T - Bond Enthalpy (Lesson)

Bond Enthalpy

In previous lessons, Hess's Law was utilized to demonstrate how enthalpy changes for chemical reactions could be calculated using tabulated enthalpy of formation values. A related method is discussed in this lesson, but instead of utilizing standard enthalpies of formation, this method relies on the energy associated with the chemical bonds themselves. The bond energy, also known as bond dissociation energy, is the energy required to break a mole of covalent bonds within a molecule in the gaseous state. For example, a hydrogen molecule is comprised of one single bond, and the energy required to break this bond is 436 kJ/mol.

H₂ (g)  $\longrightarrow$  2 H (g)     $\Delta H$ = 436 kJ

All bond enthalpies are represented as endothermic processes with a positive sign for the enthalpy change because energy is always required to break a covalent chemical bond.  A table of common bond enthalpies Links to an external site. is accessible from the OpenStax textbook.

Calculating Bond Enthalpy

Hess's law can be utilized to calculate the change in enthalpy for a reaction in conjunction with the bond enthalpies listed in the table linked above. As previously mentioned, the breaking of chemical bonds is always an endothermic process whereas forming them is always exothermic.  When expressed in the form of an equation the relationship looks like this:

$\Delta H_{rxn}=\Sigma($bond energies of broken bonds) $-\Sigma$(bond energies of bonds formed)

Take, for example, the reaction between gaseous oxygen and hydrogen to form liquid water.

2 H₂ (g) + O₂ (g)  $\longrightarrow$  2 H₂O (l)   $\Delta H_{rxn}=$ ?

Because this method of calculating enthalpy of reaction relies on bond energies as opposed to enthalpies of formation for whole substances, it is necessary to understand precisely what types of bonds are present in both the reactant and products. Therefore, Lewis structures must be drawn in order to visualize this. The reaction below represents the same chemical process, but it expresses it in terms of Lewis structures instead of chemical formulas.

$\Delta H_{rxn}=\Sigma($bond energies of broken bonds) $-\Sigma$(bond energies of bonds formed)

$\Delta H_{rxn}=\Sigma$(2)(436) +(1)(498)] $-$ $\Sigma$[(4)(464)]

$\Delta H_{rxn}=$ 486 kJ/mol

It should be noted that while using bond enthalpies to calculate enthalpy of reaction, it results in a value that can be less accurate than that obtained from the use of enthalpies of formation. Calculating $\Delta H$ from enthalpies of formation is more accurate because the values for the actual substances involved in the chemical reaction are utilized. However, for bond enthalpy an average bond enthalpy is used. The bond enthalpies of covalent bonds may change from substance to substance, and the values represented in the table are simply averages from a wide variety of bonds from various compounds.  

You Try It!

In the following self-assessment activity, calculate $\Delta H_{rxn}.$ Click the plus sign to check your answer!

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