R1 What drives chemical reactions?

Reactivity 1: measuring and calculating enthalpy changes, energy from fuels, and entropy and Gibbs energy, with exactly what HL adds to each subtopic.

IB Chemistry
SL and HL

4 subtopics in What drives chemical reactions?

Each one has its own page once something is published for it.

R1.1Measuring enthalpy changes122 questions · notes · 61 quick fire
R1.2Energy cycles in reactions144 questions · notes · 65 quick fire
R1.3Energy from fuels106 questions · notes · 60 quick fire
R1.4Entropy and spontaneityHL only97 questions · notes · 56 quick fire

Key terms introduced here

Each is in the Chemistry glossary, with every other term in the subject.

Born–Haber cycle: An energy cycle for the formation of an ionic compound from its elements by way of gaseous atoms and gaseous ions, used to find a lattice enthalpy or another unknown term.
Entropy: A measure of the dispersal or distribution of matter and/or energy in a system, symbol $S$ , in $\text{J K}^{-1}$ (or $\text{J K}^{-1}\ \text{mol}^{-1}$ ). The more ways the particles can be arranged and their energy distributed, the higher the entropy.
standard enthalpy change of formation: The enthalpy change when one mole of a substance is formed from its elements in their standard states under standard conditions; zero for an element in its standard state.

An overview of R1 What drives chemical reactions?

5 min readLast reviewed 14 Oct 2026

Reactivity 1 is about the energy changes that accompany chemical reactions. It shows how enthalpy changes are measured and calculated, how the energy released by fuels is used and what it costs the environment, and, at HL, why some reactions happen on their own while others do not.

How the subtopics fit together

Flow diagram of Reactivity 1. R1.1 Measuring enthalpy changes leads to R1.2 Energy cycles (enthalpy change from experiment) and R1.3 Energy from fuels (enthalpy of combustion). — How the subtopics of Reactivity 1 build on each other. Green subtopics are the same at SL and HL; R1.2 (blue) has HL content; R1.4 (orange) is studied only at HL.

R1.1 Measuring enthalpy changes. Energy is transferred between a chemical system and its surroundings while the total is conserved. Exothermic reactions release energy and warm the surroundings; endothermic reactions absorb it. Energy profiles show the relative stability of reactants and products. Calorimetry gives the heat transferred, $Q = mc\Delta T$ , and from it the enthalpy change, $\Delta H = -Q/n$ .
R1.2 Energy cycles in reactions. Breaking bonds absorbs energy and forming bonds releases it, so average bond enthalpies give an estimate of $\Delta H$ . Hess's law states that the enthalpy change does not depend on the route, so the enthalpy change of a reaction that cannot be measured directly can be found from other reactions.
R1.3 Energy from fuels. Combustion of reactive metals, non-metals and organic compounds, and the incomplete combustion of hydrocarbons and alcohols, which produces carbon monoxide and carbon. Coal, crude oil and natural gas compared, including the carbon dioxide each adds to the atmosphere and its link to the greenhouse effect. Biofuels as renewable alternatives, and fuel cells, which convert the chemical energy of a fuel directly into electrical energy.
R1.4 Entropy and spontaneity (HL). Entropy as a measure of the dispersal of matter and energy; standard entropy changes; the Gibbs energy change $\Delta G^{\ominus} = \Delta H^{\ominus} - T\Delta S^{\ominus}$ and its sign as the test of spontaneity; and the link between Gibbs energy and the equilibrium constant.

Big ideas across the topic

Conservation of energy. Every calculation in the topic rests on the principle that energy is neither created nor destroyed. Calorimetry follows energy from the reaction into the water; Hess's law follows it around a cycle. The same principle explains why melting and boiling are endothermic and freezing and condensing exothermic (S1.1).

Energy and structure. Bond enthalpies and lattice enthalpies measure the strength of the bonds described in Structure 2, so the energy change of a reaction is the balance between the bonds broken and the bonds formed. The very strong triple bond in ${{\mathrm{N}{\vphantom{X}}_{\smash[t]{2}}}}$ explains why the combustion of nitrogen is endothermic, and the strength of a carbon–halogen bond affects how fast a halogenoalkane is substituted (R3.4).

Energy and rate are separate questions. A large negative $\Delta H$ shows that products are more stable than reactants, but says nothing about how fast they form. A fuel is useful because it has a high activation energy (R2.2) and does not ignite until heated; alkanes are thermodynamically unstable in air but kinetically stable (R3.3).

Fuels and redox. Combustion is the oxidation of a fuel by oxygen (R3.2), and a fuel cell carries out the same reaction as two separate half-reactions, one at each electrode. The absorption of infrared radiation by carbon dioxide and methane, which makes them greenhouse gases, is explained by bond vibrations (HL S3.2).

The direction of change (HL). The Gibbs energy change brings enthalpy and entropy together to predict whether a reaction is spontaneous. It connects to the equilibrium constant through $\Delta G^{\ominus} = -RT\ln K$ (R2.3) and to cell potentials through $\Delta G^{\ominus} = -nFE^{\ominus}_{\text{cell}}$ (R3.2).

Standard level and higher level

Reactivity 1 contains the only subtopic in the course that is studied entirely at HL, R1.4 Entropy and spontaneity. Of the other three, only R1.2 has HL content, in understandings that follow its SL ones. Reactivity 1 takes 12 teaching hours at SL and 22 at HL.

Subtopic	Studied at SL and HL	Added at HL
R1.1 Measuring enthalpy changes	All of it	Nothing
R1.2 Energy cycles in reactions	Bond breaking and forming; calculating $\Delta H$ from average bond enthalpies; Hess's law with given equations and enthalpy changes	Definitions and equations for the standard enthalpy change of formation and of combustion; calculating $\Delta H^{\ominus}$ from $\Delta H_{\text{f}}^{\ominus}$ or $\Delta H_{\text{c}}^{\ominus}$ data; Born–Haber cycles for compounds of univalent and divalent ions
R1.3 Energy from fuels	All of it	Nothing
R1.4 Entropy and spontaneity	Nothing	All of it: entropy and predicting the sign of $\Delta S$ ; calculating $\Delta S^{\ominus}$ from $S^{\ominus}$ values; $\Delta G^{\ominus} = \Delta H^{\ominus} - T\Delta S^{\ominus}$ ; the sign of $\Delta G$ and the temperature at which a reaction becomes spontaneous; $\Delta G = \Delta G^{\ominus} + RT\ln Q$ and $\Delta G^{\ominus} = -RT\ln K$

The boundary in R1.2 is easy to misplace. SL students apply Hess's law to cycles built from equations whose enthalpy changes are given, and they use bond enthalpies; they do not use tables of $\Delta H_{\text{f}}^{\ominus}$ or $\Delta H_{\text{c}}^{\ominus}$ values or the summation equations built on them. Lattice enthalpy is SL as a measure of ionic bond strength (S2.1), while calculating it from a Born–Haber cycle is HL. At SL, enthalpy changes of combustion are measured by calorimetry (R1.1) and compared for different fuels (R1.3). At SL the energy profile and the sign of $\Delta H$ describe a reaction; the HL ideas of entropy and Gibbs energy explain why endothermic changes, such as ice melting above 0 °C, can still happen spontaneously.

Flashcards for R1

The official decks for this part of the syllabus.

R1.1 Measuring enthalpy changes164 cards
R1.2 Energy cycles in reactions150 cards
R1.3 Energy from fuels199 cards
R1.4 Entropy and spontaneity161 cards

From the data booklet

The entries of the chemistry data booklet that R1 uses. Each opens in the explained booklet.

And 188 more in the data booklet.

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How the subtopics fit together

How the subtopics of Reactivity 1 build on each other. Green subtopics are the same at SL and HL; R1.2 (blue) has HL content; R1.4 (orange) is studied only at HL.

R1.1 Measuring enthalpy changes. Energy is transferred between a chemical system and its surroundings while the total is conserved. Exothermic reactions release energy and warm the surroundings; endothermic reactions absorb it. Energy profiles show the relative stability of reactants and products. Calorimetry gives the heat transferred, $Q = mc\Delta T$ , and from it the enthalpy change, $\Delta H = -Q/n$ .

R1.2 Energy cycles in reactions. Breaking bonds absorbs energy and forming bonds releases it, so average bond enthalpies give an estimate of $\Delta H$ . Hess's law states that the enthalpy change does not depend on the route, so the enthalpy change of a reaction that cannot be measured directly can be found from other reactions.

R1.3 Energy from fuels. Combustion of reactive metals, non-metals and organic compounds, and the incomplete combustion of hydrocarbons and alcohols, which produces carbon monoxide and carbon. Coal, crude oil and natural gas compared, including the carbon dioxide each adds to the atmosphere and its link to the greenhouse effect. Biofuels as renewable alternatives, and fuel cells, which convert the chemical energy of a fuel directly into electrical energy.

R1.4 Entropy and spontaneity (HL). Entropy as a measure of the dispersal of matter and energy; standard entropy changes; the Gibbs energy change $\Delta G^{\ominus} = \Delta H^{\ominus} - T\Delta S^{\ominus}$ and its sign as the test of spontaneity; and the link between Gibbs energy and the equilibrium constant.

Big ideas across the topic

{{\mathrm{N}{\vphantom{X}}_{\smash[t]{2}}}}

explains why the combustion of nitrogen is endothermic, and the strength of a carbon–halogen bond affects how fast a halogenoalkane is substituted (R3.4).

Energy and rate are separate questions. A large negative

\Delta H

shows that products are more stable than reactants, but says nothing about how fast they form. A fuel is useful because it has a high activation energy (R2.2) and does not ignite until heated; alkanes are thermodynamically unstable in air but kinetically stable (R3.3).

The direction of change (HL). The Gibbs energy change brings enthalpy and entropy together to predict whether a reaction is spontaneous. It connects to the equilibrium constant through

\Delta G^{\ominus} = -RT\ln K

(R2.3) and to cell potentials through

\Delta G^{\ominus} = -nFE^{\ominus}_{\text{cell}}

(R3.2).

Standard level and higher level

Subtopic	Studied at SL and HL	Added at HL
R1.1 Measuring enthalpy changes	All of it	Nothing
R1.2 Energy cycles in reactions	Bond breaking and forming; calculating $\Delta H$ from average bond enthalpies; Hess's law with given equations and enthalpy changes	Definitions and equations for the standard enthalpy change of formation and of combustion; calculating $\Delta H^{\ominus}$ from $\Delta H_{\text{f}}^{\ominus}$ or $\Delta H_{\text{c}}^{\ominus}$ data; Born–Haber cycles for compounds of univalent and divalent ions
R1.3 Energy from fuels	All of it	Nothing
R1.4 Entropy and spontaneity	Nothing	All of it: entropy and predicting the sign of $\Delta S$ ; calculating $\Delta S^{\ominus}$ from $S^{\ominus}$ values; $\Delta G^{\ominus} = \Delta H^{\ominus} - T\Delta S^{\ominus}$ ; the sign of $\Delta G$ and the temperature at which a reaction becomes spontaneous; $\Delta G = \Delta G^{\ominus} + RT\ln Q$ and $\Delta G^{\ominus} = -RT\ln K$

The boundary in R1.2 is easy to misplace. SL students apply Hess's law to cycles built from equations whose enthalpy changes are given, and they use bond enthalpies; they do not use tables of

\Delta H_{\text{f}}^{\ominus}

\Delta H_{\text{c}}^{\ominus}

values or the summation equations built on them. Lattice enthalpy is SL as a measure of ionic bond strength (S2.1), while calculating it from a Born–Haber cycle is HL. At SL, enthalpy changes of combustion are measured by calorimetry (R1.1) and compared for different fuels (R1.3). At SL the energy profile and the sign of

\Delta H

describe a reaction; the HL ideas of entropy and Gibbs energy explain why endothermic changes, such as ice melting above 0 °C, can still happen spontaneously.