Higher and foundation tier
Consider the reaction between hydrogen and oxygen to make hydrogen oxide (water).
Two molecules of hydrogen react with one molecule
oxygen to make two molecules of water. If you study the image carefully, you will notice that
the hydrogen atoms, which were once part of a
hydrogen molecule in the reactants, are now separated and joined to an atom of oxygen,
in the products. Similarly, the two oxygen atoms,
which were joined together in a molecule of oxygen, are now separated
and joined to atoms of hydrogen.
This tells us that before any reaction can happen all the covalent bonds holding the atoms together must be broken. However, the breaking of bonds is an endothermic process; it requires an input of energy. You can imagine that dismantling and breaking apart molecules consisting of strong covalent bonds requires a lot of energy.
The table below list the bond energies for H-H, O=O and O-H bonds, the bond energy is the amount of energy needed to break 1 mole of bonds to form individual atoms. The higher the bond energy, the stronger the bond, as more energy is needed to break it and separate the atoms.
|Bond energy (kJ/mol)||436||498||463|
You can see that you need 498 kilojoules of energy to breaks 1 mole of O=O bonds and separate them into individual atoms.
Remember the law of conservation of energy. Energy cannot be created or destroyed, only change from one form to another. If it takes 498 kJ/mol of energy to break the covalent bonds holding the oxygen molecule together, then what do you think will happen if you reverse the above equation and join the two atoms together to form a molecule of oxygen?
Bond formation is exothermic, it releases heat energy to
the surroundings. It is simply the opposite of bond breaking in terms of energy change.
If a bond has a bond energy of 100 kJ/mol then it needs 100 kJ/mol to
break the bonds and
100 kJ/mol will be released if you form these bonds.
We can draw an energy profile diagram for the reaction below:
We can simplify the diagram above to give a simple graph to show the difference between exothermic and endothermic reactions in terms of the enthalpy of reaction (that is the amount of heat energy release), see image below:
The energy profile diagrams show how the energy stored in the reactants
and products changes as
the reaction takes place. For all chemical reactions, both exothermic
and endothermic, the
reactants need to be supplied with energy to break the
bonds in the reactants, this is the
activation energy. Once all the bonds in the
broken new bonds can form in the products,
remember bond formation releases energy, the
stronger the bonds formed in the products, the more
energy will be released.
In an endothermic reaction, more energy is required to break the bonds in the reactants than is released by bond formation in the products. So the products have more energy stored in their bonds than the reactants. This additional energy is absorbed from the surroundings. In an exothermic reaction more energy is released by bond formation than is required to break the bond in the reactants. This additional energy is released back into the surroundings as heat.
The actual amount of energy released is simply the difference between the amount of energy needed for bond breaking and the amount released by bond formation in the products. (Note higher tier students will need to be able to calculate the energy changes taking place during reactions using bond energy data.)