Thermochemistry
Thermochemistry: heat, enthalpy, and Hess's law
Energy is conserved. Enthalpy change ΔH is heat transferred at constant pressure; q = mcΔT gives the heat gained or lost by a substance, with m in grams and c in J/(g·°C). Coffee-cup calorimetry measures ΔH for solution reactions; bomb calorimetry measures q at constant volume, which equals ΔU, not ΔH. For reactions with gas moles changing, ΔH = ΔU + ΔnRT.
Sign conventions: q is negative when the system releases heat (exothermic, ΔH < 0) and positive when it absorbs heat (endothermic). Bond breaking is always endothermic; bond formation is always exothermic. Standard enthalpy of formation ΔHf° is zero for elements in their standard states, and ΔHrxn° = ΣnΔHf°(products) − ΣnΔHf°(reactants). Hess's law allows reactions to be added; when reversing a step, change the sign of ΔH, and when multiplying, multiply ΔH by the same factor.
Traps: use the stoichiometric coefficient, not the subscript, when summing formation enthalpies. Do not mix J and kJ. For a heating curve, temperature is constant during a phase change while added heat goes into breaking intermolecular forces (ΔHfus or ΔHvap). Also, a spontaneous reaction can be endothermic if entropy drives it-thermochemistry alone does not decide spontaneity.
Worked example: 50.0 g water warms from 20.0 °C to 30.0 °C. q = (50.0)(4.18)(10.0) = 2.09 × 10^3 J, or 2.09 kJ absorbed.
Rule of thumb: write the target reaction, then manipulate given reactions to match it.