Question 10.16: Using Bond Energies to Predict Exothermic and Endothermic Re...
Using Bond Energies to Predict Exothermic and Endothermic Reactions
One of the steps in the formation of monochloromethane (Example 10-15) is the reaction of a gaseous chlorine atom (a chlorine radical) with a molecule of methane. The products are an unstable methyl radical and HCl(g). Is this reaction endothermic or exothermic?
CH_4 + \text{.}Cl(g) \longrightarrow \text{.}CH_3 (g) + HCl(g)
Analyze
In the reaction, one \begin{matrix} C-H \end{matrix} bond is broken for every \begin{matrix} H-Cl \end{matrix} bond formed. Thus, we must compare the bond energies for the \begin{matrix} C-H \end{matrix} and \begin{matrix} H-Cl \end{matrix} bonds to decide whether the reaction is endothermic or exothermic
For every molecule of CH_4 that reacts, one \begin{matrix} C-H \end{matrix} bond breaks, requiring 414 kJ per mole of bonds; and one \begin{matrix} H-Cl \end{matrix} bond forms, releasing 431 kJ per mole of bonds. Because more energy is released in forming new bonds than is absorbed in breaking old ones, we predict that the reaction is exothermic.
Assess
In the example above, we had to break only \begin{matrix} C-H \end{matrix} bonds and form only \begin{matrix} H-Cl \end{matrix} bonds. Most reactions involve breaking and forming several types of bonds, and so it is usually not obvious whether the reaction will be exothermic or endothermic. In such cases, we must calculate ΔH, by using equation (10.26), to see whether ΔH > 0 or ΔH < 0.
\Delta H =\Delta H(\text { bond breakage })+\Delta H \text { (bond formation) } \approx \Sigma BE \text { (reactants) }-\Sigma BE \text { (products) } (10.26)