0.13 Reaction energy and bond energy  (Page 7/7)

These observations reinforce our observation that there is no single energy associated with a single type of bond, e.g. a C-H bond. On the other hand, we can see that the bond energies for similar bonds in similar molecules are close to one another. The C-H bond energies in the three chloromethanes above illustrate this quite well. The average C-H bond energy in the three chloromethanes molecules is 412 kJ/mol, which is close to each of the individual bond energies. Likewise, the average of the C-H bond energies in methane is (1663 kJ/mol)/4 = 416 kJ/mol, which is also close to the value above and is not that far from the value for breaking the first C-H bond.

If we observe the bond energy for a specific type of bond (e.g. a C-H bond) in many molecules, we generally find that we can approximate the bond energy in any specific molecule by an average of the energies of similar bonds over many molecules. This means that it is valuable to have a table of the average energy of each type of bond, since these will be close to the value of the energy of that bond in any given molecule. These data are shown in [link] .

These average bond energies are very informative because they can be used to estimate the heat of a reaction without measuring all of the required bond energies.

Consider for example the combustion of methane to form water and carbon dioxide:

We can estimate the heat of this reaction by using average bond energies. We must break four C-H bonds at an energy cost of approximately 4 × 412 kJ/mol and two O2 bonds at an energy cost of approximately 2 × 496 kJ/mol. Forming the bonds in the products releases approximately 2 × 743 kJ/mol for the two C=O double bonds and 4 × 463 kJ/mol for the O-H bonds. Net, the heat of reaction is thus approximately ΔHº = 1648+992-1486-1852 = –698 kJ/mol. This is only a rough approximation to the actual heat of combustion of methane, –890 kJ/mol. Therefore, we cannot use average bond energies to predict accurately the heat of a reaction. But we can get an estimate, and this may be sufficiently useful. Moreover, we can use these calculations to gain insight into the energetics of the reaction. For example, [link] is strongly exothermic, which is why methane gas (the primary component in natural gas) is an excellent fuel. From our calculation, we can see that the reaction involved breaking six bonds and forming six new bonds. The bonds formed, particularly the C=O bonds, are substantially stronger than the bonds broken, and this accounts for the net release of significant energy during the reaction.

Review and discussion questions

1. Consider the reaction

   N ₂ O ₄ (g) → 2 NO ₂ (g)

   Draw Lewis structures for each of N ₂ O ₄ and NO ₂ . On the basis of these structures, predict whether the reaction is endothermic or exothermic, and explain your reasoning.

2. Why is the bond energy of H ₂ not equal to ΔH _f º of H ₂ ? For what species is the enthalpy of formation related to the bond energy of H ₂ ?
3. Suggest a reason why ΔHº for the reaction

   CO ₂ (g) → CO (g) + O (g)

   is not equal to ΔHº for the reaction

   CO (g) → C (g) + O (g)

4. Determine whether the reaction is exothermic or endothermic for each of the following circumstances:
   1. The combustion of the products is more exothermic than the combustion of the reactants.
   2. The enthalpy of formation of the products is greater than the enthalpy of formation of the reactants.
   3. The total of the bond energies of the products is greater than the total of the bond energies for the reactants.