Bond energy calculations

In a nutshell

Each type of bond has a bond energy. Bond energies can be used to work out the overall energy change of a reaction and determine if a reaction is endothermic or exothermic. Exothermic reactions have a negative overall energy change. Endothermic reactions have a positive overall energy change.

Breaking and forming bonds

Chemical reactions involve breaking bonds in the reactants and forming new bonds to produce products. Energy transfers occur due to the breakage or formation of bonds.

Breaking an existing bond requires energy, therefore it is an endothermic process which absorbs energy from the surroundings.

Forming new bonds releases energy into the surroundings and is an exothermic process.

A reaction is exothermic reaction if more energy is released when new bonds form than is absorbed to break old bonds. A reaction is endothermic if more energy is absorbed to break old bonds than is released when new bonds form.

Bond energies

Definition

Every type of bond has a bond energy. Bond energy refers to the amount of energy required to break one mole of a specific bond. The bond energy of a particular bond will not be the same across all compounds; bond energies will vary slightly depending on the compound.

The overall energy change of a reaction can be calculated using bond energies.

$Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space form$

procedure

1.	First work out the energy needed to break all the bonds present. This is done by adding the bond energies of all the bonds in the reactants.
2.	Work out the energy released when new bonds form. This is done by adding the bond energies of all the bonds in the products.
3.	To get the overall energy change of a reaction, subtract the energy released when new bonds formed from the energy required to break old bonds.

The unit for energy change is $kJ \space mol^{-1}$ .

Bond energy for exothermic reactions

If the overall energy change of a reaction is negative, more energy is released than absorbed, which means that the reaction is exothermic.

Example

Hydrogen reacts with oxygen to form water.

Bond	$H-H$	$H-O$	$O=O$
Bond energy $(kJ \space mol^{-1})$ $(KJ \space mol^{-1})$ $(KJ \space mol^{-1})$	$436$	$464$	$498$

$_{H-H }^{H-H} \space + \space O=O \longrightarrow \space H-O-H \space \space+ \space\space H-O-H$

First, calculate the amount of energy needed to break all the bonds present in the reactants. There are two hydrogen molecules and one oxygen molecule:

$Energy \space needed \space to \space break \space bonds= 2(H-H) + 1(O=O)=\space 2(436) \space + \space 498 =1370 \, kJ \space mol^{-1}$

Then calculate the energy released when the products form. There are two water molecules:

$Energy \space released \space when \space bonds \space form= 4(O-H)=\space 4(464) =1856 \, kJ \space mol^{-1}$

Finally, subtract the energy released when products form from the energy required to break the bonds in the reactants:

$Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space form$ $Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space for$

$Overall \space energy \space change= 1370 \space - \space 1856 = \space -486 \, kJ \space mol^{-1}$

The overall energy change for this reaction is $\underline{-486 \,kJ \space mol^{-1}}$ therefore this reaction is exothermic.

Bond energy for exothermic reaction

If the overall energy change of a reaction is positive, more energy is absorbed than is released, which means that the reaction is endothermic.

Example

Nitrogen reacts with oxygen to form nitric oxide.

Bond	$N\equiv N$	$O=O$	$N=O$
Bond energy $(kJ \space mol^{-1})$	$945$	$498$	$630$

$630\Large N\equiv N \space + \space O=O \longrightarrow \space 2N=O$ $\Large N\equiv N \space + \space O=O \longrightarrow \space 2N=O$

$N\equiv N \space + \space O=O \longrightarrow \space 2N=O$

First, calculate the amount of energy needed to break all the bonds present in the reactants. There is one nitrogen molecule and one oxygen molecule:

$Energy \space needed \space to \space break \space bonds= (N \equiv N) \space + (O=O) \space = \space 945 \space + \space 498 =1443 \, kJ \space mol^{-1}$

$Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space form$ $Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space for$

Then calculate the energy released when the products form. There are two nitric oxide molecules:

$Energy \space released \space when \space bonds \space form= 2(N=O) = 2(630)=1260 \space kJ \space mol^{-1}$

Finally, subtract the energy released when products form from the energy required to break the bonds in the reactants:

$Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space form$ $Overall \space energy \space change = Energy \space needed \space to \space break \space bonds \space - \space Energy \space released \space when \space bonds \space for$

$Overall \space energy \space change= 1443 \space - \space 1260 = \space 183 \space kJ \space mol^{-1}$

The overall energy change of this reaction is $\underline{183 \,kJ \space mol^{-1}}$ therefore this reaction is endothermic.