IB DP Chemistry Reactivity 1.1: Measuring enthalpy changes – Revision Notes

These are short-form notes for Reactivity 1.1. If any point is unfamiliar, go back to the measuring enthalpy changes study guide for the full explanation and longer worked examples.

The notes follow the IB Diploma Programme Chemistry guide (first assessment 2025, published February 2023), whose sub-topic Reactivity 1.1 contains four understandings, numbered 1.1.1 to 1.1.4, all summarised below. The guide lists nothing extra for HL in this sub-topic, so SL and HL students revise the same content; it applies to the course examined from 2025 onwards. Afterwards, attempt the practice questions, tick off the unit on the DP Chemistry checklist, browse other units from the course hub, and try a free diagnostic to find gaps.

Key definitions

Term	Meaning
System	The chemicals that react (reactants changing to products)
Surroundings	Everything else: solvent water, cup or can, thermometer, air
Energy conservation	Whatever energy the system loses, the surroundings gain, and vice versa; the total does not change
Temperature, T	Indicates the mean kinetic energy of the particles; K; does not depend on the amount present
Heat, Q	Energy that flows because of a temperature difference; J or kJ; does depend on the amount present
Exothermic	System gives energy to surroundings; ΔH < 0
Endothermic	System takes energy from surroundings; ΔH > 0
Enthalpy change, ΔH	Heat transferred at constant pressure
Standard enthalpy change, ΔH⦵	ΔH under standard conditions (100 kPa; values usually for 298 K) with every substance in its standard state; kJ mol⁻¹
Activation energy, Ea	Height of the peak on an energy profile, measured up from the reactant level

Reactivity 1.1.1 Heat versus temperature

• A teaspoon and a kettle of boiling water share one temperature, but the kettle can give out far more heat as it cools.
• Temperature does not depend on quantity; heat does.
• ΔT has the same value in K and °C: a 6.0 °C rise is a 6.0 K rise.
• Convert absolute temperatures only: T/K = t/°C + 273.15.

Reactivity 1.1.2 Reading the thermometer

	Exothermic	Endothermic
Energy moves	Out of the system	Into the system
Surroundings	Get warmer	Get colder
ΔH	Negative	Positive
Examples	Combustion, neutralisation, respiration, metal + acid	Thermal decomposition, photosynthesis, dissolving some ammonium salts

Key idea: the thermometer reads the surroundings. Reading goes up → system lost energy → ΔH negative. Reading goes down → ΔH positive.

Reactivity 1.1.3 Stability and energy profiles

• Lower potential energy means more stable.
• Products more stable than reactants → energy given out → exothermic.
• Products less stable than reactants → energy taken in → endothermic.
• The very strong triple bond makes N₂ very stable, so N₂ + O₂ forming nitrogen oxide is endothermic, unlike most combustions.

Steps: drawing an energy profile

1. Horizontal axis: reaction coordinate. Vertical axis: potential energy.
2. Two labelled horizontal levels, reactants and products, at the correct relative heights.
3. One smooth peak (transition state) linking them.
4. ΔH arrow from the reactant level to the product level (down = exothermic, up = endothermic).
5. Ea arrow from the reactant level to the top of the peak.

To read a profile: ΔH = E(products) − E(reactants) and Ea = E(peak) − E(reactants).

Reactivity 1.1.4 Calorimetry toolkit

Equation	Units	Watch out
Q = mcΔT	Q / J; m / g; c / J g⁻¹ K⁻¹; ΔT / K	m belongs to the water or solution that changes temperature
ΔH = −Q/n	ΔH / kJ mol⁻¹; Q / kJ; n / mol	n is for the limiting reactant, or the fuel burned
n = cV	c / mol dm⁻³; V / dm³	divide cm³ by 1000 first
n = m/M	m / g; M / g mol⁻¹	use the mass actually burned or added

The guide says Q = mcΔT and the specific heat capacity of water are both in the data booklet; c(water) is 4.18 J g⁻¹ K⁻¹. For dilute solutions, take the density as 1.00 g cm⁻³ and c as that of water.

Steps: reaction in solution

1. Total mass of liquid = sum of all solution volumes, in g.
2. ΔT = extrapolated peak (or lowest point) − starting temperature.
3. Q = mcΔT; divide by 1000 for kJ.
4. Work out moles of each reactant and pick the limiting one.
5. ΔH = −Q/n, with a sign, kJ mol⁻¹, three significant figures.

Steps: burning a fuel or food

1. Mass burned = burner mass before − after.
2. n(fuel) = mass burned ÷ M.
3. Q = mass of water × 4.18 × ΔT.
4. ΔH = −Q(kJ) ÷ n. For a food, give energy per gram = Q ÷ mass burned.

Mini example. 0.0150 mol of limiting reactant raises 40.0 g of solution by 4.3 K.

Q  = 40.0 × 4.18 × 4.3 = 719 J = 0.719 kJ
ΔH = −0.719 / 0.0150   = −47.9 kJ mol⁻¹

Why experimental ΔH is usually too small

Measured temperature changes tend to fall short of theory, so the calculated ΔH is less negative than the data value. Causes:

• energy passing to the air and bench, especially from a flame;
• energy warming the can, cup or thermometer, which Q ignores;
• incomplete combustion, shown by soot;
• fuel evaporating from the wick, so the mass loss is larger than the mass burned;
• products not in standard states, such as steam instead of liquid water.

For a reaction in solution, plot a temperature–time graph with readings before and after mixing. Extend the steady post-peak cooling trend back to the mixing time to estimate the true maximum.

Must-know distinctions

• Heat vs temperature: energy on the move (J) vs a measure of mean particle kinetic energy (K).
• System vs surroundings: the reacting chemicals vs everything else, solvent included.
• Q vs ΔH: Q is heat gained by the water in one experiment (J); ΔH is per mole (kJ mol⁻¹) and carries the opposite sign.
• ΔH vs Ea: gap between reactant and product levels vs gap from reactants up to the peak.
• Limiting vs excess: the energy released depends on the reactant that runs out.

Quick self-test

1. Give the SI unit of heat and of temperature.
2. 65.0 g of water rises in temperature by 7.2 K. Calculate Q. (c = 4.18 J g⁻¹ K⁻¹)
3. 2.40 kJ of heat goes into 150 g of water. Calculate ΔT.
4. When 0.0250 mol of a solid dissolves, the solution gets colder and the size of Q is 1.34 kJ. Give ΔH with its sign.
5. A solution goes from 18.7 °C to 25.9 °C. State ΔT in kelvin and the starting temperature in kelvin.
6. 30.0 cm³ of acid is mixed with 30.0 cm³ of alkali. What value of m goes into Q = mcΔT?
7. Burning 0.00800 mol of a fuel raises the temperature of 120 g of water by 8.5 K. Calculate ΔH of combustion, assuming no heat loss.
8. A profile shows the products above the reactants. Exothermic or endothermic? Which side is more stable?
9. Which labels does the guide expect on the two axes of an energy profile?
10. Give two reasons why a spirit-burner experiment gives a less negative ΔH than the data value.

Answers

1. Heat: J (kJ is also fine). Temperature: K.
2. Q = 65.0 × 4.18 × 7.2 = 1956 J = 1960 J (1.96 kJ).
3. ΔT = 2400 ÷ (150 × 4.18) = 3.83 K.
4. Colder surroundings, so endothermic: ΔH = +1.34 ÷ 0.0250 = +53.6 kJ mol⁻¹.
5. ΔT = 7.2 K; start = 18.7 + 273.15 = 291.85 K.
6. 60.0 g (60.0 cm³ of solution at 1.00 g cm⁻³).
7. Q = 120 × 4.18 × 8.5 = 4264 J = 4.264 kJ; ΔH = −4.264 ÷ 0.00800 = −533 kJ mol⁻¹.
8. Endothermic; the reactants are more stable.
9. Horizontal: reaction coordinate; vertical: potential energy.
10. Any two: energy lost to the air; energy taken up by the can; incomplete combustion; evaporation of fuel from the wick.

Where marks are usually lost

• Taking m as the mass of solid or fuel rather than the water or solution that was heated or cooled.
• Adding 273 to a temperature difference; ΔT is the same in K and °C.
• Leaving Q in joules before dividing by moles, or quoting J mol⁻¹ when the number is in kJ mol⁻¹.
• Writing an exothermic ΔH as positive, or putting a minus sign on ΔH when the thermometer fell.
• Using moles of the excess reagent instead of the limiting one.
• Taking the highest recorded temperature when a graph is provided for extrapolation.
• Labelling the profile axes “time” and “energy” in place of reaction coordinate and potential energy.
• Starting the Ea arrow at the product level or at the axis instead of at the reactant level.
• Defining temperature as “the amount of heat in something”.

Official syllabus

International Baccalaureate Organization, Diploma Programme Chemistry guide, first assessment 2025 (published February 2023). Material on this page: Reactivity 1.1 (Measuring enthalpy changes).