Bond Enthalpy

Bond enthalpy is the amount of energy required to break one mole of a given TYPE of bond between two atoms, with both atoms in the gaseous state. For a simple DIATOMIC molecule, bond enthalpy and bond dissociation energy are numerically identical — e.g. H2's bond enthalpy is 435.8 kJ mol⁻¹, and the heteronuclear diatomic HCl's is 431.0 kJ mol⁻¹. A larger bond enthalpy always signals a STRONGER bond — for instance, the N≡N bond in N2 is stronger than the O=O bond in O2, consistent with N2's higher bond order (3 vs 2, section 5.5.6). For a POLYATOMIC molecule, though, bond enthalpy is NOT simply equal to any one bond's dissociation energy, because breaking successive, apparently-identical bonds one after another actually costs slightly DIFFERENT amounts of energy each time — the local electronic environment around the remaining atoms changes once the first bond has broken. In water, for example, breaking the first O–H bond ($\text{H}_2\text{O}\rightarrow\text{H}+\text{OH}$) costs 502 kJ mol⁻¹, while breaking the second, now-different O–H bond ($\text{OH}\rightarrow\text{H}+\text{O}$) costs only 427 kJ mol⁻¹ — so the tabulated (average) bond enthalpy of an O–H bond, $(502+427)/2=464.5$ kJ mol⁻¹, is an AVERAGE of many such successive dissociation-energy measurements taken across a large number of dif …