 Higher tier only

### Bond energies

Hydrogen and oxygen react in a highly exothermic reaction to make hydrogen oxide or water. This is shown below. Before the reaction can start the chemical bonds in the reactants have to be broken (see diagram). The amount of energy needed to break the bonds is called the bond energy. Bond energies have been calculated for all bonds and a search of the internet or a quick look in a chemistry data book will show tables of bond energies. Do not try to remember any of the values for bond energies as they will always be given to you in any exam questions set. Bond energy data is shown below for some common bonds.

Bond Bond energykJ/mol Bond Bond energy kJ/mol
C-H 413 C-C 348
H-H 436 Cl-Cl 242
O-H 463 O=O 495

The units of bond energy are kilojoules per mole. You can see from the table that to break 1 mole of C-H bond requires an input of 413 kJ of heat energy or when 1 mole of C-H bonds are formed then 413 kJ of heat energy will be released. Remember the law of conservation of energy, energy cannot be created or destroyed e.g. consider the C-H bonds in methane. ### Enthalpy change ΔH

To calculate the energy change (amount of heat energy released) or enthalpy change (ΔH) as it is often called, for the combustion of hydrogen to form water we need to think about all the bonds that are broken and formed. To begin with until you get good at working these out energy changes I would recommend that you draw out model equations similar to the one shown below. It just allows you to keep a simple tally of all the bonds that are being broken and formed. So from the image above we can see that we need to break 2 moles of H-H bonds and 1 mole of O=O bonds to start the reaction by breaking the bonds in the reactants. Once these bonds are broken then the bonds in the products, 4x O-H bonds will form. We can put this info in a table:

Bonds broken Bond energy kJ/mol Bonds formed Bond energy kJ/mol
2 x H-H 2 x 413 =826 4 x O-H 4 x 463=1852
1 x O=O 495
Total energy required to break all the bonds in the reactants:
= 826 +495
=1321kJ.
Total energy released by bond formation:
= 1852kJ

The data for bond energies is the energy required to break the bond, when the same bond is formed as described above the same amount of energy is released but it is given a negative sign. The negative sign indicates that the system is losing energy, that is to say the process is exothermic. So in the above example the total energy released by bond formation would be -1852kJ, with the negative sign indicating that the system, or chemicals are losing energy to the surroundings. So to calculate the energy or enthalpy change for the reaction (enthalpy is simply the amount of heat released at constant pressure and the triangle sign, Δ, is the letter delta from the Greek alphabet. It is used in chemistry to mean a "change in") we simply add the two values together.

Energy change for the reaction = energy required to break bonds - energy released by bond formation. =1321 -1852 = -531 kJ.
The reaction has a negative energy change so it is exothermic.

### Example 2- The reaction of hydrogen and fluorine

Calculate the energy (enthalpy) change for the explosive reaction between hydrogen and fluorine gas. The equations for the reaction and all bond energies needed are shown below.

 Bond Bond energy (KJ/mol) H-H H-F F-F 436 567 155 Bonds broken Bond energy kJ/mol Bonds formed Bond energy kJ/mol
1 x H-H 413 2 x H-F 2 x 567=1134
1 x F-F 155
Total energy required to break all the bonds in the reactants:
= 413 + 155
=568kJ
Total energy released by bond formation:
= 1134kJ

Energy change for the reaction = energy required to break bonds - energy released by bond formation.
=568 -1134 = -566 kJ/mol.
The reaction has a negative energy change so it is exothermic.

Generally we could say - The stronger the bonds in the products the more energy is released by bond formation. The weaker the bonds in the reactants the less energy is required to break them. So for a highly exothermic reaction you need products with strong bonds and reactants with weak bonds. The opposite would be true for endothermic reactions.