is endothermic?

    Reactivity 1.1.3 → S2.2 The covalent model

    (via R1.1)

  4. Official IB

    How would you expect bond enthalpy data to relate to bond length and polarity?

    Reactivity 1.2.1 → S2.2 The covalent model

    (via R1.2)

  5. 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 R1.2)

  6. Official IB

    Would you expect allotropes of an element, such as diamond and graphite, to have different ΔH⦵f values?

    Reactivity 1.2.3 → S2.2 The covalent model

    (via R1.2)

  7. Official IB

    What are the factors that influence the strength of lattice enthalpy in an ionic compound?

    Reactivity 1.2.5 → S2.1 The ionic model

    (via R1.2)

  8. Official IB

    Why is high activation energy often considered to be a useful property of a fuel?

    Reactivity 1.3.1 → R2.2 How fast? The rate of chemical change

    (via R1.3)

  9. Official IB

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

    Reactivity 1.3.1 → R3.2 Electron transfer reactions

    (via R1.3)

  10. Official IB

    How does limiting the supply of oxygen in combustion affect the products and increase health risks?

    Reactivity 1.3.2 → R2.1 How much? The amount of chemical change

    (via R1.3)

  11. Official IB

    Why do larger hydrocarbons have a greater tendency to undergo incomplete combustion?

    Reactivity 1.3.3 → S3.2 Functional groups: Classification of organic compounds

    (via R1.3)

  12. Official IB

    Why is carbon dioxide described as a greenhouse gas?

    Reactivity 1.3.3 → S3.2 Functional groups: Classification of organic compounds

    (via R1.3)

  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 R1.3)

  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 R1.3)

  15. Official IB

    Why is the entropy of a perfect crystal at 0 K predicted to be zero?

    Reactivity 1.4.1 → S1.1 Introduction to the particulate nature of matter

    (via R1.4)

  16. 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 R1.4)

  17. Official IB

    What is the likely composition of an equilibrium mixture when ΔG⦵ is positive?

    Reactivity 1.4.4 → R2.3 How far? The extent of chemical change

    (via R1.4)

  18. Official IB

    How can Gibbs energy be used to explain which of the forward or backward reaction is favoured before reaching equilibrium?

    Reactivity 2.3.7 → R1.4 Entropy and spontaneity

    (via R1.4)

  19. 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 R1.1)

  20. 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 R1.1)

  21. 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 R1.3)

  22. Official IB

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

    Reactivity 3.2.9 → R1.3 Energy from fuels

    (via R1.3)

  23. 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 R1.4)

  24. 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 R1.2)

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 observations would you expect to make during an endothermic and an exothermic reaction?

    Reactivity 1.1.2 → Tools 1

    (via R1.1)

  2. Official IB

    How can the enthalpy change for combustion reactions, such as for alcohols or food, be investigated experimentally?

    Reactivity 1.1.4 → Tools 1

    (via R1.1)

  3. Official IB

    Why do calorimetry experiments typically measure a smaller change in temperature than is expected from theoretical values?

    Reactivity 1.1.4 → Tools 1

    (via R1.1)

  4. Official IB

    What might be observed when a fuel such as methane is burned in a limited supply of oxygen?

    Reactivity 1.3.2 → Inquiry

    (via R1.3)

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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