Thermochemical Changes
Core ideas
Thermochemistry tracks energy as heat and enthalpy, ΔH. Define the system and surroundings; q has opposite signs for the two. Calorimetry uses q = mcΔT for the surroundings, often the solution, then q_system = -q_surroundings. Molar enthalpy is ΔH per mole of a specified substance, so identify the limiting reactant and the stoichiometric coefficient. Exothermic processes release heat, raise temperature, and have ΔH < 0; endothermic processes absorb heat and have ΔH > 0.
Hess's law lets you add reactions and their ΔH values. When reversing a reaction, reverse the sign; when multiplying, multiply ΔH. Standard enthalpy of formation, ΔH°f, is for one mole of compound from elements in standard states; elements in their standard states have ΔH°f = 0. Bond energies estimate gas-phase reactions: energy absorbed to break bonds minus energy released forming bonds.
Traps: using the mass of solute instead of total solution; forgetting the calorimeter heat capacity; mixing kJ and J; reporting ΔH per gram when per mole is asked; and assuming a temperature rise always means the reaction itself is exothermic without considering dissolution or phase change.
Worked example: 100.0 g of solution (c = 4.19 J/g°C) warms by 5.00°C: q = (100.0)(4.19)(5.00) = 2.10×10^3 J = 2.10 kJ absorbed by the solution, so the reaction q = -2.10 kJ. If 0.0500 mol reacted, ΔH = -2.10/0.0500 = -42.0 kJ/mol.