Topic 1 of 3
Meaning of energy
Define a process and read its profile.
A-Level 8873, revised syllabus (2026-2027)
Breaking needs energy; forming releases it
A reaction can release energy overall and still need an initial barrier to be crossed.
Chemical reactions rearrange bonds. Breaking bonds absorbs energy; forming bonds releases energy. The enthalpy change is the net difference. At constant pressure, an exothermic reaction transfers heat to the surroundings and has negative ΔH; an endothermic reaction absorbs heat from them and has positive ΔH.
Read the barrier and the overall drop
Reactants are at zero, the barrier is 120 kJ mol-1 above them, and products are 80 kJ mol-1 below them. Forward activation energy is 120, reaction enthalpy is -80, and reverse activation energy is 200.
Read ΔH = H(products) - H(reactants). Read the forward activation energy from the reactant level to the barrier, not from the page baseline. In the figure, the reverse reaction starts at -80, so its barrier is 120 - (-80) = 200 kJ mol-1. A catalyst changes the pathway and barrier, but not the reactant/product levels or ΔH.
Worked example
Construct an endothermic profile
Sketch a reaction with Delta H = +30 kJ mol-1 and forward activation energy 80 kJ mol-1.
- Label the vertical axis enthalpy and the horizontal axis reaction progress. Choose the reactants as a zero reference.
- Place products 30 kJ mol-1 above the reactants and the barrier 80 kJ mol-1 above them.
- Join the levels with a smooth rise to the barrier and fall to the products. Label Delta H from reactants to products and Ea from reactants to the barrier.
The product level is higher than the reactant level. The reverse activation energy is 80 - 30 = 50 kJ mol-1.
Name the exact process and amount
An enthalpy value belongs to a stated equation, physical states and conditions.
A standard enthalpy change refers to species in their standard states at a stated temperature, commonly 298 K, with standard pressure 100 kPa. Aqueous standard data use a specified standard concentration, conventionally 1 mol dm-3 at this level. Standard does not by itself mean a reaction occurs at 0 degrees C. Always retain the states and temperature supplied with the data.
| Term | One-mole reference | Illustration |
|---|---|---|
| Enthalpy change of reaction | The reaction as its equation is written | Doubling the equation doubles its enthalpy change. |
| Standard enthalpy of formation | One mole of compound from its elements in their standard states | C(graphite,s) + O2(g) → CO2(g) |
| Standard enthalpy of combustion | One mole of substance completely burned in oxygen under standard conditions | CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) |
| Enthalpy of neutralisation | One mole of water formed when an acid reacts with a base | H+(aq) + OH-(aq) → H2O(l) |
| Bond energy | One mole of specified gaseous covalent bonds broken | H2(g) → 2H(g); positive |
| Lattice energy | One mole of ionic solid formed from separated gaseous ions | Na+(g) + Cl-(g) → NaCl(s); negative |
The formation enthalpy of an element in its standard state is zero by convention. This is not a claim that the element contains no energy. For standard formation of water, H2(g) + ½O2(g) → H2O(l) forms exactly one mole; fractional coefficients are therefore useful.
Lattice energy becomes greater in magnitude when ionic charges increase or ion radii decrease: the opposite charges attract more strongly at shorter separation. MgO, with Mg2+/O2-, has a much more negative lattice energy than NaCl, with Na+/Cl-. For equal charges, a smaller ion generally makes lattice formation more exothermic.