R1.1
Measuring enthalpy changes
5 SL (SL 1.1.1-1.1.4)
Unverified
← View on the mapParent topic: What drives chemical reactions?
Guiding question
Guiding questionWhat can be deduced from the temperature change that accompanies chemical or physical change?
The one big idea
Every chemical reaction transfers energy between the system and its surroundings. The temperature change you measure lets you calculate the enthalpy change, because heat at constant pressure is enthalpy.
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.
- EditorialS1.4 Counting particles by mass: The moleΔH = −Q/n requires the mole.
- EditorialS1.1 Introduction to the particulate nature of matterHeat vs temperature and Q = mcΔT rest on the kinetic model.
What rests on this
- EditorialR1.2 Energy cycles in reactions
- EditorialR1.3 Energy from fuels
- EditorialR1.4 Entropy and spontaneity
- EditorialR2.2 How fast? The rate of chemical change
- EditorialR3.1 Proton transfer reactions
Core concepts that must be mastered
- Energy transfer and conservation(1.1.1)A chemical reaction transfers energy between the system (the reactants and products) and the surroundings (everything else, including the calorimeter and the solvent). Total energy is always conserved: if the system loses energy, the surroundings gain it, and vice versa. Heat is energy transferred due to a temperature difference; temperature is a measure of the average kinetic energy of the particles. They are not the same thing — 100 g of water at 80 °C has more heat than 10 g of water at 80 °C, even though both have the same temperature.
- Exothermic and endothermic reactions(1.1.2)In an exothermic reaction the system releases heat to the surroundings: the surroundings get warmer and ΔH is negative. In an endothermic reaction the system absorbs heat from the surroundings: the surroundings get colder and ΔH is positive. The direction of the temperature change of the surroundings is your experimental evidence: the thermometer reading goes up → the reaction released heat → exothermic → ΔH < 0. This sign confusion — "the temperature went up so the reaction absorbed energy" — is the single most common error here. The temperature of the surroundings rose because the system released heat.
- Energy profiles and relative stability(1.1.3)Whether a reaction is exo- or endothermic depends on the relative stability of reactants and products. More stable means lower potential energy. If the products are more stable than the reactants, the system falls to a lower energy and releases the difference as heat — exothermic. If the products are less stable, the system must absorb energy to get there — endothermic. You must sketch and interpret energy profiles: the x-axis is the reaction coordinate (progress from reactants to products), the y-axis is potential energy. An exothermic profile starts higher and ends lower; an endothermic profile starts lower and ends higher. The activation energy Ea is the hill between reactants and the transition state, and it is always positive regardless of whether the overall reaction is exo- or endothermic — a catalyst lowers this hill without changing the start and end points.
- Standard enthalpy change and calorimetry(1.1.4)The standard enthalpy change ΔH⦵ is the heat transferred at constant pressure under standard conditions — standard temperature (298 K) and standard pressure (100 kPa), with all substances in their standard states. You determine it from the temperature change of a known mass of a pure substance using Q = mcΔT, where m is the mass of the substance being heated (typically the solution, treated as water), c is its specific heat capacity (c(water) = 4.18 J g⁻¹ K⁻¹, in the data booklet), and ΔT is the temperature change. Then ΔH = −Q/n, where n is the amount of reactant in moles. The negative sign converts heat absorbed by the surroundings into enthalpy lost by the system: if the surroundings gained heat (Q > 0, temperature rose), the reaction released it (ΔH < 0, exothermic). Units of ΔH are kJ mol⁻¹. Watch the sign, watch the units (Q comes out in J if you use c in J g⁻¹ K⁻¹), and divide by 1000 before reporting ΔH in kJ mol⁻¹.
- How this sub-topic connectsMeasuring enthalpy is the entry point to thermodynamics. The calorimetry technique (Q = mcΔT) connects to Tool 1 (experimental techniques) and is the basis for investigating combustion of alcohols and foods. Energy profiles reappear in R2.2 (activation energy, catalysis, Maxwell–Boltzmann curves) and at HL in R2.2.7 (multistep energy profiles with intermediates and transition states). The concept of ΔH as a state function underpins Hess's law (R1.2) and the entropy–Gibbs framework (R1.4, wholly HL). A catalyst does not change ΔH — it lowers Ea — so the energy profile with and without a catalyst has the same start and end but a different hill. Bond enthalpies (R1.2) explain why reactions are exo- or endothermic: breaking bonds costs energy, forming bonds releases it.
Quantitative non-negotiables
- Distinguish heat from temperature and identify a reaction as exothermic or endothermic from the direction of energy transfer.
- Sketch and interpret energy profiles: x-axis = reaction coordinate, y-axis = potential energy.
- Apply Q = mcΔT to find the heat transferred, then ΔH = −Q/n to find the standard enthalpy change (units kJ mol⁻¹).
- Relate the sign of ΔH to the relative stability of reactants and products.
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.
What "HL standard" actually looks like
No HL extension. Reactivity 1.1 is entirely SL and shared — every student must be able to do calorimetry and read an energy profile. The HL depth comes in R1.2 (Hess's law, Born–Haber cycles) and R1.4 (entropy and Gibbs energy).
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.
- Official IB
What is the relationship between temperature and kinetic energy of particles?
- Official IB
Most combustion reactions are exothermic; how does the bonding in N₂ explain the fact that its combustion is endothermic?
- Official IB
Why does the acid strength of the hydrogen halides increase down group 17?
- Official IB
Neutralization reactions are exothermic. How can this be explained in terms of bond enthalpies?
3 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.
- Official IB
What observations would you expect to make during an endothermic and an exothermic reaction?
- Official IB
How can the enthalpy change for combustion reactions, such as for alcohols or food, be investigated experimentally?
- Official IB
Why do calorimetry experiments typically measure a smaller change in temperature than is expected from theoretical values?
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.