Mini-Lesson

What drives reactions — energetics

This mini-lesson covers Reactivity 1 at HL: enthalpy change, measuring it by calorimetry (q = mcΔT), Hess's law, bond enthalpies, and a first look at entropy and spontaneity. This HL lesson also builds in the Additional Higher Level (AHL) material.

Work through each screen, answer the questions as you go (some are wordy, some are calculations) and collect ⭐ stars. Press Start when you're ready.

Reactivity 1.1

Enthalpy change (ΔH)

Enthalpy change ΔH is the heat exchanged at constant pressure.

Exothermic: releases heat, ΔH negative, products lower in energy.
Endothermic: absorbs heat, ΔH positive, products higher in energy.

On an energy profile, the activation energy is the barrier that must be climbed before reactants become products.

Quick check

Exo or endo?

?A reaction releases heat to the surroundings. Its ΔH is:

Reactivity 1.1

Measuring ΔH by calorimetry

Burn or react a known amount and measure the temperature rise of water. The heat gained by the water is:

q = m c ΔTm = mass of water (g), c = 4.18 J g⁻¹ °C⁻¹, ΔT = temperature change

Then divide by the moles reacted to get ΔH per mole.

Calculate

Heat released

1Burning a fuel raises the temperature of 100 g of water by 20.0 °C (c = 4.18 J g⁻¹ °C⁻¹). Calculate q, in joules.

J

q = m c ΔT = 100 × 4.18 × 20.0.

Calculate

Enthalpy per mole

2That 8.36 kJ came from burning 0.010 mol of fuel. Calculate ΔH of combustion, in kJ mol⁻¹ (remember the sign).

kJ mol⁻¹

ΔH = −q ÷ n = −8.36 ÷ 0.010 (negative because combustion is exothermic).

Reactivity 1.2

Hess's law and energy cycles

Hess's law: the enthalpy change of a reaction is independent of the route taken — it depends only on the initial and final states. This lets us add enthalpies around a cycle to find a value we cannot measure directly, using standard enthalpies of formation or combustion.

Reactivity 1.2

Bond enthalpies

Breaking bonds absorbs energy; making bonds releases energy. For gas-phase reactions:

ΔH = Σ(bonds broken) − Σ(bonds formed)using average bond enthalpies

Calculate

Using bond enthalpies

3For H₂(g) + ½O₂(g) → H₂O(g), use H–H = 436, O=O = 498, O–H = 463 kJ mol⁻¹. Calculate ΔH.

kJ mol⁻¹

Broken = 436 + ½(498) = 685; formed = 2 × 463 = 926; ΔH = 685 − 926.

Quick check

Bond energy signs

?Which statement is correct?

AHL — Reactivity 1.4

Entropy and spontaneity

Entropy (S) measures the number of ways particles and energy can be arranged — its 'disorder'. Entropy tends to increase when solids melt or dissolve, and especially when gases form.

A change is spontaneous when it is thermodynamically favoured. Exothermic reactions and increases in entropy both favour spontaneity — a balance quantified at HL by Gibbs energy.

AHL check

Entropy change

?Which change involves the biggest INCREASE in entropy?

AHL — Reactivity 1.2 & 1.4

Energy cycles, entropy & Gibbs energy

AHL quantifies spontaneity. Born-Haber cycles use Hess's law to find lattice enthalpies. Entropy change ΔS = ΣS(products) − ΣS(reactants). Combining enthalpy and entropy gives the Gibbs energy:

ΔG = ΔH − TΔSa change is spontaneous when ΔG < 0

AHL check

Sign of ΔG

?A reaction is thermodynamically spontaneous when its Gibbs energy change ΔG is:

Sort it

Energy and disorder

Tap a change, then the box it fits.

🔥 Exothermic

❄️ Endothermic

🎲 Entropy increases

Match it

Match the term to its meaning

Tap an item on the left, then its match on the right.

Term

Meaning

Recap

The big ideas to know

ΔH: exothermic (−, releases heat) vs endothermic (+, absorbs heat)

Calorimetry: q = mcΔT, then ΔH = q ÷ n

Hess & bonds: ΔH is route-independent; ΔH = bonds broken − bonds formed

Entropy: disorder increases favour spontaneity (quantified by Gibbs energy at HL)

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