Learning outcomes
You should be able to:
■■ calculate oxidation numbers of elements in
compounds and ions
■■ describe and explain redox processes in terms of
electron transfer and changes in oxidation number
■■ use changes in oxidation numbers to help balance
chemical equations.
Introduction
Some types of reactions can cost a lot of money due
to the damage they cause. Rusting is an oxidation
reaction that destroys about 20% of iron and steel
every year. Rust is hydrated iron(III) oxide. This forms
when iron reacts with oxygen in the presence of water.
Another costly example of oxidation is the reaction
between hydrogen and oxygen that is used to propel
some types of rockets into space. In this reaction, the
hydrogen is oxidised – but the oxygen is also reduced.
In fact, oxidation and reduction always take place
together, in what we call redox reactions.
Figure below A redox reaction is taking place when the fuel in
the Space Shuttle’s rockets burns.
A simple definition of oxidation is gain of oxygen by an element.
For example, when magnesium reacts with oxygen, the magnesium combines with oxygen to form magnesium oxide. Magnesium has been oxidised.
2Mg(s) + O2(g)--> 2MgO(s)
A simple definition of reduction is loss of oxygen. When copper(II) oxide reacts with hydrogen, this is the equation for the reaction:CuO(s) + H2(g)--> Cu(s) + H2O(l)
Copper(II) oxide loses its oxygen. Copper(II) oxide has been reduced.
But if we look carefully at the copper oxide/hydrogen equation, we can see that oxidation is also taking place.
The hydrogen is gaining oxygen to form water. The hydrogen has been oxidised.
We can see that reduction and oxidation have taken place together.
Oxidation and reduction always take place together.
We call the reactions in which this happens redox reactions.
Redox reactions are very important. For example, one redox reaction – photosynthesis – provides food for the entire planet, and another one – respiration – keeps you alive. both are redox reactions.
We can also define reduction as addition of hydrogen to a compound and oxidation as removal of hydrogen from a compound. This is often seen in the reaction of organic compounds .
There are two other ways of finding out whether or not a substance has been oxidised or reduced during a chemical reaction:
■■ electron transfer
■■ changes in oxidation number.
Redox and electron transfer
Half-equations
We can extend our definition of redox to include reactions involving ions.
Oxidation Is Loss of electrons.
Reduction Is Gain of electrons.
The initial letters shown in bold spell oIL RIG. This may help you to remember these two definitions
In another example iron reacts with copper(II) ions,
Cu2+ in solution to form iron(II) ions, Fe2+ and copper.
There are two other ways of finding out whether or not a substance has been oxidised or reduced during a chemical reaction:
■■ electron transfer
■■ changes in oxidation number.
Redox and electron transfer
Half-equations
We can extend our definition of redox to include reactions involving ions.
Oxidation Is Loss of electrons.
Reduction Is Gain of electrons.
The initial letters shown in bold spell oIL RIG. This may help you to remember these two definitions
Sodium reacts with chlorine to form the ionic compound sodium chloride.
2Na(s) + Cl2(g) 2NaCl(s)
We can divide this reaction into two separate equations,
one showing oxidation and the other showing reduction.
We call these half-equations.
When sodium reacts with chlorine:
■■ Each sodium atom loses one electron from its outer shell.
Oxidation is loss of electrons (OIL). The sodium atoms have been oxidised.
Na à Na+ + e-
This half-equation shows that sodium is oxidised.
It is also acceptable to write this half-equation as:
Na – e-à Na+
■■ Each chlorine atom gains one electron to complete its outer shell.
one showing oxidation and the other showing reduction.
We call these half-equations.
When sodium reacts with chlorine:
■■ Each sodium atom loses one electron from its outer shell.
Oxidation is loss of electrons (OIL). The sodium atoms have been oxidised.
Na à Na+ + e-
This half-equation shows that sodium is oxidised.
It is also acceptable to write this half-equation as:
Na – e-à Na+
■■ Each chlorine atom gains one electron to complete its outer shell.
Reduction is gain of electrons (RIG).
The chlorine atoms have been reduced.Cl 2 + 2e- à 2Cl–
This is a half-equation showing chlorine being reduced.
There are two chlorine atoms in a chlorine molecule, so
two electrons are gained.
This is a half-equation showing chlorine being reduced.
There are two chlorine atoms in a chlorine molecule, so
two electrons are gained.
In another example iron reacts with copper(II) ions,
Cu2+ in solution to form iron(II) ions, Fe2+ and copper.
Fe(s)+ Cu2+ (aq) à Fe2+(aq) + Cu(s)
■■ Each iron atom loses two electrons to form an Fe2+ ion. The
iron atoms have been oxidised.
Fe àFe2+ + 2e–
It is also acceptable to write this half-equation as:
Fe – 2e– à Fe2+
■■ Each iron atom loses two electrons to form an Fe2+ ion. The
iron atoms have been oxidised.
Fe àFe2+ + 2e–
It is also acceptable to write this half-equation as:
Fe – 2e– à Fe2+
■ Each copper(II) ion gains two electrons. The copper ions
have been reduced.
Cu2+ + 2e– àCu
have been reduced.
Cu2+ + 2e– àCu
Balancing half-equations
We can construct a balanced ionic equation from two half equations by balancing the numbers of electrons lost and gained and then adding the two half-equations together.
The numbers of electrons lost and gained in a redox reaction must be equal.
We can construct a balanced ionic equation from two half equations by balancing the numbers of electrons lost and gained and then adding the two half-equations together.
The numbers of electrons lost and gained in a redox reaction must be equal.
