Exothermic reactions release heat to the surroundings (temperature rises). Examples: respiration, combustion, neutralisation, dissolving NaOH in water. (p.116)
Endothermic reactions absorb heat (temperature drops). Examples: photosynthesis, thermal decomposition of metal carbonates and nitrates, dissolving ammonium salts in water. (p.116)
Heat of reaction, ΔH = Hproducts − Hreactants, unit kJ mol−1. ΔH is negative for exothermic and positive for endothermic reactions. (p.117)
Exothermic: products have lower energy than reactants, e.g. Mg + H2SO4 → MgSO4 + H2, ΔH = −467 kJ mol−1. (p.118)
Endothermic: products have higher energy than reactants, e.g. N2 + O2 → 2NO, ΔH = +180 kJ mol−1. (p.118)
Exothermic: heat released in forming bonds is greater than heat absorbed to break bonds. Endothermic: heat absorbed to break bonds is greater than heat released on bond formation. (p.119)
Steps to calculate ΔH: find moles (n); calculate heat change Q = mcθ; find heat change per 1 mole; write ΔH with sign and kJ mol−1. (p.120)
n is the moles of precipitate (precipitation), metal displaced (displacement), water formed (neutralisation) or fuel burnt (combustion). (p.120)
Heat of precipitation is the heat change when 1 mole of precipitate is formed from its ions in aqueous solution, e.g. Ba2+ + SO42− → BaSO4, ΔH = −42 kJ mol−1. (p.121)
Assumptions: solution density = 1 g cm−3; specific heat capacity c = 4.2 J g−1 °C−1; no heat lost to surroundings; no heat absorbed by apparatus. (p.121)
Moles from solution: n = MV/1000 (V in cm3). Worked example: PbSO4 precipitation, Q = 200 × 4.2 × 3.0 = 2520 J, ΔH = −25.2 kJ mol−1. (p.122)
Heat of displacement is the heat change when 1 mole of a metal is displaced from its salt solution by a more electropositive metal, e.g. Fe + CuSO4 → FeSO4 + Cu, ΔH = −250 kJ mol−1. (p.125)
Worked example: Mg with FeSO4, Q = 840 J, n = 0.0125 mol, ΔH = −67.2 kJ mol−1. (p.126)
Heat of neutralisation is the heat change when 1 mole of water is formed from an acid and an alkali: H+ + OH− → H2O. (p.127)
Weak acids and alkalis ionise partially; some heat released is absorbed to ionise them completely, so the heat released is lower. (p.130)
H2SO4 (diprotic) + 2NaOH releases −114 kJ because 2 moles of water form, but heat of neutralisation per mole of water is still −57 kJ. (p.131)
Heat of combustion is the heat released when 1 mole of a substance is completely burnt in excess oxygen, e.g. CH4 + 2O2 → CO2 + 2H2O, ΔH = −394 kJ mol−1 (as given in the book). (p.133)
Heat of combustion of alcohols increases as the number of carbon atoms per molecule increases, because more CO2 and H2O are formed; successive members differ by one CH2. (p.134)
Worked example: C2H5OH, ΔH = −1376 kJ mol−1; heating 200 cm3 water by 50 °C needs 42 kJ, i.e. 1.38 g ethanol. (p.135)
3.3 Application of Exothermic and Endothermic Reactions
Exothermic applications: combustion of fuel, thermite reaction, Hot Can, gel pads. Endothermic applications: self-cooling beverages, sorbet, cold packs. (p.136)
Fuel value is the heat energy released when 1 g of fuel is completely burnt in excess oxygen. Fuel value (kJ g−1) = heat of combustion / molar mass. (p.137)
Example: octane, ΔH = −5500 kJ mol−1, molar mass 114, fuel value = 48.2 kJ g−1. (p.137)