Bond Breaking and Bond Making in Chemical Reactions

When a chemical reaction has taken place it is usual that the products have either less stored energy than the reactants (overall exothermic) or more stored energy than the reactants (overall endothermic).

That is, with regard to the balanced chemical equation, the total molar enthalpy of the products is less (overall exothermic) or more (overall endothermic) than that of the reactants.

DH = SHproducts - SHreactants

Enthalpy level diagrams illustrate this:

Bond breaking and bond making...

When chemical reactions take place, chemical bonds are broken in the reactants and new chemical bonds are formed in the products. It is as a result of these processes that a reaction is overall exothermic (energy is given out to the surroundigs) or overall endothermic (energy is absorbed from the surroundings).

Bond breaking is obviously an endothermic process.

It is not quite so obvious that...

Bond making is an exothermic process.

But, if a particular bond in a molecule is broken and then reformed the same amount of energy must be involved, because the First Law of Thermodynamics (also known as the 'law of conservation of energy') must apply.

Thermodynamics in the study of transforming energy from one form to another.
The First Law of Thermodynamics states that: Energy can be transformed from one form to another , but cannot be created or destroyed. The total amount of energy in the universe never changes. Hess's Law (G.H. Hess, 1840) of 'constant heat summation' follows from the First Law.

F2(g)    ®    2F(g)      DH° = +158 kJ mol-1
2F(g)    ®    F2(g)     DH° = -158 kJ mol-1

Bond Enthalpy...

The molar bond enthalpy is the energy required to break one mole of bonds between pairs of atoms in the 'gaseous molecule'. Bond enthalpy values are usually expressed in kJ mol-1 of bonds broken.

Molar bond enthalpy is also called bond dissociation enthalpy, but there are other terms too that refer to this.

Bond enthalpies are often averaged values...

However, for more complicated molecules a precise bond enthalpy for a particular chemical bond depends to some extent on the environment in the molecule where the bond exists. That is, a precise bond enthalpy depends on what other atoms are attached to the two atoms of the bond to be broken. For this reason, tables of bond enthalpy values are averaged over those found in a large number of different compounds.


Bond enthalpy values apply to molecules in the gaseous state.

Because bond enthalpy values are averaged and also apply to molecules in the gaseous state, when they are used in calculations the answer will be only approximate. However, the approximation is often very good and such calculations are useful in predicting the overall enthalpy change, DH, for a chemical reaction. There is often close agreement with the standard value for an enthalpy change given in a chemical data book.

Using bond enthalpies...

Bond enthalpies can be used to calculate an overall enthalpy change, DH, for a chemical reaction. It is an over-simplification, but bonds are broken in the reactants and new bonds are formed in the products.

If it takes more energy to break bonds in the reactants than is released when new bonds form in the products then the reaction will be overall endothermic. A reaction is overall exothermic if more energy is released when new bonds form in products than is used when bonds in the reactants are broken.

This idea is illustrated below for the combustion of ethanol:

CH3CH2OH(g)  +  3O2(g)  ®  2CO2(g)  +  3H2O(g)

Note that all reactants and products are gaseous.

In the method above all the chemical bonds are broken in the reactants and new bonds formed in the products. There are very few reactions in which the reactant molecules are completely broken up into atoms, but this does not matter here because Hess's Law applies.

Calculating an Enthalpy Change of Reaction, DH...

This is quite straigtforward. First write out the balanced equation for the reaction showing full structural formulae. Now simply add together the molar bond enthalpies involved for the reactants to obtain a total endothermic value. Do the same for the products to obtain a total exothermic value. Finally, add the total endothermic and exothermic values together to find DH for the reaction.

Average Molar
Bond Enthalpy
kJ mol-1
Here bond enthalpies are defined endothermically.

Bond breaking:

Total endothermic value = (+347 x 1) + (+413 x 5) + (+358 x 1) + (+464 x 1) + (+498 x 3) = +4728 kJ

Bond making:

Total exothermic value = (-464 x 6) + (-805 x 4) = -6004 kJ

Sum total of bond breaking and bond making:

DHc = +4728 + - 6004 = -1276 kJ mol-1

Now compare this with the value of -1368 kJ mol-1 given in a chemical data book. The data book value is more exothermic by 92 kJ mol-1. This is explained by the use of averaged molar bond enthalpy values and that all reactants and products are gaseous, in the calculation of DH using bond enthalpies. For example, when water vapour forms liquid water the process is exothermic.
Now check your own calculations below...
Enthalpy Change, DH =

kJ mol-1


Bond Enthalpies and Bond Lengths

BondBond Length
Bond Enthalpy
kJ mol-1
C - C0.154+347
C = C0.134+612
C º C0.120+838
Bond Length and Bond Enthalpy are averaged values.

F - F0.142+158.0
Cl - Cl0.199+243.4
Br - Br0.228+192.9
I - I0.267+151.2

H - F0.092568.0
H - Cl0.127+432.0
H - Br0.141+366.3
H - I0.161+298.3

The molar bond enthalpies and bond lengths in the table above show that in general...

The longer a chemical bond the weaker it is.

For example, hydrogen chloride shows little tendencey to decompose into its constituent elements when heated. Strong heating of hydrogen bromide produces a brown colour of bromine vapour, while copious violet fumes of iodine form when a hot glass rod is plunged into a gas jar of hydrogen iodide.

HI(g)    ®    ½H2(g)   +   ½I2(g)

The stability of the hydrogen halides to thermal decomposition therefore decreases in the order:

HCl   >   HBr   >   HI

and this is due to the progressive decrease in the H ¾ X bond enthalpy.

glass rod

Hydrogen Iodide

With regard to the diatomic halogen molecules (X2), fluorine has an abnormally low molar bond enthalpy. This low value is in part explained by the repulsion between the non-bonding electrons on the fluorine atoms, the other halogen molecules having longer bond lengths making this repulsive force is less significant.