Bond Enthalpies

Chemical reactions involve breaking bonds in the reactants and making new bonds in the products.Breaking a bond needs energy, so it is endothermic. Making a bond releases energy, so it is exothermic. The overall enthalpy change depends on the balance between the two: if more energy is released making bonds than is taken in breaking them, the reaction is exothermic.Bond enthalpy is the enthalpy change when one mole of a particular bond is broken, with all species in the gaseous state. For example, the H-H bond enthalpy is +436 kJ mol^-1:H₂(g) → 2H(g) ΔH = +436 kJ mol^-1Bond enthalpies are always positive, because breaking a bond always needs energy. A larger bond enthalpy means a stronger bond.For a diatomic molecule such as H₂ or Cl₂, the value is exact because that bond is only found in one substance. Most bonds, however, occur in many different compounds. A C-H bond in methane is not exactly the same strength as a C-H bond in ethane or ethanol, because neighbouring atoms affect it. Even the four C-H bonds in methane need different energies to break one after another.So data books give a mean bond enthalpy: the enthalpy change to break one mole of a given type of bond, averaged over a range of compounds. For C-H it is about +413 kJ mol^-1.Because they are averages, calculations using mean bond enthalpies give approximate answers. The gaseous-state condition matters too: the energy needed to separate molecules in a liquid or solid is not included.