D8.1 Bond Energy and Reaction Enthalpy Change

Earlier you learned that for two argon atoms a long distance apart the energy is higher than when the same two atoms are 377 pm apart. This difference in energy is related to the fact that vaporization of argon is an endothermic process—an input of energy is required as vaporization occurs or the temperature of the liquid decreases. You also learned that the energy of two bonded atoms is lower than when the two atoms are separated. Overcoming intermolecular attractions and breaking a chemical bond are similar processes in that both require an input of energy. Forming a chemical bond, on the other hand, releases energy from the chemical system to its surroundings. These ideas can be applied to predicting the enthalpy change when a chemical reaction occurs.

The enthalpy change for a gas-phase chemical reaction, Δ_rH, equals the sum of the enthalpy required to break each of the bonds in the reactant molecules (energy in, positive sign) plus the sum of the enthalpy released when each of the bonds in the product molecules forms (energy out, negative sign). This can be expressed mathematically as:

Because the bond energy values provided in the Appendix are averaged over many different molecules for each type of bond, such a calculation using those values is not exact. But it provides a very good estimate of the enthalpy change of a reaction. For example, consider this balanced reaction:

One H–H bond (436 kJ/mol) and one Cl–Cl bond (242 kJ/mol) are broken; two H-Cl bonds (431 kJ/mol each) are formed. Representing bond enthalpies by D_bond, we have:

The number of bonds formed here is the same as the number of bonds broken. But because the bonds in the products are stronger than those in the reactants, the reaction has a net release (negative sign) of 184 kJ for every mole of reaction as written. The energy released increases the temperature of the surroundings (the reaction is exothermic).

Note that bond enthalpy calculations assume all molecules are far from each other (which means that reactants and products must be in the gas phase). If the molecules were closer together, then there would be additional enthalpy effects due to the intermolecular attractions. Enthalpy changes related to intermolecular attractions are not accounted for by bond enthalpies.

Below are two general rules for predicting whether a chemical reaction releases energy (is exothermic):

• If there are more bonds in the product molecules than in the reactant molecules and the bonds have about the same strengths, the reaction is likely exothermic.
• If there are stronger bonds in the product molecules than in the reactant molecules and the number of bonds is the same in reactants and products, the reaction is likely exothermic.