Enthalpy change, ΔH

· Enthalpy change, ΔH = the heat energy change of a reaction at constant pressure.
· Chemical reactions are accompanied by energy transfers between the reacting system and surroundings.
· Exothermic reactions release heat energy; ΔH is negative.
· Endothermic reactions absorb heat energy; ΔH is positive.
· Units for enthalpy change are usually kJ mol⁻¹.
· Always include the sign of ΔH: e.g. −57 kJ mol⁻¹, not just 57 kJ mol⁻¹.

This diagram shows how reactants and products differ in enthalpy. Products lower than reactants indicate an exothermic reaction, while products higher than reactants indicate an endothermic reaction. Source

Reaction pathway diagrams

· A reaction pathway diagram shows the change in enthalpy as reactants form products.
· The vertical axis is enthalpy / energy; the horizontal axis is progress of reaction.
· Activation energy, Ea = minimum energy needed for a reaction to occur.
· In an exothermic profile, products are at a lower enthalpy than reactants; ΔH < 0.
· In an endothermic profile, products are at a higher enthalpy than reactants; ΔH > 0.
· ΔH is the vertical energy difference between reactants and products.
· Ea is the vertical energy difference from reactants to the highest point / transition state on the curve.

Standard conditions and enthalpy notation

· Standard conditions in this syllabus: 298 K and 101 kPa.
· The standard symbol is ⦵, e.g. ΔH⦵.
· Standard enthalpy change means the enthalpy change measured under standard conditions.
· Use correct symbols:
· ΔHr = enthalpy change of reaction.
· ΔHf = enthalpy change of formation.
· ΔHc = enthalpy change of combustion.
· ΔHneut = enthalpy change of neutralisation.

Key enthalpy definitions

· Enthalpy change of reaction, ΔHr = enthalpy change when a reaction occurs according to the stoichiometric equation.
· Enthalpy change of formation, ΔHf = enthalpy change when 1 mole of a compound is formed from its elements.
· Enthalpy change of combustion, ΔHc = enthalpy change when 1 mole of a substance is completely burned in oxygen.
· Enthalpy change of neutralisation, ΔHneut = enthalpy change when 1 mole of water is formed from an acid–alkali neutralisation reaction.
· Exam definitions must include 1 mole, correct reaction type, and standard conditions if the symbol ⦵ is used.

Bond breaking and bond making

· Energy transfers occur because chemical reactions involve breaking bonds and making bonds.
· Bond breaking is endothermic because energy is absorbed; ΔH is positive.
· Bond making is exothermic because energy is released; ΔH is negative.
· Overall ΔH depends on the balance between energy absorbed to break bonds and energy released when new bonds form.
· If more energy is released making bonds than absorbed breaking bonds, the reaction is exothermic.
· If more energy is absorbed breaking bonds than released making bonds, the reaction is endothermic.

These diagrams show why breaking bonds requires energy and forming bonds releases energy. They help explain why bond enthalpy calculations compare the total energy for bonds broken with the total energy for bonds formed. Source

Calculating ΔHr using bond energies

· Bond energy = energy required to break 1 mole of a particular covalent bond in the gaseous state.
· Bond energies are quoted as positive values because they refer to bond breaking.
· Formula: ΔHr = Σ bond energies of bonds broken − Σ bond energies of bonds formed.
· Count bonds carefully using the balanced equation.
· Multiply each bond energy by the number of bonds broken or formed.
· Some bond energies are exact; many are average bond energies, so calculated ΔH values may be approximate.
· Bond energy calculations usually apply to reactions involving gaseous covalent substances.

Experimental enthalpy calculations

· Use q = mcΔT to calculate heat energy transferred.
· q = heat energy transferred, usually in J.
· m = mass of solution, usually in g.
· c = specific heat capacity, usually 4.18 J g⁻¹ K⁻¹ for aqueous solutions unless told otherwise.
· ΔT = temperature change in K or °C; the size of the change is the same in both.
· Convert from J to kJ by dividing by 1000.
· Use ΔH = −mcΔT / n to find molar enthalpy change.
· n = number of moles of the substance reacting, usually the limiting reagent or the substance named in the enthalpy definition.
· The negative sign links temperature rise to an exothermic reaction: if solution temperature increases, ΔH is negative.

Calorimetry uses a measured temperature change to calculate heat transferred. This supports the exam formula q = mcΔT and helps explain why insulation and heat loss affect experimental ΔH values. Source

Practical and exam method tips

· Always record initial temperature, final temperature, and calculate ΔT correctly.
· For a temperature rise, the reaction is usually exothermic, so ΔH is negative.
· For a temperature fall, the reaction is usually endothermic, so ΔH is positive.
· Use total solution mass if solutions are mixed; if density is assumed 1.00 g cm⁻³, then volume in cm³ ≈ mass in g.
· Use the balanced equation to decide the correct number of moles for ΔH.
· Quote final answers with suitable significant figures and units.
· Common errors: missing the minus sign, using cm³ as mass without stating density assumption, forgetting to convert J to kJ, or using the wrong n.