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-dissociation energy
The bond dissociation energy is an enthalpy change of a particular chemical process, namely homolytic bond cleavage, and "bond strength" as measured by the BDE should not be regarded as an intrinsic property of a particular bond type but rather as an energy change that depends on chemical context. For instance, Blanksby and Ellison cites the example of ketene (HC=CO), which has a C=C bond dissociation energy of 79 kcal/mol, while ethylene (HC=CH) has a bond dissociation energy of 174 kcal/mol. This vast difference is accounted for by the thermodynamic stability of carbon monoxide (CO), formed upon C=C bond cleavage of ketene. The difference in availability of spin states upon fragmentation further complicates the use of BDE as a measure of bond strength for head-to-head comparisons, and force constants have been suggested as an alternative.