Higher and foundation tiers
Consider the reaction between hydrogen and oxygen to make hydrogen oxide (water).
Two molecules of hydrogen react with one molecule
of
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 take place all the covalent bonds holding the atoms
together in the reactants must be broken. However
the breaking of covalent 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 the H-H, O=O and O-H bonds; the bond energy is the amount of energy needed
to break 1 mole of bonds
in a molecule to form individual atoms. The higher the bond energy the stronger the bond and the greater amount of energy needed to break it and separate the atoms.
Bond | H-H | O=O | O-H |
---|---|---|---|
Bond energy (kJ/mol) | 436 | 498 | 463 |
You can see that you need 498 kilojoules of energy to break 1 mole of O=O bonds and separate the oxygen molecule into two individual atoms.
Remember the law of conservation of energy. Energy cannot be created or destroyed; it can only change from one form to another. If it takes 498 kJ/mol of energy to break the covalent bonds holding the oxygen molecules together then what do you think will happen if you reverse the above equation and join the two moles of oxygen atoms together to form 1 mole of oxygen molecules
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 chemical 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
reactants are
broken new bonds can form in the products;
remember bond formation releases energy and 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.)