Non-covalent interactions
A hydrogen bond (H-bond), is a specific type of interaction that involves dipole-dipole attraction between a partially positive hydrogen atom and a highly electronegative, partially negative oxygen, nitrogen, sulfur, or fluorine atom (not covalently bound to said hydrogen atom). It is not a covalent bond, but instead is classified as a strong non-covalent interaction. It is responsible for why water is a liquid at room temperature and not a gas (given water's low molecular weight). Most commonly, the strength of hydrogen bonds lies between 0 - 4 kcal/mol, but can sometimes be as strong as 40 kcal/mol In solvents such as chloroform or carbontetrachloride one observes e.g. for the interaction between amides additive values of about 5 kJ/mol. According to L. Pauling the strength of a hydrogen bond is essentially determined by the electrostatic charges. Measurements of thousands complexes in chloroform or carbontetrachloride have led to additive free energy increments for all kind of donor-acceptor combinations.
Bond energy
In chemistry, bond energy ("E") or bond enthalpy ("H") is the measure of bond strength in a chemical bond. IUPAC defines bond energy as the average value of the gas-phase bond dissociation energies (usually at a temperature of 298 K) for all bonds of the same type within the same chemical species. For example, the carbon–hydrogen bond energy in methane "H"(C–H) is the enthalpy change involved with breaking up one molecule of methane into a carbon atom and four hydrogen radicals, divided by 4. Tabulated bond energies are generally values of bond energies averaged over a number of selected typical chemical species containing that type of bond. Bond energy ("E") or bond enthalpy ("H") should not be confused with bond-dissociation energy. Bond energy is the average of all the bond-dissociation energies in a molecule, and will show a different value for a given bond than the bond-dissociation energy would. This is because the energy required to break a single bond in a specific molecule differs for each bond in that molecule. For example, methane has four C–H bonds and the bond-dissociation energies are 435 kJ/mol for "D"(CH–H), 444 kJ/mol for "D"(CH–H), 444 kJ/mol for "D"(CH–H) and 339 kJ/mol for "D"(C–H). Their average, and hence the bond energy, is 414 kJ/mol, even though not a single bond required specifically 414 kJ/mol to be broken.