Chemistry Reference
In-Depth Information
•
Compounds have no charge overall. Hence the
oxidation states of all the individual elements in a
compound must add up to 0. The oxidation states
of elements in compounds can vary, as seen in
Table 3.6. It is possible to recognise which of the
different oxidation states a metal element is in by
the colour of its compounds (Figure 3.18).
In this reaction the iron(ii)
ii
) ions have been
oxidised to iron(
iii
) ions (increase in oxidation state)
and the manganate(
vii
) ions have been reduced to
manganese(
ii
) ions (decrease in oxidation state)
which are very pale pink. Hence the manganate(
vii
)
ions are the oxidising agent and the iron(ii)
ii
) ions are
the reducing agent.
It is possible to recognise redox processes by looking
at the oxidation states on the two sides of the chemical
equation for a reaction. For example, in the reaction
between magnesium and dilute sulfuric acid, the
magnesium dissolves and hydrogen gas is produced.
Both magnesium metal and hydrogen gas are free
elements and so have an oxidation state of 0. In sulfuric
acid, hydrogen has an oxidation state of +1 since the
overall charge on the sulfate ion is −2. Similarly, the
oxidation state of magnesium in magnesium sulfate is +2.
Figure 3.18
Iron(
ii
) sulfate is pale green, whilst iron(
iii
) sulfate is yellow.
•
An increase in the oxidation state, for example from
+2 to +3 as in the case of Fe
2+
to Fe
3+
, is oxidation.
•
However, a reduction in the oxidation
state, for example from +6 to +3 as in the case
of Cr
6+
(in CrO
4
2−
) to Cr
3+
, is reduction.
During a redox reaction, the substance that brings
about oxidation is called an
oxidising agent
and is
itself reduced during the process. A substance that
brings about reduction is a
reducing agent
and is
itself oxidised during the process (see pp. 14, 39).
For example, if a dilute solution of acidifi ed
potassium manganate(
vii
) (pale purple) is added to
a solution of iron(ii)
ii
) sulfate, a colour change takes
place as the reaction occurs (Figure 3.19). The
iron(ii)
ii
) sulfate (pale green) changes to pale yellow,
showing the presence of iron(
iii
) ions.
magnesium + sulfuric → magnesium + hydrogen
acid sulfate
Mg(
s
) + H
2
SO
4
(
aq
) → MgSO
4
(
aq
) + H
2
(
g
)
Oxidation 0
+1
+2
0
states
The sulfate ions are unchanged by the reaction and
so can be ignored.
As you can see, the oxidation state of magnesium
has increased from 0 to +2. Therefore the
magnesium has been oxidised by the sulfuric acid
and so sulfuric acid is the oxidising agent. The
oxidation state of hydrogen has decreased from +1 in
the sulfuric acid to 0 in the free element. Therefore
the hydrogen has been reduced by the magnesium
and so magnesium is the reducing agent.
Question
1
Identify the oxidising and reducing agents in the
following reactions.
a
Zn(
s
)
+
H
2
SO
4
(
aq
)
→
ZnSO
4
(
aq
)
+
H
2
(
g
)
b
Cu
2
+
(
aq
)
+
Zn(
s
)
→
Cu(
s
)
+
Zn
2
+
(
aq
)
c
Mg(
s
)
+
Cu(NO
3
)
2
(
aq
)
→
Mg(NO
3
)
2
(
aq
)
+
Cu(
s
)
d
Fe(
s
)
+
H
2
SO
4
(
aq
)
→
FeSO
4
(
aq
)
+
H
2
(
g
)
The 'cross-over' method
A less scientifi c but simpler method to work out
the formula of compounds is called the 'cross-over'
method. In this method it is only necessary to 'swap'
the valencies of the elements or groups of atoms
Figure 3.19
Manganate(
vii
) ions (oxidising agent) and iron(ii)
ii
) ions
(reducing agent) are involved in a redox reaction when mixed.
