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Loading home…An exothermic reaction releases heat to the surroundings, so the temperature of the surroundings rises. For exothermic reactions, ΔH is negative because the system loses enthalpy. An endothermic reaction absorbs heat from the surroundings, so the surroundings cool. For endothermic reactions, ΔH is positive. Combustion of fuels is typically exothermic; thermal decomposition of carbonates can be endothermic. Enthalpy profile diagrams show reactants higher than products for exothermic changes.
Hess’s law states that the total enthalpy change for a reaction is independent of the pathway taken. If a reaction can be written as the sum of steps, ΔH overall equals the sum of ΔH for the steps. It allows calculation of ΔH for reactions that are difficult to measure directly. Enthalpy cycles using formation or combustion enthalpies apply Hess’s law. The law follows from enthalpy being a state function. Born–Haber cycles for ionic lattices are an important application.
Standard enthalpy of formation ΔHf° is the enthalpy change when one mole of a compound forms from its elements in their standard states under standard conditions. ΔHf° of an element in its standard state is zero. Standard enthalpy of combustion ΔHc° is the enthalpy change when one mole of a substance burns completely in oxygen under standard conditions. Both are usually quoted at 298 K and 1 bar (or 1 atm in older data). They are used in Hess’s law cycles to find unknown reaction enthalpies. Combustion enthalpies are typically large and negative.
Bond enthalpy (bond dissociation enthalpy) is the enthalpy required to break one mole of a particular bond in gaseous molecules. Mean bond enthalpies are averages over similar compounds. Approximate ΔH = Σ(bonds broken) − Σ(bonds formed), using mean bond enthalpies. Breaking bonds is endothermic; forming bonds is exothermic. The method is approximate because mean values are not molecule-specific. It is useful when calorimetric data are unavailable.
Lattice enthalpy is the enthalpy change when one mole of an ionic crystal forms from gaseous ions (exothermic convention in many syllabuses) or the reverse endothermic dissociation. A more exothermic lattice enthalpy means a more stable ionic lattice. It increases with higher ionic charges and smaller ionic radii (stronger attraction). Born–Haber cycles determine lattice enthalpy indirectly using Hess’s law. Lattice enthalpy strongly influences melting points and solubility trends. Comparing theoretical and experimental values can reveal covalent character.
Temperature measures the average kinetic energy of particles; it is an intensive property. Heat is energy transferred because of a temperature difference; it is extensive. Specific heat capacity c is the heat required to raise the temperature of 1 g (or 1 kg) of a substance by 1 K. In calorimetry, q = mcΔT for a temperature change without phase change. Water has a high specific heat capacity, moderating climate and laboratory temperature changes. Confusing heat with temperature is a common student error.
Enthalpy H is a thermodynamic state function; ΔH = Hproducts − Hreactants for a process at constant pressure. If ΔH < 0, the reaction is exothermic and heat is released. If ΔH > 0, the reaction is endothermic and heat is absorbed. Standard enthalpy changes ΔH° refer to standard states at a stated temperature, commonly 298 K.
An enthalpy profile for an exothermic reaction shows reactants at higher enthalpy than products, with activation energy as the barrier to the transition state. For an endothermic reaction, products lie higher than reactants, so the reverse barrier differs. Catalysts lower activation energy but do not change overall ΔH.
Calorimetry measures heat transferred and, with care about heat losses and heat capacities, estimates ΔH. Sign conventions must be consistent: heat gained by the solution often equals heat lost by the reaction for exothermic cases. Clear diagrams help students connect energy, activation, and thermodynamic favourability (noting that ΔH alone does not decide spontaneity).
Hess’s law asserts that ΔH depends only on initial and final states, not on the path. Therefore enthalpy changes can be added algebraically like steps on a cycle.
Using formation enthalpies: ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants). Elements in standard states cancel because their ΔHf° is zero. Using combustion enthalpies, a cycle can be drawn in which reactants and products both burn to the same combustion products; then ΔHrxn = ΣΔHc°(reactants) − ΣΔHc°(products).
These methods find ΔH for reactions that are slow, side-reaction prone, or otherwise hard to measure. Born–Haber cycles extend the same idea to ionic solids. Students should always write a clear cycle, include signs carefully, and check units (kJ mol⁻¹).
Simple solution calorimetry mixes reactants in an insulated container and records temperature change of a known mass of water or solution. Heat absorbed by the calorimeter contents is q = mcΔT (plus calorimeter constant terms if calibrated). For an exothermic reaction, the solution warms; the heat released by reaction equals heat gained by surroundings (approximately). Dividing by moles of limiting reactant gives ΔH per mole.
Combustion calorimetry often uses a bomb calorimeter at constant volume, measuring ΔU; ΔH is obtained with corrections involving Δn_g. Major experimental issues include heat losses, incomplete combustion, evaporation, and inaccurate temperature readings. Insulation, stirring, and calibration improve accuracy.
Results should be reported with signs and standard-state notes. Comparison with data-book values tests experimental skill. Calorimetry connects macroscopic heat measurement to chemical energetics.
The enthalpy required to break a bond in the gas phase is the bond dissociation enthalpy. Tables often list mean bond enthalpies averaged over many molecules. To estimate ΔH, sum enthalpies of bonds broken (endothermic) and subtract enthalpies of bonds formed (exothermic).
This works reasonably for gas-phase reactions involving similar bonds. Limitations: mean values ignore molecular environment; reactants/products may not all be gases; resonance and strain are not included; phase changes need separate enthalpy terms. Therefore calorimetric or Hess’s-law values are preferred when available.
Still, bond-enthalpy estimates quickly show whether a reaction is likely exothermic and help explain fuel energy content. Stronger bonds in products than in reactants favour exothermic ΔH. Students should treat bond-enthalpy ΔH as approximate.
A Born–Haber cycle applies Hess’s law to the formation of an ionic solid from elements. Steps typically include atomisation of metal and non-metal, ionisation of metal, electron affinity of non-metal, and lattice enthalpy. The sum of the cycle equals ΔHf° of the ionic compound. Rearrangement yields an experimental lattice enthalpy.
Lattice enthalpy becomes more exothermic with higher ionic charges and smaller ions because electrostatic attraction increases. Comparing experimental and theoretical (purely ionic) lattice enthalpies indicates covalent character when polarisation is significant.
Born–Haber analysis explains relative stabilities of ionic compounds and supports discussion of Fajans’ rules. It is a capstone application of energetics, bonding, and Hess’s law together.