Chapter 9, Problem 9.102QE

Interpretation Introduction

Interpretation:

The approximate enthalpy change for the following reaction has to be determined.

  ${\text{Cl}}_{\text{2}}\left(\text{g}\right)+\text{CO}\left(\text{g}\right)\to {\text{Cl}}_{\text{2}}\text{CO}\left(\text{g}\right)$

Concept Introduction:

The bond dissociation energy or bond energy $\left(D\right)$ is defined as energy needed to break one mole of bonds in a gaseous species.

The equation that describes the bond dissociation for ${\text{H}}_{\text{2}}$ is as follows:

  ${\text{H}}_{\text{2}}\left(\text{g}\right)\to \text{H}\left(\text{g}\right)+\text{H}\left(\text{g}\right)$

Bond energies are always endothermic and have a positive sign. It takes energy to break a bond.

The enthalpy of reaction can be determined by the sum of the bond dissociation energies of all the reactants minus the sum of the bond dissociation energies of all products present in chemical reaction.

The equation to calculate enthalpy of reaction is as follows:

  $\Delta H_{\text{reaction}}=\left[\begin{array}{c}{\displaystyle \sum \left(\text{moles of bond broken}\times \text{bond energies of bond broken}\right)}\\ -{\displaystyle \sum \left(\text{moles of bond formed}\times \text{bond energies of bond formed}\right)}\end{array}\right]$

Negative sign in the equation depicts that bonds will form in the products. It is an exothermic process, so the energy charge is the negative of bond energy.