R1.1Reactivity 1 — What drives chemical reactions?

Measuring enthalpy changes — IB Chemistry SL Revision Notes & Topical Questions

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Lesson summary

R1.1 Measuring enthalpy changes — IB CHEMISTRY SL notes: Distinguish heat from temperature, classify exothermic and endothermic reactions, interpret energy profiles, define standard enthalpy changes and use calorimetry.Temperature — a measure of the average kinetic energy of the …

Comparison
Intensive and extensive
System and surroundings
Open system
Closed system
Isolated system

Part 1 of 9Heat and temperature

Temperature

Temperature a measure of the average kinetic energy of the particles, in K or °C.

Heat

Heat energy transferred from a hotter to a colder body, and a measure of energy content that depends on the amount.

Comparison

Comparison a beaker and a pipette drop of boiling water have the same temperature, but the beaker has far more heat content.

Intensive and extensive

Intensive and extensive temperature is independent of the amount, whereas heat depends on it.

Interval

Interval a change of 1 K equals a change of 1 °C.

Part 2 of 9Systems and conservation of energy

Systems and conservation of energy

System and surroundings

System and surroundings the system is the reacting substances and the surroundings are everything else, such as the flask and air.

Open system

Open system exchanges matter and energy with the surroundings, e.g. an uncovered beaker, and is the most common.

Closed system

Closed system exchanges energy but not matter, e.g. a sealed flask.

Isolated system

Isolated system exchanges neither, which a calorimeter tries to approach.

Law of conservation of energy

Law of conservation of energy energy cannot be created or destroyed, only transferred.

Part 3 of 9Exothermic and endothermic reactions

Enthalpy change ΔH H(products) − H(reactants) at constant pressure, in kJ mol⁻¹.

Feature	Exothermic	Endothermic
Energy flow	Out of the system	Into the system
Surroundings	Temperature rises	Temperature falls
Sign of ΔH	Negative	Positive
Products	Lower in enthalpy, more stable	Higher in enthalpy, less stable
Origin	Bonds formed release more energy than bonds broken	Bonds broken need more energy than bonds formed
Examples	Combustion, neutralisation, respiration	Thermal decomposition of CaCO₃, photosynthesis, dissolving NH₄NO₃

Part 4 of 9Energy profiles

1Diagram

Diagram plots enthalpy against reaction progress for reactants, transition state and products.

2Transition state

Transition state the highest-energy point, with bonds partly broken and partly formed;
- it is unstable, exists fleetingly and cannot be isolated.

3Activation energy Ea

Activation energy Ea the minimum energy colliding particles need for a successful collision, measured from the reactants to the transition state.

4Enthalpy change

Enthalpy change measured between the reactant and product levels.

5Exothermic profile

Exothermic profile products lie below the reactants, and E_a is generally lower.

6Endothermic profile

Endothermic profile products lie above the reactants.

7Catalyst

Catalyst lowers the hump but does not change ΔH.

8Kinetic stability

Kinetic stability an exothermic reaction with high E_a may be too slow to observe, e.g. petrol and air need a spark.

Part 5 of 9Standard enthalpy changes

Standard conditions

100 kPa, solutions at 1 mol dm⁻³, substances in their standard states, usually quoted at 298 K.

Symbol

ΔH^⦵ shows standard conditions.

Enthalpy change	Definition	Sign
Reaction, ΔH_r	Reactants in the stoichiometric equation form products	Either
Formation, ΔH_f	One mole of a compound formed from its elements	Either
Combustion, ΔH_c	One mole of a substance burned completely in excess oxygen	Always negative
Neutralisation, ΔH_neut	One mole of water formed from an acid and an alkali, about −57 kJ mol⁻¹ for strong acid and base	Always negative

Scaling
- ΔH refers to the amounts in the equation as written, so doubling the equation doubles ΔH and reversing it changes the sign.

Part 6 of 9Calorimetry calculation

Heat transferred

Heat transferred q = mcΔT, with q in J, m in g, c in J g⁻¹ K⁻¹ and ΔT in K.

What the symbols mean: E, Q or q = energy transferred (J) · m = mass (kg; g in chemistry calorimetry) · c = specific heat capacity (J kg⁻¹ °C⁻¹; water 4.18 J g⁻¹ °C⁻¹) · ΔT or Δθ = temperature change (°C or K)

Water

Water c = 4.18 J g⁻¹ K⁻¹ .
- (IB data booklet)

Enthalpy change per mole

ΔH = −q/n, with q converted from J to kJ by dividing by 1000.

What the symbols mean: ΔH = enthalpy change (kJ mol⁻¹) · q = heat energy measured (J, divide by 1000 for kJ) · n = amount of the limiting reactant (mol); negative sign = exothermic

Sign

ΔH is negative if the temperature rose.

Temperature units

Temperature units no conversion is needed because ΔT in °C equals ΔT in K.

Measure the mass of water or solution.

Record the corrected temperature change.

Calculate q = mcΔT in joules.

Find the moles of the limiting reactant or fuel burned.

Calculate ΔH = −q/n in kJ mol⁻¹.

Compare with the literature value and evaluate errors.

Part 7 of 9Calorimetry of reactions in solution

Calorimetry of reactions in solution

Design

Design one reagent is in excess and the temperature change is measured over a few minutes in an insulated cup.

Assumptions

Assumptions the solution has the specific heat capacity (4.18 J g⁻¹ K⁻¹) and density (1.00 g cm⁻³) of water, the cup absorbs negligible heat, the reaction is complete and heat losses are negligible.

Mass

Mass m is the total solution mass.

Temperature correction graph

Temperature correction graph take readings before mixing, add the second reactant,
- then extrapolate the cooling line back to the time of mixing to find the true ΔT.

Endothermic reactions

Endothermic reactions the extrapolated line is the warming section.
Assumption — the rate of cooling is constant.

Part 8 of 9Calorimetry of combustion

Method

Method a spirit burner heats water in a copper can, and the mass of fuel burned is found from the mass loss.

Energy transfer

Energy transfer not all heat reaches the water, as some is lost to the surroundings and some is absorbed by the calorimeter.

Errors

Errors heat loss and incomplete combustion make the measured |ΔH| smaller than the literature value.

Error	Effect on \|ΔH\|	Improvement
Heat loss to surroundings	Too small	Insulate, add a lid, shield from draughts, extrapolate
Incomplete combustion	Too small	Supply more oxygen
Heat absorbed by the calorimeter	Too small	Include its heat capacity
Evaporation of fuel or water	Uncertain mass	Cap the burner and weigh promptly

Random error
- arises from thermometer readings.

Part 9 of 9More from the specification

1Calorimetry heat equation (q

Calorimetry heat equation (q mc Delta T)

Calculates thermal energy transferred in calorimetry experiments using mass (m), specific heat capacity (c), and maximum temperature change (Delta T).

2Enthalpy change definition (Delta H)

The energy difference between the enthalpy of products and reactants under constant pressure conditions (Delta H = H_products - H_reactants).

3Sources of calorimetry experimental error

Heat losses to the surrounding atmosphere and thermal absorption by the calorimeter container underestimate true enthalpy changes.

4Temperature extrapolation graphs

Cooling curves plotted against time allow correction for thermal loss by extrapolating back to the theoretical maximum temperature.

Key terms

Term	Meaning
Heat	Energy transferred between a system and surroundings because of a temperature difference.
Enthalpy change (ΔH)	The heat energy transferred at constant pressure during a reaction.
Exothermic	A reaction that releases heat; ΔH is negative.
Endothermic	A reaction that absorbs heat; ΔH is positive.
Activation energy	The minimum energy colliding particles need to react.
Specific heat capacity	The energy needed to raise the temperature of 1 g of a substance by 1 K.
Calorimetry	The measurement of heat changes from the temperature change of a known mass of water or solution.

Avoid these mistakes

Using temperature as a measure of heat — temperature measures average kinetic energy, not energy content.

Showing E_a from reactants to products — it is from reactants to the transition state.

Writing a positive ΔH for an exothermic reaction — it is negative.

Forgetting to convert J to kJ — ΔH is quoted in kJ mol⁻¹.

Last updated: 14 October 2026