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Q: What does H2 Chemistry Notes: Topic 7 - Chemical Energetics cover? A: Integrate enthalpy changes, bond energies, Born-Haber cycles, entropy, and Gibbs free energy for Core Idea 3 (Chemical Energetics) in the 2026 H2 Chemistry syllabus.
Core Idea
CORE IDEA 3 - TRANSFORMATION, 7 CHEMICAL ENERGETICS: THERMOCHEMISTRY AND THERMODYNAMICS (GIBBS FREE ENERGY AND ENTROPY)
Energetics connects molecular behaviour to spontaneous change. Mastering enthalpy cycles, entropy arguments, and ΔG calculations is essential for structured questions across Papers 2 and 3. Use this note alongside worksheets stored at https://eclatinstitute.sg/blog/h2-chemistry-notes. For the full topic map and paper weightings, see our H2 Chemistry Syllabus 2026-27 overview.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-01-13. Core Idea 3 Topic 7 is assessed across Papers 1–3.
Quick revision box
What this topic tests: Enthalpy cycles, entropy, Gibbs free energy, and feasibility arguments.
Top mistakes to avoid: Sign errors in ΔH/ΔS/ΔG; weak Hess-cycle setup; qualitative feasibility without thermodynamic linkage.
20-minute sprint plan: 5 min sign convention recap; 10 min Hess/ΔG calculations; 5 min spontaneity explanation drill.
Enthalpy change when one mole of compound forms from elements in standard states.
HX2(g)+21OX2(g)HX2O(l)
Standard enthalpy of combustion ΔHc∘
Enthalpy change when one mole of substance burns completely in excess oxygen.
CHX4(g)+2OX2(g)COX2(g)+2HX2O(l)
Enthalpy of atomisation ΔHatom
Enthalpy required to form one mole of gaseous atoms from the element.
21ClX2(g)Cl(g)
Bond dissociation enthalpy
Enthalpy to break one mole of a specified gaseous bond.
HX2(g)2H(g)
State conditions: 298K, 1bar, substances in standard states. Use bond energies (and a few constants like c for water) from the SEAB Chemistry Data Booklet where relevant; ΔH and S∘ data tables are typically provided in the question when needed.
2 Hess' Law and Enthalpy Cycles
Hess' law uses the state function property-route independent.
2.1 Formation/Combustion Cycle
Draw arrows from elements to compound (formation) or compounds to oxides (combustion).
Combine given ΔH∘ values; reverse sign if reversing a reaction; multiply when scaling.
2.2 Bond Energy Approach
Approximate reaction enthalpy:
ΔH≈∑(Ebroken)−∑(Eformed)
Ensure all species are gaseous; adjust for physical states if necessary (include enthalpy of vaporisation or atomisation).
When using mean bond energies, take the data from the SEAB Chemistry Data Booklet unless the question provides specific values.
Methane: standard bond-counting reference for combustion enthalpy calculations.
A fast bond-energy check uses methane combustion: count four C−H and two O=O bonds broken, then four C=O and four O−H bonds formed.
3 Lattice Energy and Born-Haber Cycles
Born-Haber cycle for ionic compounds involves:
Atomisation of metal and non-metal.
First (and second) ionisation energies of metal.
Electron affinity of non-metal.
Lattice energy (formation of solid from gaseous ions).
Construct cycles carefully, labelling each step with values and sign convention (exothermic negative). Remember standard enthalpy of formation sits at bottom of cycle.
Use the standard molar entropy values provided in the question data; convert to kJ⋅mol−1⋅K−1 before combining with enthalpy terms.
4.2 Gibbs Free Energy
ΔG=ΔH−TΔS
ΔG<0: spontaneous under given conditions.
Temperature dependence: if both ΔH and ΔS are positive, spontaneity increases with temperature.
4.3 Equilibrium Connection
At the threshold where ΔG=0, T≈ΔSΔH. This approximation is useful for estimating the temperature where spontaneity flips (e.g., some decomposition reactions) when ΔH and ΔS are treated as roughly constant over the range.
5 Worked Example
Question:
Given the data below, calculate the lattice energy of NaCl.
ΔHf∘(NaCl)=−411kJmol−1
Enthalpy of atomisation of sodium: +108kJmol−1
Bond dissociation enthalpy of ClX2: +242kJmol−1
First ionisation energy of sodium: +496kJmol−1
Electron affinity of chlorine: −349kJmol−1
Solution:
Construct the cycle:
Na(s)Na(g): +108kJmol−1.
Na(g)NaX+(g)+eX−
21ClX2(g)Cl(g)
Cl(g)+eX−ClX−(g)
NaX+(g)+ClX−(g)NaCl(s)
Apply Hess' law:
ΔHf∘=108+496+121−349+U
Solve for U:
U=ΔHf∘−(108+496+121−349)=−787kJmol−1
State final answer: lattice energy of NaCl is −787kJmol−1.
6 Experimental Considerations
Calorimetry (Paper 4):
Use polystyrene cup to minimise heat loss.
Record initial and maximum temperatures; extrapolate to correction for heat loss if required.
Compute q=mcΔT with appropriate units c=4.18Jg−1K−1 for water (value from the SEAB Chemistry Data Booklet).
Error sources: heat exchange with environment, incomplete reaction, inaccurate mass/volume measurements. Suggest improvements (insulation, repeating and averaging, using lid).
In spirit-burner practicals, ethanol is a common fuel for comparing enthalpy change setups across repeat trials.
Ethanol: common liquid fuel reference in calorimetry planning and error analysis.
7 Common Misconceptions
Forgetting to convert ΔS to kJ⋅mol−1⋅K−1 when combining with ΔH (divide by 1000).
Reversing sign for electron affinity (exothermic values negative).
Using bond energies from data booklet without adjusting for stoichiometric coefficients.
Assuming ΔH<0 guarantees spontaneity; emphasise the TΔS term.
8 Quick Drills
Determine ΔH for 2CX2HX6(g)+7OX2(g)4COX2(g)+6HX2O(l) using bond energies (Data Booklet) and compare with a tabulated ΔHc∘ value when it is provided.
For CaCOX3(s)CaO(s)+COX2(g)
Sketch a Born-Haber cycle for MgClX2 including the second ionisation energy of magnesium and the second electron affinity of chlorine (endothermic). Annotate each energy change.
Common exam mistakes
Unit mismatch when calculating ΔG: ΔH is typically in kJ⋅mol−1 while ΔS is in J⋅mol−1⋅K−1; forgetting to divide ΔS by 1000 before substituting into ΔG=ΔH−TΔS is the most common numerical error in this topic.
Reversing the sign of electron affinity: Electron affinity is exothermic (negative) for most elements; treating it as positive in a Born-Haber cycle gives the wrong lattice energy and is penalised heavily.
Applying bond energies to non-gaseous species: Mean bond energy calculations assume all species are in the gaseous state; failing to account for enthalpy of vaporisation or atomisation when a liquid or solid is involved gives an inaccurate answer.
Assuming ΔH < 0 guarantees spontaneity: A reaction is only guaranteed spontaneous when ΔG<0, which depends on both ΔH and TΔS; ignoring the entropy term is a recurring conceptual error.
Mishandling stoichiometric coefficients with bond energies: Bond counts must match the balanced equation exactly; students often forget to multiply bond energies by the correct coefficient (e.g. for four O−H bonds in two HX2O molecules).
Wrong arrow direction in Born-Haber cycles: Each step must be drawn in the correct direction with the correct sign; a common slip is drawing atomisation or ionisation steps as exothermic, when they are endothermic.
Ignoring second ionisation energy for Group 2 compounds: MgClX2 Born-Haber cycles require both the first and second ionisation energies of magnesium; omitting one gives an incorrect lattice energy.
Frequently asked questions
How do I decide whether a reaction is spontaneous at a given temperature? Calculate ΔG=ΔH−TΔS at the stated temperature (using consistent units). If ΔG<0, the reaction is spontaneous under those conditions. If ΔG>0, it is non-spontaneous. At the temperature where ΔG=0, the system is at equilibrium.
What is the difference between bond dissociation enthalpy and mean bond energy? Bond dissociation enthalpy is the energy to break a specific bond in a specific molecule (e.g. the first C−H bond in methane). Mean bond energy is an average value across many compounds containing that bond type. The Data Booklet provides mean bond energies; use them for approximate calculations and acknowledge they introduce some error.
Why do Born-Haber cycles use gaseous ions rather than aqueous ions? Lattice energy is defined as the enthalpy change when one mole of ionic solid forms from its constituent gaseous ions at infinite separation. Using gaseous ions isolates the lattice formation step from solvation effects, allowing the pure electrostatic interaction to be calculated via Hess' law.
When is a reaction spontaneous despite being endothermic? When the entropy increase (positive ΔS) is large enough that TΔS>ΔH, making ΔG negative. This typically occurs at high temperatures. A classic example is the thermal decomposition of calcium carbonate: endothermic but spontaneous above a threshold temperature because the release of gaseous COX2 greatly increases entropy.
Struggling with Chemical Energetics? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
