THE ELECTRONIC THEORY OF
OXIDATION & REDUCTION
1. Oxidation & Reduction
Simple definitions of oxidation and reduction are based on the loss/gain of oxygen or
the loss/gain of hydrogen. Oxidation is the gain of oxygen or the loss of
hydrogen; reduction is the loss of oxygen or the gain of hydrogen. These
definitions can only be used when a chemical reaction involves hydrogen and
oxygen, and therefore their usefulness is limited.
A more basic and more useful definition of oxidation and
reduction is based on the loss/gain of electrons.
OXIDATION IS LOSS OF ELECTRONS
REDUCTION IS GAIN OF ELECTRONS
In reactions involving simple ions, it is usually easy to tell
whether electrons are lost or gained, but it is less easy to
tell when complex ions or covalent molecules are
involved. Oxidation number is a useful concept for helping
to decide in these more awkward cases.
2. Oxidation State (Number)
The oxidation state is used to express the oxidation state of an element, whether as
the uncombined element or when combined in a compound; it consists of a + or –
sign followed by a number, or it is zero. Oxidation states of metals are usually written
in roman numerals eg Iron (III) chloride.
Atoms of elements have no overall charge and are therefore given an oxidation state
of zero. When two elements combine, the atoms or ions of the more electropositive
element have a positive oxidation state, and those of the more electronegative
element a negative oxidation state. Elements become more electronegative the
higher their Group number and the lower their Period number; therefore, the most
electronegative element is fluorine.
The oxidation state of an element in a compound is equal to the charge which a
particle of the element would carry in the compound, assuming the compound is
ionic. This is a purely theoretical idea, and it is does not matter whether the
compound in question is really ionic or covalent.
IGCSE TOPIC 10.7b: REDOX REACTIONS
1
e.g.
compound
oxidation state
NaCl
CCl4
HBr
H2S
Na
C
H
H
+1
+4
+1
+1
Cl
Cl
Br
S
-1
-1
-1
-2
The following general rules are useful:
All free elements (i.e. those not combined with another element) have an
oxidation state of 0.
e.g. Na, Mg, Br2
In simple ions, the charge on the ion is equal to the oxidation state.
e.g.
ion
oxidation number
Na+
+1
Fe3+
+3
Br-1
O2-2
Since fluorine is the most electronegative element, it always has an oxidation
state in its compounds of –1.
Combined oxygen always has an oxidation state of –2, except when in
combination with fluorine or in peroxides.
e.g.
compound
oxidation numbers
Fe2O3
Mn2O7
CrO3
Fe +3
Mn +7
Cr +3
O -2
O -2
O -2
H2O2
F2O
H +1
O +2
O -1
F -1
BUT ….
Group I elements in their compounds always have an oxidation state of +1.
Group II elements in their compounds always have an oxidation state of +2.
Hydrogen in its compounds always has an oxidation state of +1, except when
it has combined with a reactive metal.
e.g.
compound
oxidation numbers
H2O
HCl
CH4
H +1
H +1
H +1
O -2
Cl -1
C -4
NaH
Na +1
H -1
BUT ….
The sum of the oxidation states of all the atoms in an uncharged molecule is
zero: in an ion, the sum is equal to the charge on the ion.
IGCSE TOPIC 10.7b: REDOX REACTIONS
2
e.g.
compound/ion oxidation states
NH3
NH4+
H2SO4
H +1 N -3
H +1 N -3
H +1 S +6 O -2
(-3+1+1+1=0)
(-3+1+1+1+1=+1)
(+1+1+6-2-2-2-2=0)
3. Number of Oxidation States
Many elements have several oxidation states:
e.g.
sulphur
chlorine
H2S
S
SCl2
SO2
SO3
+1 -2
0
+2 -1
+4 -2
+6 -2
HCl
Cl2
HOCl ClF3 KClO3 KClO4
+1-1
0
+1-2+1 +3-1
+1+5-2 +1+7-2
4. Redox Reactions
If, during a chemical reaction, the oxidation state of an element increases (i.e.
becomes more positive or less negative), then the element has lost electrons and
has been OXIDISED.
Conversely, if the oxidation state of an element decreases (i.e. becomes less
positive or more negative), then the element has gained electrons and has been
REDUCED.
Oxidation
-5 -4 -3 -2 -1 0 +1 +2 +3 +4 +5
Reduction
REDuction and OXidation occur together in what is called a REDOX reaction.
IGCSE TOPIC 10.7b: REDOX REACTIONS
3
Example:

work out the oxidation states for all the elements in the reaction

identify the oxidation and reduction steps

work out the total number of electrons transferred in each step: they should be
equal
oxidation: loss of 2e- x 1
Mg + 2HCl
0
MgCl2 + H2
+1 -1
+2 -1
0
reduction: gain of 1e- x 2
check: 2e- x 1 = 1e- x 2
Complete the equations for the remaining reactions in a similar way.
Mg + 2HCl
0
MgCl2 + H2
+1 -1
+2 -1
2Fe2+ + Cl2
+2
2Fe3+ + 2Cl-
0
+3
2Fe3+ + S2+3
-2
+2
0
Zn2+ + Cu
0
IGCSE TOPIC 10.7b: REDOX REACTIONS
-1
2Fe2+ + S
Cu2+ + Zn
+2
0
+2
4
0
OXIDATION & REDUCTION
QUESTION SHEET 1
Work out the oxidation state of each element in the following chemical formulae.
1.
Cu
…………………………………………………………………..
2.
NaCl
………………………………………………………………….
3.
CuS
………………………………………………………………….
4.
Cl2
………………………………………………………………….
5.
Fe3+
………………………………………………………………….
6.
Cr2O3
………………………………………………………………….
7.
H2O
………………………………………………………………….
8.
HNO3
………………………………………………………………….
9.
MnO4-
………………………………………………………………….
10.
K2CO3
………………………………………………………………….
11.
NaClO3
………………………………………………………………….
12.
SO42-
………………………………………………………………….
13.
NaNO2
………………………………………………………………….
14.
SOCl2
………………………………………………………………….
15.
Cu(NO3)2
………………………………………………………………….
16.
K2Cr2O7
………………………………………………………………….
17.
Al(OH)3
………………………………………………………………….
18.
K2MnO4
………………………………………………………………….
19.
Na3PO4
………………………………………………………………….
20.
Fe3O4
………………………………………………………………….
IGCSE TOPIC 10.7b: REDOX REACTIONS
5
OXIDATION & REDUCTION
QUESTION SHEET 2
Work out the oxidation state of each element in the following chemical formulae
1.
Mg
…………………………………………………………………..
2.
KBr
………………………………………………………………….
3.
CaS
………………………………………………………………….
4.
Br2
………………………………………………………………….
5.
Ba2+
………………………………………………………………….
6.
Fe2O3
………………………………………………………………….
7.
Na2O
………………………………………………………………….
8.
LiNO3
………………………………………………………………….
9.
MnO42-
………………………………………………………………….
10.
Rb2CO3
………………………………………………………………….
11.
KBrO3
………………………………………………………………….
12.
IO4-
………………………………………………………………….
13.
KNO2
………………………………………………………………….
14.
POCl3
………………………………………………………………….
15.
Sr(NO3)2
………………………………………………………………….
16.
K2CrO4
………………………………………………………………….
17.
Cr(OH)3
………………………………………………………………….
18.
KMnO4
………………………………………………………………….
19.
H3PO4
………………………………………………………………….
20.
CuCl2-
………………………………………………………………….
21.
Pb3O4
………………………………………………………………….
IGCSE TOPIC 10.7b: REDOX REACTIONS
6
OXIDATION & REDUCTION
QUESTION SHEET 3
Work out the oxidation state of each element in the following chemical formulae
1.
Ca
…………………………………………………………………..
2.
LiF
………………………………………………………………….
3.
MgO
………………………………………………………………….
4.
I2
………………………………………………………………….
5.
Cr3+
………………………………………………………………….
6.
Al2O3
………………………………………………………………….
7.
HF
………………………………………………………………….
8.
Ni(NO3)2
………………………………………………………………….
9.
CrO42-
………………………………………………………………….
10.
SrCO3
………………………………………………………………….
11.
NaClO4
………………………………………………………………….
12.
SO32-
………………………………………………………………….
13.
NaIO3
………………………………………………………………….
14.
XeF4
………………………………………………………………….
15.
Pb(OH)2
………………………………………………………………….
16.
K2MnO4
………………………………………………………………….
17.
Al2(SO4)3
………………………………………………………………….
18.
NaVO3
………………………………………………………………….
19.
H3PO3
………………………………………………………………….
20.
NH4+
………………………………………………………………….
IGCSE TOPIC 10.7b: REDOX REACTIONS
7
OXIDATION & REDUCTION
QUESTION SHEET 4
Examine each of the following redox equations. Work out the state of each element
in all the atoms, ions and molecules. Using these numbers, explain with reasons
which substance is oxidised and which substance is reduced
SO2 + 2Mg
2MgO + S
C + H2O
CO + H2
SnCl2 + HgCl2
Hg + SnCl4
H2 + Cl2
2HCl
2Fe3+ + 2I-
IGCSE TOPIC 10.7b: REDOX REACTIONS
2Fe2+ + I2
8
OXIDATION & REDUCTION
QUESTION SHEET 5
Examine each of the following reactions.

For each equation, work out the oxidation state of each element in all the
atoms, ions and molecules. Use these numbers to decide whether the change
taking place is a redox reaction or not.

Where a redox reaction occurs, indicate, with reasons, which species is
oxidised and which is reduced

Where the change is not a redox reaction, describe in one word the type of
change taking place.
1.
Mg(s) + Cl2(g)
2.
NaCl(s)
3.
2Fe3+(aq) + 2I-(aq)
4.
C(s) + H2O(g)
5.
Ag+(aq) + Cl-(aq)
MgCl2(s)
Na+(l) + Cl-(l)
IGCSE TOPIC 10.7b: REDOX REACTIONS
2Fe2+(aq) + I2(s)
CO(g) + H2(g)
AgCl(s)
9
6.
NaOH(aq) + HCl(aq)
7.
2H2(g) + O2(g)
8.
Cu(s) + 4HNO3(aq)
IGCSE TOPIC 10.7b: REDOX REACTIONS
NaCl(aq) + H2O(l)
2H2O(l)
Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)
10
REDOX EQUATIONS
Constructing Half –Equations
The half-equation shows either the oxidation or the reduction step of a redox
change. In a half-equation:

only one element changes its oxidation state

The loss/gain of electrons responsible for the change in oxidation state is
shown

the atoms of each element must balance

the total charge on both sides of the equation must be the same

the two half equations can be added together to give the full ionic equation
When sodium reacts with chlorine gas, sodium atoms lose an electron and are
oxidised to sodium ions; whereas chlorine atoms gain an electron and are reduced to
chloride ions. Since chlorine is a molecule (Cl2) both constituent atoms must gain an
electron resulting in 2 ions.
Na+ + e-
Na
Cl2
+
2e-
2Cl-
Since this latter process requires 2 electrons, two sodium atoms must also be
involved hence;
2Na+ + 2e-
2Na
Now the two half equations can be added together to show the full ionic equation
remembering to cancel the electrons.
Cl2

+
2e2Na
2Cl2Na+ + 2e2Na+
Cl2 + 2Na
IGCSE TOPIC 10.7b: REDOX REACTIONS
11
+
2Cl-
If during the addition of two half equations the same ion appears on both sides of the
equation i.e. has not changed in any way, these spectator ions should also be
cancelled.
Eg in reaction of potassium bromide with chlorine
Cl2(s) + 2KBr(aq)
2KCl(aq) + Br2
Chlorine atoms are gaining electrons to form chloride ions as in the previous
example.
Cl2 +
2e2Cl-
Whilst bromide ions in potassium bromide are losing electrons to turn into bromine
2K+
2Br -
+
Br2
2e- +
+
2K+
So combining the half equation gives
2K+
+
Cl2
2Br +
Br2
2e-
+
2e- +
2Cl-
Overall
Cl2
+
2Br -
IGCSE TOPIC 10.7b: REDOX REACTIONS
Br2
12
+
2Cl-
2K+
OXIDATION & REDUCTION
QUESTION SHEET 6
Examine each of the following reactions.

For each equation, write the two half equations and combine them to get the
overall ionic equation
1.
Mg(s) + Cl2(g)
2.
Br2(s)
3.
2H2(g) + O2(g)
4.
SO2 + 2Mg
5.
NaOH(aq) + HCl(aq)
MgCl2(s)
+ 2KI
IGCSE TOPIC 10.7b: REDOX REACTIONS
2KBr(aq) + I2
2H2O(l)
2MgO + S
NaCl(aq) + H2O(l)
13
Oxidising and reducing agents
An oxidising agent oxidises another substance and is itself reduced in the process.
A reducing agent reduces another substance and is itself oxidised in the process.
In the reaction
CuO
+
H2
Cu
+
H2O
CuO is losing oxygen and so is reduced. This happens when heated with hydrogen.
Hydrogen has reduced CuO to copper metal and has itself gained oxygen and
therefore been oxidised. You could also consider the oxidation states of each
substance and would draw the same conclusion.
Potassium dichromate is an oxidising agent and can oxidise ethanol to a substance
called ethanol. During this process the dichromate Cr2O72- (which is orange) is itself
reduced to green Cr3+.
Cr2O72-
Oxidation state +6
IGCSE TOPIC 10.7b: REDOX REACTIONS
Cr3+
oxidation state +3
14
Reduction
Summary Questions
Topic 7b Redox
1
2
IGCSE TOPIC 10.7b: REDOX REACTIONS
15
3
4
5
Chlorine reacts with iron (II) ions as shown below
Cl2 + 2Fe2+
2Fe3+ + 2Cl-
What type of reagent is chlorine in this reaction? Give a reason for your answer.
…………………………………………………………………………………………..
…………………………………………………………………………………………….
(2)
IGCSE TOPIC 10.7b: REDOX REACTIONS
16
6
7
IGCSE TOPIC 10.7b: REDOX REACTIONS
17
8
(b) Use the data in the table above to calculate the oxidation number of copper in copper
carbonate CuCO3.
…………………………………………………………………………………….
IGCSE TOPIC 10.7b: REDOX REACTIONS
18
(1)