JEE Chemistry Thermodynamics Notes Set 05

Bond Enthalpy

1. Bond enthalpy is the energy needed to break one mole of a bond in the gas phase. It is always positive, and the value in a given molecule is an average over similar compounds.
2. \(\Delta_rH=\sum(\text{bonds broken})-\sum(\text{bonds formed})\) for gas-phase reactions.
3. Enthalpy of formation from atoms: \(\Delta_fH=\text{atomisation energy of the elements}-\text{sum of bond enthalpies of the compound}\). The atomisation energy of the elements uses the sublimation enthalpy of graphite for carbon and half the dissociation enthalpy of \(\ce{H2}\) per hydrogen atom.
4. Count bonds from the structure: ethane has 6 C-H and 1 C-C; ethene has 4 C-H and 1 C=C; ethyne has 2 C-H and 1 triple bond; water has 2 O-H bonds.
5. Resonance energy = actual \(\Delta_fH\) minus the value calculated for the structure with localised bonds. A negative value (extra stability) is typical, for example about -150 kJ/mol for benzene.

Born-Haber Cycle

1. Applies Hess's law to an ionic solid: \(\Delta_fH=\Delta_{sub}H(M)+IE(M)+\tfrac12\Delta_{diss}H(X_2)+\Delta_{eg}H(X)+\text{(lattice enthalpy of formation, negative)}\).
2. Lattice enthalpy is the energy released when gaseous ions form one mole of the crystal (or the same magnitude needed to separate it).
3. Enthalpy of solution \(=-\text{(lattice enthalpy of formation)}+\Delta H_{hyd}(\text{cation})+\Delta H_{hyd}(\text{anion})\).
4. The second electron gain enthalpy of oxygen is positive, so lattice enthalpy of oxides must be very large to stabilise \(\ce{O^2-}\).

Remember

1. Standard state of \(\ce{H+(aq)}\) formation enthalpy is zero.
2. Monoclinic sulphur, diamond and black phosphorus are not reference forms, so their \(\Delta_fH\) is not zero.
3. Use \(\Delta U\) and \(C_V\) in a rigid adiabatic vessel; the final temperature is found by assuming all the heat of reaction heats the products.
4. A reaction with mixed exothermic and endothermic parts stays at constant temperature when the heat released equals the heat absorbed.