Enthalpy change
Section: 5 Chemical Energetics | Syllabus: Cambridge AS Level Physics 9702
Exothermic and Endothermic Reactions Chemical reactions are accompanied by enthalpy changes (ΔH) - the heat energy transferred between the system (the reaction) and the surroundings at constant pressure.
Exothermic Reaction A reaction that releases heat energy to the surroundings. ΔH is negative. The products are at a lower energy than the reactants. Endothermic Reaction A reaction that absorbs heat energy from the surroundings.
ΔH is positive. The products are at a higher energy than the reactants. Exothermic examples: combustion, neutralisation, many oxidation reactions, respiration. Endothermic examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate in water.
ΔH = H(products) − H(reactants). If products are lower in energy → ΔH negative → exothermic. Reaction Pathway Diagrams A reaction pathway diagram (energy profile diagram) shows how the enthalpy of the system changes as reactants are converted to products, via the transition state.
Activation energy (Eₐ) - the minimum energy that colliding particles must have for a reaction to occur. It is the energy required to break bonds and reach the transition state (the peak of the curve).
For an exothermic reaction: the products are lower in enthalpy than the reactants. ΔH = negative. For an endothermic reaction: the products are higher in enthalpy than the reactants. ΔH = positive. The activation energy is always positive regardless of whether the reaction is exo- or endothermic.
Figure 5.1: Reaction Pathway Diagrams - Exothermic and Endothermic (Two diagrams side by side. Left - Exothermic: reactants at higher enthalpy level, products lower, smooth curve peaking at transition state.
Label Eₐ (from reactants to peak) and ΔH (negative, from reactants to products). Right - Endothermic: reactants lower, products higher. Label Eₐ and ΔH (positive). Both diagrams: x-axis = reaction coordinate / progress of reaction, y-axis = enthalpy / H.) Exam Tip ΔH on the diagram is the overall enthalpy change from reactants to products - it does not include the activation energy hump.
Eₐ is measured from the reactants level to the top of the curve (the transition state), not from the products. Standard Conditions and Standard Enthalpy Changes Standard Conditions (⦵) A temperature of 298 K (25°C) and a pressure of 101 kPa (approximately 1 atm).
Standard conditions are denoted by the symbol ⦵ (plimsoll symbol). Term Symbol Definition Standard enthalpy change of reaction ΔH⦵ᵣ The enthalpy change when molar quantities of reactants, as shown in the equation, react under standard conditions.
Standard enthalpy change of formation ΔH⦵f The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions. By definition, ΔH⦵f of any element in its standard state = 0.
Standard enthalpy change of combustion ΔH⦵c The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions. Always exothermic (negative). Standard enthalpy change of neutralisation ΔH⦵neut The enthalpy change when an acid and a base react to form one mole of water under standard conditions.
Always exothermic. For strong acid + strong base ≈ −57 kJ mol⁻¹. Key Point - Formation ΔH⦵f must produce exactly one mole of the compound from elements in their standard states. For example: C(s) + O₂(g) → CO₂(g) ΔH⦵f = −394 kJ mol⁻¹.
The equation must be written to give 1 mol CO₂ - use fractions if necessary (e.g. ½O₂). Calculating Enthalpy Changes from Experimental Data Enthalpy changes can be measured experimentally using a calorimeter.
The heat transferred to or from the surroundings (usually water or solution) is measured and used to calculate ΔH. Specific Heat Capacity, c The energy required to raise the temperature of 1 g of a substance by 1 K (or 1°C).
Units: J g⁻¹ K⁻¹. For water/dilute aqueous solution: c = 4.18 J g⁻¹ K⁻¹. The two key equations are: q = mcΔT - heat transferred (J), where m = mass of solution (g), c = specific heat capacity (J g⁻¹ K⁻¹), ΔT = temperature change (K or °C).
ΔH = −q ÷ n - enthalpy change per mole (J mol⁻¹ or kJ mol⁻¹), where n = moles of limiting reagent. The negative sign converts from the surroundings' perspective (q) to the system's perspective (ΔH). Sign Convention If the temperature of the surroundings increases , the reaction is exothermic → q is positive → ΔH is negative.
If temperature decreases , the reaction is endothermic → q is negative → ΔH is positive. Enthalpy of Neutralisation 50.0 cm³ of 1.00 mol dm⁻³ HCl is mixed with 50.0 cm³ of 1.00 mol dm⁻³ NaOH. Temperature rises by 6.8°C.
m = 100.0 g (total solution, assuming density = 1 g cm⁻³) | c = 4.18 J g⁻¹ K⁻¹ q = 100.0 × 4.18 × 6.8 = 2842 J = 2.84 kJ n(H₂O formed) = 0.0500 × 1.00 = 0.0500 mol ΔH = −2.84 ÷ 0.0500 = −56.8 kJ mol⁻¹ Enthalpy of Combustion Burning 0.46 g of ethanol (M = 46 g mol⁻¹) heats 200 g of water from 20.0°C to 33.5°C.
q = 200 × 4.18 × 13.5 = 11 286 J = 11.29 kJ n(ethanol) = 0.46 ÷ 46 = 0.0…
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