What Is Enthalpy? 什么是焓？

What Is Enthalpy? 什么是焓？

Enthalpy (H) is the total heat content of a system at constant pressure. It is a state function — its value depends only on the current state, not on how the system reached it. In OCR Chemistry, what matters is the enthalpy change (ΔH): the heat transferred during a reaction at constant pressure. The fundamental relationship is ΔH = ΔU + PΔV. For reactions in solution with negligible volume change, ΔH ≈ ΔU.

焓 (H) 是恒压下系统的总热含量，是一个状态函数。OCR考试的核心是焓变 (ΔH)——恒压反应中传递的热量。你永远不需要测量绝对焓值，只需要计算变化量。

Exothermic vs Endothermic 放热与吸热

Exothermic (ΔH < 0): Energy released to surroundings. Products are more stable (lower enthalpy) than reactants. Temperature of surroundings increases. Examples: combustion, neutralisation, reaction of Group 1 metals with water. In energy profile diagrams, products sit below reactants.

Endothermic (ΔH > 0): Energy absorbed from surroundings. Products have higher enthalpy. Temperature of surroundings decreases. Examples: thermal decomposition (CaCO₃ → CaO + CO₂), photosynthesis, dissolving ammonium nitrate. Products sit above reactants in energy profile diagrams.

Key principle: Bond breaking is always endothermic (requires energy); bond forming is always exothermic (releases energy). The net ΔH depends on the balance. OCR examiners frequently test this by asking you to interpret energy profile diagrams.

Standard Enthalpy Changes 标准焓变

All standard enthalpy changes are measured at 298 K, 100 kPa, with substances in their standard states. The plimsoll symbol (⦵) denotes standard conditions. OCR requires precise definitions — always include "one mole" and standard conditions.

ΔH⦵f — Standard Enthalpy of Formation (标准生成焓): Enthalpy change when one mole of a compound forms from its elements in their standard states. ΔH⦵f of any element = 0. E.g. C(s) + O₂(g) → CO₂(g), ΔH⦵f = −393.5 kJ mol⁻¹.

ΔH⦵c — Standard Enthalpy of Combustion (标准燃烧焓): Enthalpy change when one mole of a substance burns completely in excess oxygen. Always exothermic. E.g. CH₄ + 2O₂ → CO₂ + 2H₂O(l), ΔH⦵c = −890 kJ mol⁻¹. Water is liquid under standard conditions.

ΔH⦵neut — Standard Enthalpy of Neutralisation (标准中和焓): Enthalpy change when one mole of water forms from acid + base. Strong acid–strong base: ≈ −57 kJ mol⁻¹. Weaker acids/bases give less exothermic values (energy used for ionisation).

Hess's Law 盖斯定律

Hess's Law: Total enthalpy change for a reaction is independent of the route taken (consequence of enthalpy being a state function). You can calculate unknown ΔH values by combining known enthalpy changes along alternative pathways.

Formation method (生成焓法): ΔH⦵r = ΣΔH⦵f(products) − ΣΔH⦵f(reactants). Construct a cycle with elements at the bottom — arrow up to products minus arrow up to reactants.

Combustion method (燃烧焓法): ΔH⦵r = ΣΔH⦵c(reactants) − ΣΔH⦵c(products). Construct a cycle with combustion products (CO₂ + H₂O) at the bottom — arrow down from reactants minus arrow down from products.

OCR pitfall: Formation is products minus reactants; combustion is reactants minus products. Never memorise blindly — always draw the cycle.

Bond Enthalpy 键焓

Bond enthalpy is the energy required to break one mole of covalent bonds in the gaseous state. OCR distinguishes between mean bond enthalpy (average across many compounds, e.g. C–H ≈ +413 kJ mol⁻¹) and actual bond enthalpy (specific to one molecule — breaking the first C–H in CH₄ requires more energy than the mean).

Calculating ΔH: ΔH = Σ(bonds broken) − Σ(bonds formed). Bond breaking is endothermic (+), bond forming is exothermic (−).

OCR classic question: "Why does the calculated ΔH differ from the experimental value?" Answer: mean bond enthalpies are averages, not specific to the molecule's exact bond environment.

Calorimetry 量热法

The core equation: q = mcΔT, where q = heat transferred (J), m = mass of solution (g, ≈ volume in cm³ for dilute aqueous solutions), c = 4.18 J g⁻¹ K⁻¹ (specific heat capacity of water), ΔT = temperature change.

Convert q to ΔH: ΔH = −q / n, where n = moles of limiting reactant. The negative sign is critical — forgetting it is the #1 OCR calorimetry error.

Common sources of error: Heat loss to surroundings (use polystyrene cup + lid + draught shield), incomplete combustion (use copper calorimeter), evaporation (cover with lid), and assuming solution's specific heat capacity equals pure water. OCR also tests cooling curve extrapolation to compensate for heat loss.

Born-Haber Cycles 玻恩-哈伯循环

Born-Haber cycles apply Hess's Law to ionic compound formation, breaking ΔH⦵f into discrete steps: atomisation of metal and non-metal (endothermic), ionisation energies (endothermic), electron affinities (1st EA exothermic for most non-metals, 2nd EA always endothermic), and lattice enthalpy (always exothermic).

Key relationship: ΔH⦵f = ΔH⦵at(metal) + ΔH⦵at(non-metal) + ΣIE + ΣEA + ΔH⦵L

OCR trick: The second electron affinity of oxygen (O⁻ + e⁻ → O²⁻) is endothermic due to electron-electron repulsion. This appears almost every year.

Gibbs Free Energy 吉布斯自由能

Enthalpy alone cannot predict reaction feasibility. Gibbs free energy combines enthalpy and entropy: ΔG = ΔH − TΔS. A reaction is feasible when ΔG < 0. But feasibility ≠ rate — a reaction can be thermodynamically feasible yet kinetically hindered (high activation energy). OCR examiners love testing this distinction.

Four temperature-dependence scenarios:
• ΔH < 0, ΔS > 0 → feasible at all temperatures.
• ΔH > 0, ΔS < 0 → never feasible.
• ΔH < 0, ΔS < 0 → feasible only at low T (|ΔH| > |TΔS|).
• ΔH > 0, ΔS > 0 → feasible only at high T (|TΔS| > |ΔH|).

The temperature where ΔG = 0: T = ΔH / ΔS. Remember to convert ΔS from J K⁻¹ mol⁻¹ to kJ K⁻¹ mol⁻¹.

OCR Exam Tips 考试必备技巧

1. Always state standard conditions (298 K, 100 kPa) when defining standard enthalpy changes. Missing this costs marks.

2. Draw the Hess cycle before every calculation. OCR awards marks for correctly drawn cycles and it prevents sign errors.

3. Units matter. Calorimetry gives q in joules, but ΔH is quoted in kJ mol⁻¹. Divide by 1000.

4. The "one mole" rule. All standard enthalpy definitions specify exactly one mole. Omit this and lose marks.

5. Mean bond enthalpies give approximate ΔH. This question appears yearly — the answer: averages across many compounds, not specific to the molecule.

6. Feasibility (ΔG < 0) does not guarantee reaction occurs. Kinetics (activation energy) is a separate consideration.

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