Oxydation-Reduction Reactions

Balancing Oxydation-Reduction Reactions

Voltaic Cells
Cell EMF
Spontaneity of Redox Reactions
Effect of Concentration of Cell EMF
Batteries
Corrosion
Electrolysis

Oxidationreactions

the
Oxidation-reduction
transfer of electrons from one species to
another.
The oxidation state of one or more
substances in the reaction changes.
The transfer of electrons can be used to
produce energy in the form of electricity.
Electrochemistry is the study of the
relationships between chemical reactions
and electrical energy.

We identify a reaction as oxidationreduction by comparing the oxidation

numbers of atoms in the reactants and
products.
If the oxidation numbers change, the
reaction is an oxidation-reduction.

zinc metal loses electrons to become a

cation, and thus zinc has been oxidized.
Hydrogen gains electrons to become
hydrogen gas.
Hydrogen has been reduced.
Zinc is the reducing agent,
agent or reductant,
and hydrogen ion is the oxidizing agent,
agent or
oxidant.
oxidant

The amount of each element must be the

same on both sides of the equation.
Balancing
the
number
of
electrons
transferred. In the equation

Balancing the mass automatically balances

the total charge on each side.
In many reactions, though, balancing the
mass does not result in balancing charge.

In the next equation the law of conservation

of mass appears to have been obeyed.
However, note that the total charge on the
left side is +3, while the total charge on the
right side is only +2.

The two processes, oxidation of manganese

metal and reduction of chromium(III) ion, do
not correspond to the transfer of the same
number of electrons.
If we multiply the manganese on each side by
3 and the chromium on each side by 2, we
get

It is not always convenient to balance oxidationreduction reactions by inspection.

The equation

would be impossible to balance by inspection.

To balance such an equation, we use a technique
known as the method of halfhalf-reactions.
reactions
A half-reaction is an equation that shows either
oxidation or reduction alone.

The electrical energy produced by a spontaneous

oxidation-reduction reaction can be "harnessed"
using a voltaic cell (also called a galvanic cell).
In a voltaic cell the two half-reactions are made
to occur in separate compartments (half-cells).
The electrons can be transferred only through
the wire.
oxidation and reduction half-reactions occur at separate
electrodes
electric current flows through the wire

Anode - the electrode at which oxidation

takes place.
the negative (-) electrode
produces electrons

Cathode - the electrode at which reduction

takes place.
the positive (+) electrode
consumes electrons

Oxidation-reduction reactions part I

Galvanic cells I: the copper-zinc cell

Galvanic cells II: the zinc-hydrogen cell

EXAMPLE:
Describe how you would construct a galvanic
cell based on the following reaction:
Pb2+ (aq) + Zn (s) Pb (s) + Zn2+ (aq)

Pb2+ (aq) + 2 e- Pb (s)

Zn (s) Zn2+ (aq) + 2 eLooking at the two half-reactions, we find that
the Pb2+ is being reduced, and the Zn is being
oxidized.
Therefore, the anode compartment of our cell
would consist of a strip of zinc metal immersed
in a solution containing Zn2+ ions (such as zinc
nitrate).
The cathode compartment would consist of a
strip of lead immersed in a solution containing
Pb2+ ions (such as lead (II) nitrate).

Pb2+ (aq) + 2 e- Pb (s)

Zn (s) Zn2+ (aq) + 2 eThe two half-cells would be connected to
each other with a salt bridge and an external
wire.
Electrons flow through the wire from the zinc
anode to the lead cathode.
Anions move from the cathode compartment
towards the anode while cations migrate from
the anode compartment toward the cathode.

Redox chemistry of iron and copper

Single vertical line, , represents a phase

boundary.
Double vertical line, , represents a salt
bridge.

Shorthand for the anode half-cell is always

written on the left of the salt-bridge symbol,
followed on the right of the symbol by the
shorthand for the cathode half-cell.
Reactants in each half cell are written first, followed
by products.
Electrons move through the external circuit from
left to right.
For Zn (s) + Cu2+ (aq) Zn2+ (aq) + Cu (s):
Zn (s)Zn2+ (aq) Cu2+ (aq)Cu (s).

Cell involving a gas.

Additional vertical line due to presence of
additional phase.
List the gas immediately adjacent to the appropriate
electrode.

Detailed notation includes ion concentrations

and gas pressures.

EXAMPLE:
Give the shorthand notation for a galvanic cell
that employs the overall reaction
Pb(NO3)2(aq) + Ni (s) Pb(s) + Ni(NO3)2(aq)
Give a brief description of the cell.

SOLUTION: The two half-reactions for this

overall reaction are:
Pb2+ (aq) + 2 e- Pb (s)
Ni (s) Ni2+ (aq) + 2 eFrom these half-reactions, we know that lead
is being reduced and nickel is being oxidized.

Therefore, Ni is the anode and Pb is the

cathode. The cell notation is:
Ni(s)Ni2+(aq)Pb2+(aq)Pb(s)
This cell would consist of a strip of nickel as
the anode dipping into an aqueous solution
of Ni(NO3)2 and a strip of Pb as the cathode
dipping into an aqueous solution of Pb(NO3)2.
The two half-cells would be connected by a
salt bridge and a wire.

Electrons spontaneously flow from one

species to anotherand through the wire
from the anode to the cathode of a voltaic
cellbecause of a difference in potential
energy, or a potential difference.
The potential difference between two
electrodes is measured in volts.
One volt is equal to one joule per coulomb. 1

The potential difference

that drives electrons
through the wire in a voltaic
cell is called the
electromotive force or
emf.
For a voltaic cell the emf is
denoted Ecell and referred
to as the cell potential.

Cell potential - measured with a voltmeter.

Gives a positive reading when the + and - terminals
of the voltmeter are connected to cathode (+) and
anode (-), respectively.
can use voltmeter-cell connections to determine which
electrode is the anode and which is the cathode

The value of a cell potential depends on what

half-reactions are taking place in the two
compartments of the cell.
The cell potential measured under standard
conditions, Ecell (25C, 1 M concentrations,
and 1 atm pressures), is the standard cell
potential or standard emf.
emf

For the zinc and copper voltaic cell in Figure

20.5, Ecell is 1.10 V.
That is for the reaction

at 25C, where the concentrations of copper

and zinc ions are both 1 M.

Such
potentials
can
be
measured
experimentally, but many of them can be
calculated from tabulated standard reduction
potentials,
potentials Ered values.

The more positive the standard cell potential,

the greater the driving force for electrons to
flow from the anode to the cathode.
Because the cathode of a voltaic cell is always
the half-reaction with the more positive (or
less negative) standard reduction potential,
the standard cell potential of a voltaic cell is
always positive.

The standard reduction potentials for the various

half-reactions are measured against a standard
hydrogen electrode (SHE).
The half-reaction of interest and the SHE, both
under standard conditions, are made into a
voltaic cell, as shown in Figure 20.11, and the
cell potential is measured experimentally.
The standard potential of the standard hydrogen
electrode's half-reaction is arbitrarily assigned a
value of zero, so the measured potential
corresponds to the half-reaction being evaluated.

The standard reduction potentials in Table

20.1 can be used to compare the oxidizing
power or reducing power of a substance.
The more positive the value of Ered for a
species, the more readily it undergoes
reduction and the better oxidizing agent it is.
As Ered becomes more negative, the species
on the right side of the arrow becomes a
stronger reducing agent.

Standard reduction potentials

It is possible to use standard reduction

potentials to predict whether or not a given
oxidation-reduction
reaction
will
be
spontaneous.
For instance, will copper solid be oxidized by
a solution containing iron(II) ions?

To determine the answer, we calculate the

standard cell potential for the cell as we have
described it.
In our description, copper is being oxidized
so the copper electrode is the anode.
Iron is being reduced, making the iron
electrode the cathode.
Ered values for the two half-reactions are

Ered = 0.34 V

Ered = - 0.44 V

The negative Ecell tells us that this reaction, as

written, will not occur spontaneously. (Its reverse
reaction will be spontaneous.)

Our ability to predict the spontaneity of a

chemical reaction by calculating a standard
cell potential points to a relationship between
the sign of Ecell and the sign of G , the
standard change in Gibbs free energy.
Quantitatively, this relationship is expressed
as

n is the number of moles of electrons

transferred in the reaction,
F, called Faraday's constant, is the quantity of
electrical charge on a mole of electrons.

Both n and F are always positive numbers.

Therefore, a positive value for E will always
correspond to a negative value for G , both
denoting a spontaneous reaction.
The relationship between cell potential and
free energy change holds at conditions other
than standard, as well.

0.0592
E=E
log Q
n
o

Nernst equation:
0.0592
E=E
log Q
n
o

(in volts at 25oC).

Enables us to calculate cell

nonstandard-state conditions.

potentials

under

Calculate Ecell for the following cell reaction:

2 Cr(s) + 3Pb2+(aq) 2 Cr3+ (aq) + 3Pb(s)
[Pb2+] = 0.15 M;
[Cr3+] = 0.50 M

2 Cr (s) 2 Cr3+ + 6 e+0.74 V

3 Pb2+ (aq) + 6 e- 3 Pb (s)

Eo
Eo=-0.13V

o
E cell
= +0.74 + ( 013
. ) = +0.61 V

0.0592
E=E
log Q
n
o

[ ]
[ ]
3+ 2

Ecell =

Ecell

o
Ecell

Cr
0.0592

log
6
Pb 2 +

0.0592
(0.5)2
= +0.61
log
3 = 0.59
6
(0.15)

Three different ways to determine the value

of an equilibrium constant K:
[C]c[D]d
K=
[A]a [B]b
G o
ln K =
RT

nFE o
ln K =
RT

Equilibrium constants for redox reactions

tend to be either very large or very small in
comparison with equilibrium constants for
acid-base reactions.
Positive value of Eo corresponds to K > 1.
Negative value of Eo corresponds to K < 1.

Important application of Nernst equation electrochemical determination of pH using a

pH meter.
Consider a cell with a hydrogen electrode as
the anode and a second reference electrode
as the cathode.
Pt (s)H2 (1 atm)H+ (? M)reference cathode.

Ecell = 0.0592pH + Eref

Ecell = 0.0592pH + Eref

pH =

Ecell Eref
0.0592

can measure the pH of a solution by

measuring Ecell
Actual pH measurements use a glass
electrode with a calomel electrode as the
reference.

EXAMPLE::
EXAMPLE
The following cell has a potential of 0.49 V.
Calculate the pH of the solution in the anode
compartment.
Pt(s) H2(g) (1 atm)H+(pH = ?)Cl-(aq) (1M)Hg2Cl2
(s) Hg (l)

The cell reaction is

Hg2Cl2 (s) + H2 (g) 2 Hg (l) + 2 Cl- (aq) + 2
H+ (aq)

The cell reaction is

Hg2Cl2 (s) + H2 (g) 2 Hg (l) + 2 Cl- (aq) + 2
H+ (aq)
E o = E Ho

2 H

o
+ E Hg

2 Cl 2 Hg, Cl

= 0 .00 V + 0 .28 V = 0 .28 V

Ecell Eref
pH =
0.0592

0 . 49 V 0 . 28 V
pH =
= 3 .55
0 .0592

EXAMPLE::
EXAMPLE
Calculate the equilibrium constant for the
following reaction at 25oC.
5 S2O82- (aq) + I2 (s) + 6 H2O (l) 10 SO42- (aq) + 2 IO3- (aq) +
12 H+ (aq)

S2O82- (aq) + 2 e- 2 SO42- (aq)

Eo = +2.01
V
I2(s) + 6H2O (l) 2IO3- (aq) + 12 H+(aq) +10e- Eo
= -1.20 V
o
Ecell = 2.01+ (1.20 ) = +0.81 V

The value of n for this reaction is 10.

0.0592
E =
log K
n
o

(10)(0.81)
log K =
= 137
(0.0592)

K = 10137

Most important practical application of

galvanic cells is their use as batteries.
Features required in a battery depend on the
application.
C. General features.
Compact and lightweight.
Physically rugged and inexpensive.
Provide a stable source of power for relatively long
periods of time.

A battery is a self-contained source of

electrochemical energy made from one or
more voltaic cells.
Ordinary flashlight batteries consist of a
single voltaic cell, while car batteries are six
identical voltaic cells connected in series.
It is worth noting that electrochemistry is one
of only a very few commercially viable
methods of generating electricity.

Car Battery

Cathode:
Anhode:

Car Battery

Cathode:
Anhode:

Corrosion - the oxidative deterioration of a

metal.
Well-known
example
of
corrosion
conversion of iron to rust.
Requires both oxygen and water.
Involves pitting of the metal surface.
rust is deposited at a location physically separated
from the pits

Corrosion is the undesirable oxidation of a metal.

A familiar example of corrosion is the rusting of
iron.
Iron metal is oxidized to Fe2+ by oxygen.

The Fe2+ is then further oxidized to Fe3+ in a

hydrated form of Fe2O3, what we know as rust.

Corrosion of iron can be prevented by coating

it with paint or with another metal such as tin
or zinc.
Coating with paint or with a less readily
oxidized metal such as tin protects the iron
simply by preventing oxygen and water from
reaching the iron surface.
If such a coating is damaged and the iron is
exposed, corrosion of iron will occur in the
exposed area.

Some metals create their own sealant by

oxidation.
Aluminum, for example, oxidizes fairly
readily.
Ered = - 1.66 V

Coating iron with a more easily oxidized

metal such as zinc, also prevents oxygen and
water from reaching the iron surface.
But unlike tin, zinc protects the iron even if
the zinc coating is damaged.
It does so by making the iron serve as the
cathode in an voltaic cell.

When the zinc coating is damaged and iron is

exposed, the zinc itself is oxidized rather
than the iron.
Ered = - 0.44 V
Ered = - 0.76 V

Iron, with the more positive (less negative)

Ered is more easily reduced and therefore,
less easily oxidized.
Zinc will be oxidized, serving as the sacrificial
anode.
This type of corrosion prevention is called
cathodic protection.
protection
Ered = - 0.44 V
Ered = - 0.76 V

One of the many uses of cathodic protection is to

prevent corrosion of underground pipes and
storage tanksmany of which are iron.
A more easily oxidized metal is placed in electrical
contact with the object to be protected by an
insulated copper wire.
The iron object becomes the cathode (and the more
easily oxidized metal the sacrificial anode) in a
voltaic cell.
Oxidation eventually consumes the sacrificial
anode, and it must be replaced periodically.

A voltaic cell is one in which a spontaneous

chemical reaction is used to generate a
voltage.
Electrolysis is the use of a voltage to drive a
nonspontaneous reaction.
Reactions that are driven by an externally
supplied voltage are called electrolysis
reactions,
reactions and electrochemical cells designed
for the purpose of carrying out electrolysis
reactions are called electrolytic cells

Sodium metal is produced commercially by

electrolysis of molten sodium chloride.
Electrodes are immersed in molten sodium
chloride, and a voltage source drives
electrons from the anode to the cathode.
Sodium is reduced at the cathode to molten
sodium metal.
Chloride ions are oxidized to chlorine gas at
the anode.

Cathode:

Ered = - 2.71 V

Anode:

Ered = 1.36 V
Ecell = - 4.07 V

A negative cell potential means that

the reaction is nonspontaneous.
The value of the cell potential tells us
that a minimum of 4.07 volts must be
applied to the cell to drive the reaction
in the desired direction.

Cathode:

Ered = - 2.71 V

Anode:

Ered = 1.36 V
Ecell = - 4.07 V
Electrolysis of molten salts is used
commercially to produce a number of
active metals.

Cathode:

Ered = - 2.71 V

Anode:

Ered = 1.36 V
Ecell = - 4.07 V
Electrolysis of molten salts is used
commercially to produce a number of
active metals.

Electrolysis can also be done with aqueous

solutions.
Although chlorine gas can be produced from
a solution of sodium chloride, sodium metal
cannot be produced.
This has to do with the relative ease of
reduction of water versus sodium ions.
Ered = - 2.71 V
Ered = - 0.83 V

With a far less negative reduction potential,

waterrather than sodium will be reduced
at the cathode of an electrolytic cell
containing aqueous sodium chloride.

Ered = - 2.71 V
Ered = - 0.83 V

Ered = - 2.71 V
Ered = - 0.83 V

In a comparison of reduction potentials, we

might also predict that water, and not
chloride ions, would be oxidized at the anode
of such an electrolytic cell.

Ered = 1.36 V
Ered = 1.23 V

The first half-reaction, with the more positive

reduction potential, is more apt to occur as a
reduction rather than an oxidation.
Experimentally, however, chlorine gas is
produced at the anode during electrolysis of
aqueous sodium chloride.
Ered = 1.36 V
Ered = 1.23 V

This is explained by kinetics.

Although the oxidation of water is
thermodynamically favored, the oxidation of
chloride ions is kinetically favored because of
a lower activation energy.

Ered = 1.36 V
Ered = 1.23 V

Electrolysis of water

Electroplating is an electrolytic process used

to deposit a thin layer of one metal on
another.
In such a process the metal to be deposited is
used as the anode, and the metal on which
the deposit is to be made is the cathode.
Electroplating is the process by which silverplated flatware is finished.
It is also the process used to coat iron car
bodies with zinc to protect them from rust.

When the electrode in an electrolytic cell is

involved in the reaction, it is called an active
electrode.
When the same metal is to be oxidized from
the anode and reduced (deposited) at the
cathode, the standard cell potential is zero.
It therefore requires only a very small voltage
to drive such a process.

Electroplating

Using stoichiometry, we can relate the

amount
of
metal
deposited
in
an
electroplating process to the current and the
length of time for which the current is
applied.

The charge passing through an electrolytic

cell is measured in coulombs.

How many grams of copper would be reduced

by the application of 5.00 amps to a solution
of copper sulfate for 35.0 minutes?

EXAMPLE::
EXAMPLE
How many grams of Cl2 would be produced in
the electrolysis of molten NaCl by a current of
4.25 A for 35.0 min?

Remember that a coulomb is an A.s or that an

ampere is C/s.
2 Cl- Cl2 + 2 emoles of electrons = 2

C
60 s
1 mol e - 1 mol Cl 2 70.9 g Cl 2
4.25
35.0 min

-
- = 3.28 g Cl 2
s
1 min 96,500 C 2 mol e
1 mol Cl