Enthalpy change
Section: 5 Chemical Energetics | Syllabus: Cambridge AS Level Chemistry (9701)
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 TipΔH 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. |
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).
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 Neutralisation50.0 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 CombustionBurning 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 = 0.010 mol
ΔH = −11.29 ÷ 0.010 = −1129 kJ mol⁻¹
(Literature value ≈ −1367 kJ mol⁻¹ - difference due to heat losses to surroundings.)
Figure 5.2: Simple Calorimetry Setup
(Diagram of a polystyrene cup calorimeter: styrofoam cup inside a beaker (insulation), thermometer inserted through a lid, known volume of solution inside. Label: thermometer, insulating cup, reaction mixture, lid. Beside it, a spirit lamp calorimeter for combustion: copper can of water above a spirit lamp, thermometer in water. Label: copper can, water, spirit lamp, thermometer, draught shield.)
Exam TipExperimental values for ΔH⦵c are always less exothermic than data book values because: heat is lost to the surroundings, combustion may be incomplete, the experiment is not done under standard conditions. Always state assumptions: density of solution = 1 g cm⁻³, c of solution = 4.18 J g⁻¹ K⁻¹, no heat loss.
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