Energetics II
Notes for every section, eight new exam-style questions with full worked solutions, and online recall cards for your review days. Nothing here is copied from the book.
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Read one section at a time, then answer its check question before you open the answer.
13.1 The key definitions
| Term | Definition | Sign |
|---|---|---|
| Enthalpy of atomisation | Enthalpy change when 1 mol of gaseous atoms forms from the element in its standard state | + |
| First ionisation energy | Enthalpy change to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of gaseous 1+ ions | + |
| Second ionisation energy | Removing 1 mol of electrons from 1 mol of gaseous 1+ ions | + |
| First electron affinity | Enthalpy change when 1 mol of gaseous atoms gains 1 mol of electrons to form 1 mol of gaseous 1− ions | − |
| Second electron affinity | 1 mol of gaseous 1− ions gains 1 mol of electrons | + (repulsion) |
| Lattice energy | Enthalpy change when 1 mol of a solid ionic compound forms from its gaseous ions | − |
| Enthalpy of hydration | Enthalpy change when 1 mol of gaseous ions dissolves in water | − |
| Enthalpy of solution | Enthalpy change when 1 mol of a solute dissolves in water | + or − |
The second electron affinity is endothermic because an electron is being added to a negative ion, which repels it.
Check yourself: Why is the second electron affinity of oxygen positive?
The electron is added to a negative ion, which repels it, so energy must be put in
13.2 Born–Haber cycles
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13.3 Cycles for ions with a charge of 2
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13.4 What controls the size of the lattice energy
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13.5 Experimental and theoretical lattice energies; polarisation
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13.6 Enthalpy changes of solution and hydration
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13.7 What controls the hydration enthalpy
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13.8 Measuring an enthalpy change of solution
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13.9 Why enthalpy is not the whole story
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13.10 Predicting the sign of ΔSsystem
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13.11 Four reactions to know
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13.12 Total entropy change
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13.13 Gibbs energy change, ΔG, and temperature
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13.14 ΔG and the equilibrium constant
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13.15 Feasible is not the same as fast
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Extra worked questions
Cover the solution, answer on paper, then mark yourself against each step. Each step shows the mark it earns.
All 8 worked solutions are free for members. The first one is open to everyone.
Three definitions
Define (a) the first electron affinity of chlorine, (b) the lattice energy of sodium chloride, and (c) the enthalpy of hydration of Na+.
Try it first, then show the worked solution
| (a) | Enthalpy change when 1 mol of gaseous Cl atoms gains 1 mol of electrons to form 1 mol of gaseous Cl− ions | 1 |
|---|---|---|
| (b) | Enthalpy change when 1 mol of solid NaCl forms from its gaseous ions | 1 |
| (c) | Enthalpy change when 1 mol of gaseous Na+ ions dissolves in water | 1 |
Answer: Each definition needs "1 mol" and the correct state symbols (gaseous).
Examiner tip: Write the equation alongside, e.g. Cl(g) + e− → Cl−(g). It shows the states clearly.
Lattice energy of sodium chloride
Calculate the lattice energy of NaCl using (kJ mol−1): ΔfH(NaCl) −411; ΔatH(Na) +107; IE1(Na) +496; ΔatH(Cl) +122; EA1(Cl) −349.
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Lattice energy of magnesium chloride
Calculate the lattice energy of MgCl2 using (kJ mol−1): ΔfH(MgCl2) −641; ΔatH(Mg) +148; IE1(Mg) +738; IE2(Mg) +1451; ΔatH(Cl) +122; EA1(Cl) −349.
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Enthalpy change of solution of NaCl
Use these values (kJ mol−1) to calculate the enthalpy change of solution of NaCl: lattice energy −787; ΔhydH(Na+) −406; ΔhydH(Cl−) −378.
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Covalent character in silver iodide
The experimental lattice energy of AgI is much more exothermic than the value calculated from the perfect ionic model. Explain why.
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Total entropy for making ammonia
N2(g) + 3H2(g) → 2NH3(g); ΔH = −92 kJ mol−1; ΔSsystem = −199 J K−1 mol−1. Calculate ΔStotal at 298 K and at 600 K, and comment on feasibility.
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From ΔG to K
For the formation of ammonia at 298 K, ΔG = −32.7 kJ mol−1. Calculate the equilibrium constant at 298 K (R = 8.31 J K−1 mol−1) and explain why no ammonia forms when nitrogen and hydrogen are mixed at room temperature.
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Decomposing calcium carbonate
CaCO3(s) → CaO(s) + CO2(g); ΔH = +178 kJ mol−1. Standard entropies (J K−1 mol−1): CaCO3 93, CaO 40, CO2 214. (a) Calculate ΔS. (b) Show the reaction is not feasible at 298 K. (c) Find the minimum temperature for it to be feasible. (d) Explain why a feasible reaction might still not be observed.
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Video walkthroughs
Narrated walkthroughs of exam-style questions from this topic of the book, showing where each mark comes from. Try the question first, then watch. Captions are available in the player.
Born-Haber cycle for potassium bromide
Total entropy change for rusting
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When does a decomposition become feasible?
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Hydration enthalpy from an energy cycle
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From delta G to the equilibrium constant
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Ethanol: a huge K, but no fire
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Recall cards
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