Energetics — Edexcel IGCSE Chemistry Modular - UNIT 1 Revision Notes & Topical Questions

Energy is conserved in chemical reactions

• it cannot be created or destroyed, only transferred.
• If energy is transferred to the surroundings, the products must have less energy than the reactants by the same amount.
• Temperature changes are measured with a thermometer.

Worked example. A student mixes chemicals and records the temperatures: (1) 10 cm³ NaOH + 10 cm³ HCl, 19 → 26 °C; (2) 2 g citric acid + 10 cm³ NaHCO₃, 20 → 16 °C; (3) 0.5 g Mg powder + 10 cm³ CuSO₄, 19 → 31 °C; (4) 3 cm Mg ribbon + 10 cm³ H₂SO₄, 19 → 21 °C. Experiments 1, 3 and 4 are exothermic (temperature increases) and experiment 2 is endothermic (temperature decreases).

Simple calorimetry measures the energy change of a reaction. Two types of experiment are used: reactions in solution, and combustion of fuels.

Reactions in solution (neutralisation, dissolving, displacement):

Assumptions

• the solution has the specific heat capacity of water (4.18 J g⁻¹ °C⁻¹) and density 1 g cm⁻³;
  • the heat capacity of the container is ignored;
  • the reaction is complete;
  • and heat losses are negligible.

Combustion of a fuel

Sources of error are heat loss to the surroundings, heat absorbed by the calorimeter, and incomplete combustion. Reduce them by putting the can close above the flame, using a lid and draught shields.

The heat energy change is Q = mcΔT, where m is the mass of the substance being heated (in g), c is the specific heat capacity (4.18 J g⁻¹ °C⁻¹ for water) and ΔT the temperature change (the same in °C and K).

Worked example 1. Excess iron powder is added to 100.0 cm³ of 0.200 mol dm⁻³ copper(II) sulfate and the temperature rises by 7.5 °C.

Worked example 2. 1.023 g of propan-1-ol (M = 60.11 g mol⁻¹) heats 200 g of water by 30 °C. Q = 200 × 4.18 × 30 = 25 080 J = 25.1 kJ. Moles = 1.023 ÷ 60.11 = 0.0170 mol, so ΔH = −25.1 ÷ 0.0170 = −1470 kJ mol⁻¹.

Worked example 3. The energy from 0.01 mol of propanol heats 250 g of water from 25 °C to 37 °C. Q = 250 × 4.18 × 12 = 12 540 J. ΔH = −12.54 ÷ 0.01 = −1254 kJ mol⁻¹.

Energy level diagrams show the energy of reactants and products on the vertical axis, with the progress of the reaction along the horizontal axis. In an exothermic reaction the products are lower than the reactants and an arrow points down (ΔH negative); in an endothermic reaction the products are higher and the arrow points up (ΔH positive).

Bond breaking is endothermic (energy is taken in) and bond making is exothermic (energy is given out). The enthalpy change is ΔH = energy taken in to break bonds − energy given out making bonds. If bond making releases more than bond breaking takes in, the reaction is exothermic (ΔH negative).

Worked example 4. H–H + Cl–Cl → 2 H–Cl, with bond energies H–H 436, Cl–Cl 242, H–Cl 431 kJ mol⁻¹. Energy in = 436 + 242 = 678 kJ. Energy out = 2 × 431 = 862 kJ. ΔH = 678 − 862 = −184 kJ, so the reaction is exothermic.

Worked example 5. 2HBr → H₂ + Br₂, ΔH = +103 kJ.

Practical: reaction between HCl and NaOH. Place 25 cm³ NaOH in a lidded polystyrene cup, measure its temperature, add 5 cm³ portions of HCl, stir and record the highest temperature each time. Plot temperature against volume of acid; the peak shows the volume needed for neutralisation, where the largest energy is released.