Chapter 20: Electrochemsitry
A.P. Chemsitry
20.1 Oxidation-Reduction Reactions
• Oxidation-reduction
reactions (or redox reactions) involve the transfer
of electrons from one species to another. In other
words, the oxidation state of one or more
substances in the reaction changes. The transfer
of electrons inherent in an oxidation-reduction
process can be used to produce energy in the
form of electricity. Electrochemistry is the study
of the relationships between chemical reactions
and electrical energy.
20.2 Balancing Oxidation-Reduction
Equations
• A half-reaction is an
equation that shows
either oxidation or
reduction alone.
Reduction half-reaction
Zn (s)  Zn2+(aq) + 2 e-
Oxidation half-reaction
Balancing steps
1.
2.
Divide the equation into two incomplete half-reactions, one for oxidation and
the other for reduction.
Balance each half-reaction. [Assuming an acidic solution]
1.
2.
3.
4.
3.
4.
5.
6.
7.
First, balance the elements other than H and O
Next, balance the O atoms by adding H2O [OH- for basic solution]
Then, balance the H atoms by adding H+. [H2O for basic solution]
Finally, balance the charge by adding e- to the side with the greater overall positive charge.
Multiply each half-reaction by an integer so that the number of electrons lost in
one half-reaction equals the number gained in the other.
Add the two half-reactions, and simplify where possible by canceling species that
appear on both sides of the equation.
Check the equation to make sure that there are the same number of atoms of
each kind on both sides and the same total charge on both sides.
Neutralize any H+ ions in the equation with OH- ions, adding the same number of
OH- ions to the opposite side of the equation. (The combination of an H+ and an
OH- produces a water molecule.)
Cancel any resulting water molecules with water molecules on the opposite side
of the equation.
Reduction Half-Reaction Equation
Oxidation Half-Reaction Equation
Add electrons to each side to balance charge
Multiply to get same electrons in each half reaction
Then combine the half-reactions
Check to make sure there are the same number of each type of
atom on each side of the equation
20.3 Voltaic Cells
Salt Bridge
• In a voltaic cell the two
half-reactions are made
to occur in separate
compartments (half-cells)
• The electrode at which
oxidation occurs is called
the anode; in this case,
the zinc electrode. The
electrode at which
reduction takes place is
called the cathode; in this
case, the copper
electrode
20.4 Cell EMF Under Standard
Conditions
• 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.
• 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, E°cell (25°C,
1 M concentrations, and 1 atm
pressures), is the standard cell
potential or standard emf.
Calculating EMF
The standard reduction potentials for the various half-reactions are
measured against a standard hydrogen electrode (SHE). The halfreaction of interest and the SHE, both under standard conditions, are
made into a voltaic cell, as shown above, 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.
Sample Calculation
What E0 means
• 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 E°red for a species, the
more readily it undergoes
reduction and the
better oxidizing agent it is.
As E°red becomes more negative,
the species on the right side of
the arrow becomes a
stronger reducing agent.
20.5 Spontaneity of Redox Reactions
• Our ability to predict
the spontaneity of a
chemical reaction by
calculating a standard
cell potential points to a
relationship between
the sign of E°cell and the
sign of G° , the
standard change in
Gibbs free energy.
Sample Calculation
20.6 Effect of Concentration on EMF
• Just as the values
of G and G° were related, the
values of E and E° are related.
The Nernst equation derives
from the relationship
between G and G° .
• Substituting - nFE for G gives
• Solving for E gives
• The Nernst equation is most
commonly expressed in terms
of log (base 10) rather than ln
(natural log).
• When the temperature is 298
K, this expression simplifies to
Concentration Calculation
Calculate E0 
20.7 Batteries
• 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.
20.8 Corrosion
• 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.
• To undergo these oxidationreduction reactions, iron
must be in contact with
oxygen and water.
Electrolysis
•
•
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, 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.