Version 3.1
Chemistry NCEA L3
3.7 Redox
SJ Gaze
2013
Achievement Criteria
This achievement standard involves describing oxidation-reduction processes.
1 Processes include reactions in electrochemical and electrolytic cells, This includes the
use of reduction potentials and spontaneity of reactions. Knowledge of preferential
discharge in electrolytic cells is not required.
2 Calculations are limited to include the use of oxidation numbers, mole ratios and
electrode potentials.
3 Knowledge of appearance of the following reactants and their products is required.
Oxidants will be limited to: O2, Cl2, I2, Fe3+, dilute acid (with metals), H2O2, MnO4– (reacting
in acidic, basic or neutral conditions), Cu2+ , Cr2O72– / H+, OCl–, concentrated HNO3, IO3–
, MnO2.
Reductants will be limited to: metals, C, H2, Fe2+, Br–, I–, H2S, SO2, (SO32– , HSO3-), S2O32–,
H2O2, H2C2O4.
Appropriate information relating to other oxidants or reductants will be provided.
Standard reduction potentials will be included where required.
Redox Terms
 A redox reaction is where one substance is oxidised and the other
substance is reduced.
Oxidation
>loss of electrons
>loss of hydrogen
>gain of oxygen
Reduction
>gain of electrons
>gain of hydrogen
>loss of oxygen
Oxidation numbers are used to determine what is oxidised and what is reduced in
a reaction.
Electron transfer
An Iron nail left in copper sulfate
Copper is reduced – gained electrons
Oxidising agent (oxidant)
2+
Fe(s) + Cu(aq)
2+
Fe(aq) +
Iron is oxidised – lost electrons
Reducing Agent (reductant)
Cu(s)
Oxygen transfer
Iron Ore smelting
Iron oxide is reduced – lost oxygen
Oxidising agent (oxidant)
2Fe2O3(s) + 3C(s)
4Fe(s)
+
3CO2(g)
carbon is oxidised – gained oxygen
Reducing Agent (reductant)
Iron Ore
Hydrogen transfer
Sulfur production
Hydrogen sulphide is oxidised – lost hydrogen
Reducing Agent (reductant)
2H2S(g) + O2(g)
2S(s)
Oxygen gas is reduced – gained hydrogen
Oxidising agent (oxidant)
+
2H2O(l)
Half Equations
A balanced redox equation is broken into two half-equations, to show how
electrons are transferred.
Fe(s)
+
2+
2+
Cu(aq)
Fe(aq)
Cu(s)
+
Reduction half equation - oxidant is reduced to a product
2+
Fe
Fe
+
2e
-
Oxidation half equation – reductant is oxidised to a product
Cu
2+
+
2e -
Cu
Summary of terms
Reductant
Oxidant
>reducing agent
>is oxidised
>loses electrons + hydrogen
>gains oxygen
>oxidising agent
>is reduced
>gains electrons + hydrogen
>loses oxygen
Oxidants – L3
Name
Conditions
permanganate
(manganate(VII))
acidified
MnO4- + 8H+ + 5epurple (+7)
→
Mn2+ + 4H2O
colourless (+2)
manganese (II)
ion
permanganate
(manganate(VII))
neutral
MnO4- + 4H+ + 3epurple (+7)
→
MnO2 + 2H2O
brown (s) (+4)
manganese
dioxide
permanganate
(manganate(VII))
alkaline
MnO4- + epurple (+7)
→
MnO42green (+6)
manganate
Dichromate
(dichromate(VI))
acidified
Cr2O72- + 14H+ + 6eorange (+6)
→
2Cr3+ + 7H2O
green (blue) (+3)
chromic (III) ion
Hydrogen
peroxide
acidified
H2O2 + 2H+ + 2ecolourless (-1)
→
2H2O
colourless (-2)
water
Cl2 + 2epale yellow/green (0)
→
Clcolourless (-1)
chloride ion
ClO- + H2O + 2ecolourless (+1)
→
Cl- + 2OHcolourless (-1)
chloride ion
I2 + 2egrey (s) (0)
→
2Icolourless (aq)(-1)
iodide
ion
→
I- + I2
colourless (aq)(-1)
iodide
ion
chlorine
Hypochlorite
(chlorate (I))
iodine
triiodide
alkaline
Half equation / colour change /ON
I3Purple brown (-1)
Name
Oxidants – L3
Name
Conditions
Half equation / colour change /ON
chlorate
(chlorate(V))
acidified
ClO3- + 6H+ + 6ecolourless (+5)
→
Nitric acid
(nitrate ion)
concentrated
NO3- + 2H+ + ecolourless (+5)
→
Cl- + 3H2O
colourless (-1)
NO2 + H2O
brown (g) (+4)
Name
chloride
ion
nitrogen
dioxide
Iron (III) ion
Fe3+ + eorange (+3)
→
Fe2+
pale green (+3)
Copper (II) ion
Cu2+ + eblue (+2)
→
Cu+
white (+1)
oxygen
O2 + 4ecolourless (0)
→
2O2colourless (-2)
oxide ion
Dilute acid
(H+ ions)
2H+ +2eColourless (+1)
→
H2
colourless (0)
hydrogen gas
Mn2+ + 2H2O
colourless (+2)
manganese
ion
I2
grey (s)(0)
iodine
Manganese
dioxide
acidified
MnO2 + 2H+ + 2ecolourless (+4)
→
iodate
IO3- + 2eColourless (aq) (+5)
→
bromate
BrO3- + 2eColourless (aq) (+5)
→
Brcolourless (-1)
Iron (II) ions
copper (I) ion
bromide ion
Reductants – L3
Name
Half equation / colour change /ON
hydrogen
H2(g)
colourless (0)
→
2H+(aq)
+ 2ecolourless (+1)
hydrogen ion
zinc
Zn(s)
grey (0)
→
Zn2+
+ 2ecolourless (+2)
zinc (II) ion
Iron (II) ion
Fe2+(aq)
pale green (+2)
→
Fe3+(aq)
+
orange (+3)
oxalate ion
(ethanediotae (COO-)2)
C2O42-(aq)
colourless (+3)
→
2CO2(aq or g) + 2ecolourless (+4)
oxalic acid
(ethanedioic acid)
H2C2O4(aq)
colourless (+3)
sulfite ions
SO32-(aq) + H2O
colourless (+4)
→
SO42-(aq)
+ 2H+ + 2ecolourless (+6)
sulphate ion
thiosulfate
2S2O32-(aq)
colourless (+2)
→
S4O62-aq)
+ 2ecolourless solution (+2.5)
tetrathionate
magnesium
Mg(s)
grey (0)
carbon
C(s) + 2H2O
black (0)
→
→
→
Name
e-
2CO2(aq or g) + 2H+ + 2ecolourless (+4)
Mg2+(aq)
+ 2ecolourless (+2)
CO2(g)
+ 4H+ + 4ecolourless (+4)
Iron (III) ion
carbon
dioxide
carbon
dioxide
magnesium
ion
carbon
dioxide
Reductants – L3
Name
Half equation / colour change /ON
carbon monoxide
CO(g) +
H2O
colourless (+2)
iron metal
Fe(s)
silver (0)
copper metal
Cu(s)
Red/brown (0)
→
hydrogen sulphide
H2S(g)
colourless (-2)
→
S(s)
+ 2H+ + 2eyellow ppt (0)
sulphur
hydrogen peroxide
H2O2(aq)
colourless (-1)
→
O2( g)
+ 2H+ + 2ecolourless (0)
oxygen gas
iodide ion
2I-(aq)
colourless (-1)
→
I2(s)
grey (0)
bromide ion
2Br-(aq)
colourless (-1)
→
Br2(aq)
+ 2eorange (0)
bromine
Sulfur dioxide
SO2(g) + 2H2O
colourless (+4)
→
SO42-(aq)
+ 4H+ + 2ecolourless (+6)
sulphate ion
carbon
C(s)
+ H2O
black (0)
→
→
Name
CO2(g)
+ 4H+ +
colourless (+4)
→
2e-
carbon
dioxide
Fe2+
+ 2epale green (+2)
iron (II) ion
Cu2+(aq)
blue (+2)
copper (II) ion
+
2e-
+ 2e-
CO(g)
+ 2H+ + 2ecolourless (+2)
iodine
carbon
monoxide
Oxidation Numbers
 The Oxidation Number (ON) gives the ‘degree’ of oxidation or reduction of an
element.
 They are assigned to a INDIVIDUAL atom using the following rules.
Elements
=0
e.g.
ON
Hydrogen atom
Oxygen atom
(not as element)
(not as element)
ON
Fe
H2
0
0
= +1
e.g.
HCl H2SO4
+1
ON
= -2
e.g.
MnO4- CO2
+1
-2
Except peroxides
Except Hydrides
ON
ON
= -1
e.g.
LiH
-1
-2
= -1
H2O2
-1
Oxidation Numbers
•
•
The Oxidation Number (ON) gives the ‘degree’ of oxidation or reduction of
an element.
They are assigned to a INDIVIDUAL atom using the following rules.
Elements
Hydrogen atom
Oxygen atom
Each atom
(not as element)
(not as element)
ON
=0
e.g.
Fe
ON
H2
= +1
e.g.
HCl H2SO4
Except Hydrides
= -1
ON
e.g.
LiH
ON
= -2
e.g.
MnO4- CO2
ON
Except peroxides
= -1
H 2 O2
Oxidation Numbers
Monatomic ions
ON
= charge
e.g
2+
Fe
Cl
+2
Polyatomic ions
Sum of ON
=charge
Molecules
Sum of ON on atoms =charge
e.g.
CO2
e.g.
MnO4
-1
+7 -2
Because
Total charge = -1
And
Oxygen = -2
+7 + (4x -2) = -1
Mn
O
+4 -2
Because
Total charge = 0
And
Oxygen = -2
+4 +(2x-2) = 0
c
O
Oxidation Numbers
Monatomic ions
ON
= charge
e.g
Fe 3+
Cl-
Polyatomic ions
Sum of ON
=charge
Molecules
Sum of
Because
Total charge = -1
And
Oxygen = -2
+7 + (4x -2) = -1
Mn
O
=zero
e.g.
CO2
e.g.
MnO4+7 -2
ON
+4 -2
=0
Because
Total charge = 0
And
Oxygen = -2
+4 +(2x-2) = 0
C
O
Oxidation Numbers
Oxidation is a loss of electrons
Reduction is a gain of electrons
ON
and causes an increase in
Oxidation of Fe 2+
Fe 2+
+2
and causes an decrease in
ON
Reduction of MnO4 Fe 3+ + e-
Mn 2+
+7
+3
Fe has increased
MnO4 -+ 5e-
ON
(+2 to +3) caused by
a loss of electrons e-
+2
Mn has decreased ON
(+7 to +2) caused by
a gain in electrons e-
OXIDATION and REDUCTION always occur together. The electrons lost by
one atom are gained by another atom.
This is called a REDOX reaction.
Join the two half equations together.
Using ON to identify Redox Reactions
4Fe + 3CO2
2Fe2O3 + 3C
What has been oxidised and what has been reduced?
STEP ONE – write the ON for each atom
2Fe2O3 + 3C
+3 -2
0
4Fe + 3CO2
0
+4 -2
STEP TWO – Identify the atom that has had its ON increased. It is Oxidised
C has increased ON (0 to +4) so C is Oxidised.
STEP THREE – Identify the atom that has decreased ON. It is reduced.
Fe has decreased ON (+3 to 0) so Fe is Reduced.
Balancing half equations
1. Write half
equation by
identifying
reactant and
product
2. Balance atoms
that are not O or
H
3. Balance O by
4. Balance charge
adding H2O and H by adding
by adding H +
electrons
Balance the half equation for the oxidation of Fe 2+ to Fe 3+
Fe
2+
Fe
3+
Atoms already
balanced
Fe
2+
There are no O or Fe 2+ Fe +3+e
H atoms to
2+ = 3+ (-1)
balance
1 electron to
balance charge
3+
Fe
+ e
Balancing half equations
1. Write half
equation by
identifying
reactant and
product
2. Balance atoms
that are not O or
H
3. Balance O by
4. Balance charge
adding H2O and H by adding
by adding H +
electrons
Balance the half equation for the reduction of MnO4 - to Mn
MnO4 -
Mn 2+
Atoms already
balanced
MnO4 -+ 8H
2+
MnO4 -+ 8H
+
+
Total charge +7
2+
Mn + 4H2O
2+
Mn + 4H2O
Balance O by adding
4H2O and H by adding
8 H+
+
MnO4 + 8H
+
5e-
Total charge +2
Add 5 electrons (e-)
2+
Mn + 4H2O
Balancing half equations
Rules e.g.
Cr2O72→
Cr3+
1. Assign oxidation numbers and identify element oxidised or reduced.
(+6)(-2)
(+3)
Cr2O72→
Cr3+
2. Balance atom no. for element oxidised or reduced (other than oxygen and hydrogen)
Cr2O72→
2Cr3+
3. Balance the Oxygen using H2O
Cr2O72-
→
2Cr3+ + 7H2O
4. Use H+ (acidic conditions) to balance the hydrogen
14H+ + Cr2O72- + 6e→
2Cr3+ + 7H2O
5. Use OH- (in alkaline conditions) to cancel any H+ [same amount on both sides]
6. Balance charge by adding electrons (LHS on oxidants RHS on reductants)
14H+ + Cr2O72- + 6e→
2Cr3+ + 7H2O
7. Check balance of elements and charges.
Electrochemistry
Electrochemistry is the chemistry of reactions involving the transfer of electrons,
redox reactions. In year 12 the focus was on electrolytic cells or electrolysis.
This involves a non-spontaneous reaction in which an external source of
electricity provides electrons with the energy required to bring about a redox
reaction.
In year 13 the focus is on
electrochemical cells in
which spontaneous redox
reactions use the energy
released from a chemical
reaction to generate
electric current. These are
called Galvanic cells or
batteries.
Galvanic Cells and Salt Bridges
Under normal conditions a redox reaction occurs
when an oxidising agent is in contact with a reducing
agent. If the two half reactions are physically
separated, the transfer of electrons is forced to take
place through an external metal wire. As the
reaction progresses a flow of electrons occurs. This
only happens if there is a full circuit so that there is
no net build-up of charge. To complete this circuit
the separate solutions are connected using a salt
bridge which allows ions to flow and transfer charge.
Typically the salt bridge is a glass tube filled with a
gel prepared using a strong electrolyte such as
KNO3(aq) (which contains ions that do not react with
the electrodes or species in the solutions. The
anions (NO3-) and cations (K+) can move through the
salt bridge so that charge does not build up in either
cell as the redox reaction proceeds.
Galvanic cells
The oxidation and reduction reactions that occur at the electrodes are called half-cell
reactions.
Zn electrode (anode, oxidation)
Zn(s) 
Zn2+(aq) + 2e
Cu electrode (cathode, reduction) Cu2+(aq) + 2e

Cu(s)
Electromotive Force
The reduced and oxidised substances
in each cell form a redox couple. The
2 couples in this cell (the Daniel cell)
are Zn2+|Zn and Cu2+|Cu. By
convention, when writing redox
couples, the oxidised form is always
written first.
The fact that electrons flow from one
electrode to the other indicates that
there is a voltage difference between
the two electrodes. This voltage
difference is called the electromotive
force or emf of the cell and can be
measured by connecting a voltmeter
between the two electrodes. The emf
is therefore measured in volts and is
referred to as the cell voltage or cell
potential.
A high cell potential shows that the cell
reaction has a high tendency to generate a
current of electrons. Obviously the size of
this voltage depends on the particular
solutions and electrodes used, but it also
depends on the concentration of ions and
the temperature at which the cell operates.
Cell Diagrams
Galvanic cells can be represented using cell diagrams. This is a type of
short hand notation that follows a standard IUPAC convention. For the
copper/zinc cell the standard cell diagram is
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
The vertical lines represent phase boundaries and || represents the salt
bridge.
The cathode (reduction reaction) is always shown on the right hand side
and the anode (oxidation) on the left in a standard cell diagram.
The electrons thus move from left to right in the standard cell diagram,
representing a spontaneous redox reaction. The electrodes are always
written in at the beginning and end of a cell diagram. This occurs both if
the metal is involved in the redox reaction (as in the Daniel cell above
where the electrodes are the Cu and Zn), and also if an inert electrode is
used.
In each half cell the reactant appears first, followed by the product.
Cell Diagrams
An inert electrode must be used in cells in which both species in a redox
couple are in aqueous solution (MnO4- and Mn2+). The inert electrodes
are commonly either platinum, Pt(s) or graphite, C(s) electrodes. Since
the two species in the redox couple are in solution, they are separated by
a comma rather than a vertical line.
eg
Cu(s) | Cu2+(aq) || MnO4(aq), Mn2+(aq) | Pt(s)
The cell diagram shows two half cells linked. Each half cell consists of the
oxidant, the reductant and the electrode (which may be the oxidant or
reductant). The two half cells above are Cu(s)|Cu2+(aq) and MnO4(aq),
Mn2+(aq)|Pt(s).
Standard Electrode Potentials
Under standard conditions (when the pressure of hydrogen gas is 1 atm, and
the concentration of acid is 1 mol L-1) the potential for the reduction reaction is
assigned a value of zero.
2H+(aq) + 2e →
H2(g)
Eo = 0.00 V
The superscript o denotes standard state conditions. When the hydrogen
electrode acts as a cathode, H+ ions are reduced, whereas when it acts as an
anode, H2 gas is oxidised.
Standard Electrode Potential
The overall cell voltage is the sum
of the electric potential at each
electrode. If one of the electrode
potentials is known, and the overall
cell voltage is measured, then the
potential of the other electrode can
be calculated by subtraction.
Clearly it is best if all electrode
potentials are measured relative to
a particular electrode. In this way,
a scale of relative values can be
established. The standard
hydrogen electrode (SHE) is used
as the standard reference
electrode, and it has arbitrarily
been given a value of 0.00 V.
Standard Electrode Reduction Potential
For any redox couple, the standard electrode (reduction) potential is the
voltage obtained under standard conditions when that half-cell is connected
to the standard hydrogen electrode.
For example, the electrode potential of a Zn2+|Zn electrode can be measured
by connecting it to a hydrogen electrode.
Experimentally, the more positive
terminal is always where reduction is
occurring in a spontaneous reaction. In
example (a) reduction occurs in the
hydrogen electrode (positive electrode)
while oxidation occurs in the Zn2+|Zn
compartment (negative electrode). The
cell diagram for this electrochemical cell is
Zn(s) | Zn2+(aq) || H+(aq), H2(g) | Pt(s)
Standard Reduction Potentials
Using the standard reduction potentials for many half reactions have been
measured under standard conditions (at 25 oC). Standard reduction
potentials are provided in examinations.
The table can be used to decide the relative strength of species as oxidants or
reductants. The species on the left in the couple with the most positive
reduction potential, will be the strongest oxidising agent or oxidant. E.g it is
F2(g) (NOT F2 / F). This means F2 has the greatest tendency to gain electrons.
As the electrode potential decreases, the strength as an oxidant decreases.
Conversely the strongest reducing agent or reductant would have the least
positive (or most negative) e.g. Li(s). This means Li has the greatest tendency
to lose electrons.
Using reduction potentials to determine Eocell
In any electrochemical cell, the standard cell potential (voltage), E0cell, is the
difference between the reduction potentials of the two redox couples
involved. The couple with the more positive reduction potential will be the
reduction half-cell (cathode). This means that the Eocell for any combination
of electrodes can be predicted using the relationship
Eocell = Eo(reduction half-cell) - Eo(oxidation half-cell)
OR
Eocell = Eo(cathode) - Eo(anode)
OR
Eocell = Eo(RHE) - Eo(LHE)
(where RHE is the right hand electrode and LHE is the left hand electrode in
the standard cell diagram).
Using reduction potentials to determine Eocell
For example, consider the cell Zn(s) | Zn2+(aq) || Ag+(aq) | Ag(s)
Reduction reaction is Ag+(aq) + e 
Ag(s)
Oxidation reaction is Zn(s) 
Zn2+(aq) + 2e
Eocell = Eo(Ag+/Ag) - Eo (Zn2+/Zn)
Eo(Ag+/Ag) = +0.80V
Eo (Zn2+/Zn) = -0.76V
= 0.80 - (-0.76) V = +1.56V
1. Any metal that is more reactive (lower Eo value) will reduce the cation of a less
reactive metal because the Eocell value for the reaction will be positive.
2. A positive standard cell potential suggests that the reaction occurs
spontaneously from left to right, as shown in a standard cell diagram. In reality the
reaction may not appear to proceed as it may be very slow due to a high activation
energy. Another reason a reaction may not proceed is if the surface of the metal is
coated with an oxide. Aluminium oxide on the surface of aluminium can protect
the aluminium from undergoing a spontaneous oxidation reaction
Predicting whether a Reaction will occur
It is possible to use Eo values to predict whether a reaction will occur. This
simply involves identifying which species must be reduced and which
species must be oxidised if the reaction is to proceed spontaneously. The
appropriate reduction potentials are then substituted into the equation.
Eocell = Eo(cathode) - Eo(anode)
where Eo(cathode) is the reduction potential for the half cell where reduction
occurs and Eo(anode) is the reduction potential for the half cell where
oxidation occurs. If the Eocell calculated is positive, then the reaction will
occur spontaneously. Conversely, a negative cell potential means the
reaction will not proceed.
Predicting whether a Reaction will occur
Corrosion – an everyday application of a redox reaction
Corrosion is the term usually applied to the deterioration of
metals by an electrochemical process. One example is the rusting of iron
in the presence of water and oxygen.
Although the reactions involved in rusting are quite complex, the
main steps are as follows.
Step 1 - Oxidation occurs at a region of the iron’s surface, the
anode.
Fe(s)
Fe2+(aq) + 2e
Eo red = -0.44 V
Step 2 - Electrons travel to some other region of the metal’s
surface where a variety of cathode reactions can occur.
Step 3 - In acidic medium, atmospheric oxygen is quickly reduced
to H2O. The acidity of the solution can, in part, be due to the presence of
dissolved acidic gases such as CO2 and SO2.
Predicting whether a Reaction will occur
Example
Can a solution of acidified potassium permanganate oxidise the Fe2+
present in a solution of iron (II) nitrate? (Note in questions such as this you
will have to recognise that ions such as sodium and potassium are
spectator ions.) The unbalanced equation for the reaction would be:
MnO4
+ Fe2+
Reduction reaction is

Mn2+ + Fe3+
MnO4  Mn2+
Eo (MnO4 /Mn2+) = +1.51 V
Oxidation reaction is Fe2+
 Fe3+
Eo cell
= Eo (MnO4 /Mn2+) - Eo (Fe3+/Fe2+)
Eo (Fe3+/Fe2+) = +0.77 V
= +1.51 - +0.77 = 0.74 V
Since the cell potential is a positive value (>0.00 V) the reaction should
proceed. If it had been negative then the reaction would not proceed.
Discussion format of E° Cells
1.
Movement of electrons
from anode (oxidation > LEO) lower E° value
to cathode (reduction > GER) higher E° value
2. Half equations at each electrode
anode X- → X + ecathode X + e- → X-
3. Movement of salt bridge solutions
Anions → Anode
Cations → Cathode
4. Overall summary
Oxidation at the anode → salt bridge anions balance
cations
produced at the electrodes
Reduction at the cathode → salt bridge cations balance cations
removed at electrode
E° Cells Summary
AAO
n
o
d
e
n
i
o
n
LEO
x
i
d
a
t
i
o
n
CCR
a
t
h
o
d
e
GER
a
t
i
o
n
e
d
u
c
t
i
o
n