A given bond type, such as carbon-hydrogen, does not have the exact same strength in every molecule containing it. The surrounding atoms and overall molecular structure shift the bond’s strength slightly from one compound to the next. Because of this, tabulated bond energies are averages, calculated from measurements across many different compounds that share that bond type.
This makes bond energy a useful estimate rather than an exact value for any single molecule. Breaking bonds always requires energy, an endothermic process, while forming new bonds always releases energy, an exothermic process. Chemists use this pattern to estimate the overall enthalpy change of a reaction using Hess’s law.
The rule is simple. The reaction’s enthalpy change roughly equals the energy needed to break the reactant bonds. Then subtract the energy released forming the product bonds.
If more energy is released forming new bonds than was used breaking old ones, the reaction is exothermic overall. In general, higher bond energy corresponds to a stronger and typically shorter bond between two atoms. Double and triple bonds have higher bond energies than single bonds between the same two elements.
More shared electron pairs pull the atoms together more tightly. A common misconception is that bond energy values give exact reaction enthalpies. Because the values are averages, calculations using bond energies are always approximations, not precise thermodynamic results.