Everything from the 4 lessons of this topic on a few pages: the key idea of each lesson, the diagram to remember, the points to learn and the traps examiners set. Read it, cover it, test yourself.
An enthalpy change, ΔH, is the heat energy exchanged with the surroundings at constant pressure. Exothermic reactions release heat, so ΔH is negative; endothermic reactions absorb heat, so ΔH is positive. Standard values (⊖) are for 298 K, 101 kPa and 1 mol dm−3 solutions, with substances in their standard states. Learn four definitions exactly: reaction (the molar quantities in the equation), formation (1 mol of compound from its elements), combustion (1 mol burned completely in oxygen) and neutralisation (1 mol of water formed).
The standard enthalpy change of formation of water and the standard enthalpy change of combustion of hydrogen are the same equation, H2(g) + ½O2(g) → H2O(l), so they have the same value, −286 kJ mol−1. Writing 2H2 + O2 → 2H2O doubles the value and fits neither definition.
The heat transferred to or from the surroundings is q = mcΔT, where m is the mass of water or solution heated (not the mass of reactant), c = 4.18 J g−1 K−1 and ΔT is the temperature change. Divide by the amount in moles of the reactant that was used up (the limiting one) and convert J to kJ: ΔH = −q ÷ n. The minus sign means a temperature rise gives a negative (exothermic) ΔH. Heat loss to the surroundings makes experimental values less exothermic than data book values, most of all in combustion experiments.
In q = mcΔT, m is the mass of the water or solution whose temperature changes, never the mass of fuel or solid added. Remember too that q comes out in joules: divide by 1000 for kJ, and add the minus sign yourself when the temperature rose.
Bond energy is the energy needed to break one mole of a particular covalent bond in gaseous molecules. Breaking bonds is endothermic and forming bonds is exothermic, so ΔH = Σ(bonds broken) − Σ(bonds formed). For bonds such as C–H the value depends on the molecule, so data books quote a mean bond energy averaged over many compounds. That, and substances that are not gases under standard conditions, is why bond energy answers differ slightly from experimental ΔH values.
It is broken minus formed, not the other way round; reversing it flips the sign. And count bonds per molecule and per mole of equation: 2H2O contains four O–H bonds, and CO2 contains two C=O double bonds, not two single C–O bonds.
Hess's law: the total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same. Build a cycle with the unknown change along the top and a route through measurable changes. With formation data the common point is the elements below: ΔHr = ΣΔHf(products) − ΣΔHf(reactants). With combustion data the common point is the combustion products below: ΔHr = ΣΔHc(reactants) − ΣΔHc(products). Multiply each value by its coefficient and reverse the sign of any arrow you travel against.
The two formula shortcuts have opposite orders: formation is products minus reactants, combustion is reactants minus products. If unsure, draw the cycle and follow the arrows, changing the sign whenever you travel against one. Never forget to multiply by the coefficients, and remember that elements have ΔHf = 0.