Chemical energetics: thermochemistry
Study H1 Chemistry enthalpy changes and calorimetry, use Hess cycles and bond energies, and explain lattice-energy trends.
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Before you begin
Energetics is about the heat taken in or given out when bonds break and form. The lessons start with enthalpy changes and energy profiles, measure enthalpy changes by calorimetry, use Hess’s law and bond energies to calculate them, and finish with how ionic charge and size affect lattice energy.
Be comfortable with: the mole, solution concentration, ionic and covalent bonding, and rearranging simple equations.
Learning goals
- Explain that most reactions transfer energy, mainly as heat from breaking and forming bonds, and can be exothermic (∆H negative) or endothermic (∆H positive)
- Construct and interpret energy profile diagrams using the enthalpy change and the activation energy
- Explain and use standard enthalpy changes of formation, combustion and neutralisation
- Explain and use bond energy, with ∆H positive for bond breaking
- Explain and use lattice energy, with ∆H negative for gaseous ions forming a solid lattice
- Calculate enthalpy changes from experimental results, using heat change = mc∆T
- Explain how ionic charge and ionic radius affect the size of a lattice energy
- Use Hess' Law and given energy cycles to find enthalpy changes that cannot be measured directly, such as an enthalpy change of formation from enthalpy changes of combustion
- Use Hess' Law with average bond energies to calculate enthalpy changes
Syllabus statements covered
- explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)
- construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 7)
- explain and use the terms: — enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; neutralisation
- explain and use the terms: — bond energy (∆H positive, i.e. bond breaking) (see also Section 2)
- explain and use the terms: — lattice energy (∆H negative, i.e. gaseous ions to solid lattice)
- calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mc∆T
- explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
- apply Hess’ Law to carry out calculations involving given simple energy cycles and relevant energy terms (restricted to enthalpy changes of formation, combustion and neutralisation), with particular reference to: — determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion
- apply Hess’ Law to carry out calculations involving given simple energy cycles and relevant energy terms (restricted to enthalpy changes of formation, combustion and neutralisation), with particular reference to: — average bond energies [construction of energy cycles is not required]
Lessons
Work through them in order.
- Enthalpy changes and energy profilesConnect enthalpy signs, energy levels, activation energy and catalyst pathways.
- Calorimetry (q = mcΔT)Calculate heat and molar enthalpy for H1 Chemistry from temperature measurements, checking mass, sign and the amount reacting.
- Hess’ law from given simple cyclesTrace supplied formation and combustion cycles, match physical states and use route equations to find unknown enthalpy changes.
- Bond enthalpy calculationsEstimate H1 Chemistry reaction enthalpies from bonds broken and formed, and explain why mean bond enthalpies give approximate values.
- Lattice energy and ionic trendsExplain how ionic charge and radius change the size of lattice energy.
Practise and check
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Topic reference
Quick reference
| Question | First move | What to show |
|---|---|---|
| Exothermic or endothermic? | decide whether heat is given out or taken in by the reaction | exothermic: Δ H negative; endothermic: Δ H positive |
| Read an energy profile | compare the products with the reactants | Δ H is products minus reactants; Eₐ is the peak minus reactants |
| Enthalpy change from a temperature change | q = mcΔ T for the solution | divide by the moles that react; give the sign and kJ mol⁻¹ |
| Hess’s law with formation data | products minus reactants | multiply each value by its coefficient |
| Hess’s law with combustion data | reactants minus products | multiply each value by its coefficient |
| Bond energies | bonds broken minus bonds formed | an estimate, because bond energies are averages |
| Compare lattice energies | compare the charges, then the sizes of the ions | larger charges and smaller ions give a more negative lattice energy |
Core knowledge to remember
- Enthalpy change, Δ H: the heat energy change at constant pressure. Exothermic reactions have negative Δ H; endothermic reactions have positive Δ H.
- Standard conditions: 298 K and 100 kPa, with every substance in its standard state.
- Standard enthalpy change of formation: one mole of a compound is formed from its elements in their standard states.
- Standard enthalpy change of combustion: one mole of a substance is burned completely in oxygen.
- Standard enthalpy change of neutralisation: an acid and a base react to form one mole of water.
- Bond energy: the energy needed to break one mole of a covalent bond in the gas phase; always positive.
- Lattice energy: the enthalpy change when one mole of an ionic solid is formed from its separate gaseous ions; always negative.
- Hess’s law: the enthalpy change of a reaction is independent of the route taken.
Common mistakes
- Giving an exothermic reaction a positive Δ H because the solution gets hotter.
- In calorimetry, using the mass of the solid instead of the mass of the solution, or dividing by the wrong number of moles.
- Reversing an equation without reversing the sign of its Δ H, or forgetting to multiply by coefficients.
- Using products minus reactants with combustion data.
- Treating bond-energy answers as exact. Bond energies are averages.
- Calling a more negative lattice energy “greater”. Say “more negative” or “larger in magnitude”.