The enthalpy change for a gas-phase chemical reaction, ΔrH, 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:
ΔrH = ∑Ebonds broken – ∑Ebonds formed
Because the bond energy values provided in the Appendix are averaged over many different molecules for each type of bond, this calculation is not exact. But it provides a good estimate for the enthalpy change of any specific reaction. Consider this balanced reaction:
H2(g) + Cl2(g) → 2 HCl(g)
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 Dbond, we have:
ΔrH = [436 kJ/mol + 242 kJ/mol] – [2(431 kJ/mol)] = -184 kJ/mol
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 that all molecules are far from each other (which means that reactants and products must be in the gas phase). Additional enthalpy changes occur when a gas condenses to a liquid or a solid or dissolves into a solution; these transformations are not accounted for by bond enthalpies.
Check Your Learning
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 strength, 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.