Oxidation & Reduction
Oxidation and reduction take place together at the same time in the same reaction
These are called redox reactions
There are three definitions of oxidation. It is a reaction in which:
o Oxygen is added to an element or a compound
o An element, ion or compound loses electrons
o The oxidation state of an element is increased
There are three definitions of reduction. It is a reaction in which:
o Oxygen is removed from an element or a compound
o An element, ion or compound gains electrons
o The oxidation state of an element is decreased
Oxidation state
The oxidation state (also called oxidation number) is a number assigned to an atom or
ion in a compound which indicates the degree of oxidation (or reduction)
The oxidation state helps you to keep track of the movement of electrons in a redox
process
It is written as a +/- sign followed by a number.
Eg O2- means that it is an atom of oxygen that has an oxidation state of -2. It is not
written as O2- as this refers to the ion and its charge
Assigning the oxidation number
Oxidation number refers to a single atom or ion only
The oxidation number of a compound is 0 and of an element (for example Br in Br2)
is also 0
The oxidation number of oxygen in a compound is always -2 (except in peroxide R-
O-O-R, where it is -1)
For example, in FeO, oxygen is -2 then Fe must have an oxidation number of +2 as
the overall oxidation number for the compound must be 0
Ionic Equations
Ionic equations are used to show only the particles that actually take part in a reaction
These equations show only the ions that change their status during a chemical process,
i.e: their bonding or physical state changes
The other ions present are not involved and are called spectator ions
Writing ionic equations
For the neutralisation reaction between hydrochloric acid and sodium hydroxide:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
If we write out all of the ions present in the equation and include the state symbols,
we get:
H+(aq) + Cl- (aq)+ Na+(aq) + OH-(aq) → Na+ (aq)+ Cl-
(aq) + H2O(l)
The spectator ions are thus Na+ and Cl–. Removing these from the previous equation
leaves the overall net ionic equation:
H+(aq) + OH-(aq) →H2O(l)
This ionic equation is the same for all acid-base neutralisation reactions
Example redox equation: oxygen loss/gain
Zinc oxide + carbon → zinc + carbon monoxide
ZnO + C → Zn + CO
In this reaction the zinc oxide has been reduced since it has lost the carbon atom has
been oxidised since it has gained oxygen
Redox & Electron Transfer
Example redox equation: electron loss/gain and oxidation state
Zinc + copper sulphate → zinc sulphate + copper
Zn + CuSO4 → ZnSO4 + Cu
Writing this as an ionic equation:
Zn(s) + Cu2+(aq) + SO42-(aq) →Zn2+(aq) + SO42-(aq) +
Cu(s)
By analysing the ionic equation, it becomes clear that zinc has become oxidised as its
oxidation state has increased and it has lost electrons:
Zn(s) →Zn2+(aq)
Copper has been reduced as its oxidation state has decreased and it
has gained electrons:
Cu2+(aq) → Cu(s)
Exam Tip
Use the mnemonic OIL-RIG to remember oxidation and reduction in terms of the movement
of electrons: Oxidation Is Loss – Reduction Is Gain.
Oxidising & Reducing Agents
Oxidising agent
A substance that oxidises another substance, in so doing becoming itself reduced
Common examples include hydrogen peroxide, fluorine and chlorine
Reducing agent
A substance that reduces another substance, in so doing becoming itself oxidised
Common examples include carbon and hydrogen
The process of reduction is very important in the chemical industry as a means of
extracting metals from their ores
Example
CuO + H2 →Cu + H2O
In the above reaction, hydrogen is reducing the CuO and is itself oxidised, so
the reducing agent is therefore hydrogen
The CuO is reduced to Cu and has oxidised the hydrogen, so the oxidising agent is
therefore copper oxide
Redox Reactions
Identifying redox reactions
Redox reactions can be identified by the changes in the oxidation states when a
reactant goes to a product
Example
Chlorine + potassium iodide → potassium chloride +
iodine
Cl2 + 2KI → 2KCl + I2
Chlorine has become reduced as its oxidation state has decreased from 0 to -1 on
changing from the chlorine molecule to chloride ions:
Cl2(g) → 2Cl-(aq)
Iodine has been oxidised as its oxidation state has increased from -1 to 0 on
changing from iodide ions to the iodine molecule:
2I-(aq) → I2(s)
Identifying redox reactions by colour changes
The tests for redox reactions involve the observation of a colour change in the
solution being analyse
Two common examples are acidified potassium manganate (VII), and potassium
iodide
Potassium manganate (VII), KMnO4, is an oxidising agent which is often used to test
for the presence of reducing agents
When acidified potassium manganate (VII) is added to a reducing agent its colour
changes from pink-purple to colourless
Potassium iodide, KI, is a reducing agent which is often used to test for the presence
of oxidising agents
When added to an acidified solution of an oxidising agent such as aqueous chlorine or
hydrogen peroxide, the solution turns a brown colour due to the formation of iodine
