Electrochemical Reactions

In electrochemical reactions, electrons are
transferred from one species to another.

Substance which receives electrons is the
oxidising agent or oxidant and is itself reduced.

Substance which gives up its electrons is the
reducing agent or reductant and is itself
oxidised.
Oxidation and Reduction
Gain of Electrons is Reduction
GER
LEO
Loss of Electrons is Oxidation
Balancing Redox Equations (Chem 120)
Reduction Potentials
Reduction Potentials
Voltaic Cells

In spontaneous oxidation-reduction (redox)
reactions, electrons are transferred and energy is
released.
Figure 16.2
Voltaic Cells

We can use that
energy to do work if
we make the
electrons flow
through an external
device.

We call such a
setup a voltaic cell.
Figure 16.3
Voltaic Cells

A typical cell looks
like this.

The oxidation
occurs at the
anode.

The reduction
occurs at the
cathode.
Voltaic Cells

Figure 16.4
Once even one
electron flows
from the anode to
the cathode, the
charges in each
beaker would not
be balanced and
the flow of
electrons would
stop.
Voltaic Cells

Figure 16.4
Therefore, we use
a salt bridge,
usually a U-shaped
tube that contains a
salt solution, to
keep the charges
balanced:

Cations move toward
the cathode.

Anions move toward
the anode.
Voltaic Cells

In the cell, electrons leave the anode and
flow through the wire to the cathode.

As the electrons leave the anode, the
cations formed dissolve into the solution
in the anode compartment.
Voltaic Cells

As the electrons reach the cathode,
cations in the cathode are attracted to the
now negative cathode.

The electrons are taken by the cation,
and the neutral metal is deposited on the
cathode.
Electromotive Force (emf)
Figure 16.6

Water only
spontaneously flows
one way in a
waterfall.

Likewise, electrons
only spontaneously
flow one way in a
redox reaction—
from higher to lower
potential energy.
Electromotive Force (emf)

The potential difference between the
anode and cathode in a cell is called the
electromotive force (emf).

It is measured in volts (V):
J
1V=1 C

It is also called the cell potential, and is
designated Ecell.
Standard Cell Potentials

The cell potential at standard conditions
can be found through this equation:
 = Ered
 (cathode) − Ered
 (anode)
Ecell

Because cell potential is based on the
potential energy per unit of charge, it is
an intensive property.
Standard Hydrogen Electrode

Their values are referenced to a standard
hydrogen electrode (SHE).

By definition, the reduction potential for
hydrogen is 0 V.
2 H+ (aq, 1M) + 2 e−  H2 (g, 1 atm)
Figure 16.7
Standard Reduction (Half-Cell)
Potentials

Table 16.1
Reduction
potentials for
many
electrodes
have been
measured and
tabulated.
Cell Potentials

For the oxidation in this cell:
 = −0.76 V
Ered

For the reduction:
Ered
 = +0.34 V
Figure 16.8
Cell Potentials
 = Ered
 (cathode) − Ered
 (anode)
Ecell
= +0.34 V − (−0.76 V)
= +1.10 V
Strengths of Oxidising
and Reducing Agents
Figure 16.11

The strongest oxidisers
have the most positive
reduction potentials.

The strongest reducers
have the most negative
reduction potentials.
Strengths of Oxidising
and Reducing Agents

The greater the
difference between
the two, the greater
the voltage of the
cell.
Figure 16.9
Cell emf and G

G for a redox reaction can be found by using
the equation:
G = −nFE

A positive value of E and a negative value of G
both indicate that a reaction is spontaneous.

Consequently, under standard conditions:
G = −nFE
n is the number of moles of electrons transferred.
F is called Faraday’s constant: 1 F = 96,485 C/mol
= 96,485 J/V mol
Cell emf Under Non-standard
Conditions - The Nernst
Equation

Remember that:
G = G + RT ln Q

This means:
−nFE = −nFE + RT ln Q
Cell emf Under Non-standard
Conditions - The Nernst
Equation

Dividing both sides by −nF, we get the
Nernst equation:
RT
E = E − nF
ln Q
or, using base-10 logarithms,
E = E −
2.303 RT
log Q
nF
Cell emf Under Non-standard
Conditions - The Nernst
Equation
At room temperature (298 K):

2.303 RT
= 0.0592 V
F

Thus the equation simplifies to:
E = E −
0.0592
log Q
n
Concentration Cells



Notice that the Nernst equation implies that a cell
could be created that has the same substance at
both electrodes.
 would be 0, but Q would not.
For such a cell, Ecell
Therefore, as long as the concentrations are
different, E will not be 0.
Figure 16.12
Batteries and Fuel Cells The Hydrogen Fuel Cell
Figure 16.16
Corrosion of Iron
Figure 16.17
Preventing the Corrosion of
Iron
Figure 16.18
Electrolysis
Figure 16.20

Voltaic cells are based on
a spontaneous redox
reaction, but it is possible
to use electricity to cause
non-spontaneous redox
reactions to occur.

For example, the
electrolysis of molten
sodium chloride to
produce sodium metal
and chlorine gas.
Electrolysis

Electroplating is another instance where
elecrolysis is used to deposit a thin layer of one
metal on another to improve beauty or resistance
to corrosion, for example the nickel-plating of
steel.
One point worth
noting is that in both
voltaic and electrolytic
cells, the anode is
where oxidation
occurs (and the
cathode is where
reduction occurs).
Figure 16.21
Quantitative Aspects of
Electrolysis

Half-reactions show how many electrons are
needed to achieve an electrolytic process and so
it follows that the amount of a substance oxidised
(or reduced) is directly proportional to the number
of electrons passed into the cell.
Figure 16.22