Bond energy

[1][2][3] IUPAC defines bond energy as the average value of the gas-phase bond-dissociation energy (usually at a temperature of 298.15 K) for all bonds of the same type within the same chemical species.

The BDE, denoted by Dº(R—X), is usually derived by the thermochemical equation, This equation tells us that the BDE for a given bond is equal to the energy of the individual components that make up the bond when they are free and unbonded minus the energy of the components when they are bonded together.

These energies are given by the enthalpy of formation ΔHfº of the components in each state.

The enthalpy of formation of a large number of atoms, free radicals, ions, clusters and compounds is available from the websites of NIST, NASA, CODATA, and IUPAC.

Most authors use the BDE values at 298.15 K.[5] For example, the carbon–hydrogen bond energy in methane BE(C–H) is the enthalpy change (∆H) of breaking one molecule of methane into a carbon atom and four hydrogen radicals, divided by four.

[7] The bond-dissociation energies of several different bonds of the same type can vary even within a single molecule.

[11] Dividing the length of this bond by the sum of each boron atom's radius gives a ratio of

This ratio is slightly larger than 1, indicating that the bond itself is slightly longer than the expected minimum overlap between the two boron atoms' valence electron clouds.

This ratio is notably lower than 1, indicating that there is a large amount of overlap between the valence electron clouds of the two rhenium atoms.

[14] The extent of this effect is described by the compound's lattice energy, where a more negative lattice energy corresponds to a stronger force of attraction between the ions.

Generally, greater differences in electronegativity correspond to stronger ionic bonds.

Meanwhile, the compound sodium iodide (NaI) has a lower lattice energy of -704 kJ/mol with a similarly lower electronegativity difference of 1.73 between sodium and iodine.