Topic 13: Energetics II — Edexcel A-Level Chemistry

Topic 13: Energetics II builds on the foundational ideas of enthalpy changes from Year 1, diving deeper into the thermodynamic principles that govern chemical reactions. You will explore the Born-Haber cycle for ionic compounds, calculating lattice enthalpy from experimental data, and understand how enthalpy changes relate to ion size and charge. This topic also introduces entropy and free energy, explaining why some reactions are spontaneous despite being endothermic. Mastering these concepts is essential for predicting reaction feasibility and understanding the driving forces behind chemical processes.

The Born-Haber cycle is a cornerstone of this topic, allowing you to calculate lattice enthalpy indirectly using Hess's law. You'll learn to construct cycles for simple ionic compounds like NaCl and MgO, and interpret how lattice enthalpy varies with ionic radius and charge. Additionally, you'll delve into entropy (S) as a measure of disorder, calculating entropy changes for reactions and using the Gibbs free energy equation (ΔG = ΔH - TΔS) to determine spontaneity. These ideas are crucial for A-Level Chemistry, as they connect enthalpy, entropy, and temperature in a unified framework.

Energetics II is not just about calculations; it provides a deeper understanding of why reactions occur. For example, the dissolution of ammonium nitrate in water is endothermic but spontaneous due to an increase in entropy. This topic also has real-world applications, such as in designing industrial processes where temperature control can shift reaction feasibility. By the end, you should be able to predict whether a reaction is spontaneous at a given temperature and explain the factors that influence lattice enthalpy and entropy changes.

• →Born-Haber cycle: A thermochemical cycle used to calculate lattice enthalpy by summing enthalpy changes for steps like atomisation, ionisation, electron affinity, and formation.
• →Lattice enthalpy: The enthalpy change when one mole of an ionic compound forms from its gaseous ions; it is exothermic and depends on ion charge and size.
• →Entropy (S): A measure of the disorder or randomness of a system; entropy increases with temperature, phase changes (solid to liquid to gas), and increased number of particles.
• →Gibbs free energy (ΔG): Determines spontaneity; ΔG = ΔH - TΔS. A reaction is spontaneous when ΔG < 0, and at equilibrium when ΔG = 0.
• →Standard entropy change (ΔS°): Calculated from standard entropies of products minus reactants; positive ΔS° favours spontaneity.

• Correct calculation of oxidation numbers in compounds and ions, including peroxides and metal hydrides.
• Correct identification of oxidation and reduction based on electron transfer and oxidation number changes.
• Correct identification of oxidising and reducing agents.
• Correct identification of disproportionation reactions.
• Correct use of Roman numerals to indicate oxidation numbers.
• Correct construction of full ionic equations from ionic half-equations.

• 💡Always check that the sum of oxidation numbers in a neutral compound equals zero and in an ion equals the charge of the ion.
• 💡Remember that oxidising agents are reduced (gain electrons) and reducing agents are oxidised (lose electrons).
• 💡When balancing half-equations, ensure the total charge on both sides is equal.
• 💡Practice identifying oxidation numbers in various contexts, especially for s- and p-block elements.
• 💡In Born-Haber cycle questions, always label each step with its correct enthalpy change (e.g., ΔH_at, ΔH_IE, ΔH_EA) and ensure arrows point in the correct direction. Use Hess's law to sum the cycle: ΔH_f = ΔH_at + ΔH_IE + ΔH_EA + ΔH_lattice (with signs).
• 💡When calculating ΔG, remember to convert temperatures to Kelvin and use consistent units (kJ mol⁻¹). For spontaneity, check the sign of ΔG: negative means spontaneous. Also, note that ΔG = 0 at equilibrium, so you can find the temperature at which a reaction becomes spontaneous by setting ΔG = 0.
• 💡For entropy calculations, always use standard entropies (S°) from data tables. Remember that entropy is a state function, so ΔS° = ΣS°(products) - ΣS°(reactants). A common mistake is forgetting to multiply by stoichiometric coefficients.

• Confusing the direction of electron transfer in oxidation and reduction.
• Incorrectly assigning oxidation numbers in complex ions or species.
• Failing to balance both atoms and charges when constructing ionic half-equations.
• Misidentifying the species being oxidised or reduced in a disproportionation reaction.
• Misconception: Lattice enthalpy is the same as the enthalpy of formation. Correction: Lattice enthalpy is the enthalpy change when gaseous ions form a solid lattice, while enthalpy of formation is from elements in standard states. They are related via the Born-Haber cycle but are not equal.
• Misconception: A reaction with a negative ΔH is always spontaneous. Correction: Spontaneity depends on both ΔH and ΔS. For example, an exothermic reaction with a large decrease in entropy may be non-spontaneous at high temperatures (ΔG > 0).
• Misconception: Entropy always increases in a reaction. Correction: Entropy can decrease, e.g., in the formation of a solid from gases (like NH₃(g) + HCl(g) → NH₄Cl(s)), but the reaction may still be spontaneous if ΔH is sufficiently negative.