The document provides information about electrochemistry. It discusses oxidation-reduction reactions and how they involve the transfer of electrons between species. It explains how to assign oxidation numbers to keep track of electrons gained and lost. Balancing oxidation-reduction reactions using the half-reaction method is also covered. Finally, the document discusses voltaic cells, electrolytic cells, and applications of electrochemistry such as electroplating.

1. Electrochemistry
Chapter 18
Electrochemistry
In electrochemical reactions, electrons
are transferred from one species to
another.

2. Electrochemistry
Oxidation Numbers
Assigning oxidation
numbers to all
species allows us to
keep track of the
electrons lost or
gained.

3. Electrochemistry
Assigning Oxidation Numbers
1. Elements in their elemental form are given a zero.
2. Monatomic ions are the same as their charge.
3. Nonmetals tend to have negative oxidation
numbers, although there are some exceptions:
 Oxygen has an oxidation number of −2, except in
peroxide where it has a −1.
 Hydrogen is −1 when bonded to a metal, +1 when
bonded to a nonmetal

4. Electrochemistry
4. In a neutral compound the sum of the ox #s is zero..
5. In a polyatomic ion is the sum of the ox #s is the
charge on the ion.
6. Nonmetals tend to have negative oxidation numbers,
although some are positive in certain compounds or
ions.
 Fluorine always has an oxidation number of −1.
 Halogens unless bonded to oxygen in a polyatomic ion.
 ClO3
- , the chlorine is a +5

5. Electrochemistry
Oxidation and Reduction
• A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral
zinc metal to the Zn2+ ion.

6. Electrochemistry
Oxidation and Reduction
• A species is reduced when it gains electrons.
Here, each of the H+ gains an electron and they
combine to form H2.

7. Electrochemistry
Balancing Oxidation-Reduction Equations
(Half-Reaction Method)
The oxidation and reduction are treated as two separate processes.
Each one is balanced separately and then combined to balance the
overall reaction.
1. Assign oxidation numbers to determine what is oxidized and
what is reduced.
2. Write the oxidation and reduction half-reactions.
3. Balance the elements in each half reaction leaving the O
and H for last.
 Balance O by adding H2O
 Balance H by adding H+
 Balance charge by adding electrons

8. Electrochemistry
4. Multiply the half-reactions by integers so that the electrons
gained and lost are the same.
5. Add the half-reactions, subtracting things that appear on
both sides.
6. Balance for both mass and charge.
Practice problem:
Permanganate ion(MnO4
−) and oxalate ion(C2O4
2−)
MnO4
−(aq) + C2O4
2−(aq) Mn2+(aq) + CO2(aq)

9. Electrochemistry
#1.) assign oxidation numbers and determine what is reduced
and what is oxidized
MnO4
− + C2O4
2-  Mn2+ + CO2
+7 +3 +4
+2
Manganese goes from +7 to +2, it is reduced.
Carbon goes from +3 to +4, it is oxidized.

10. Electrochemistry
Oxidation ½ rxn
C2O4
2−  CO2
• To balance the carbon, add a coefficient of 2 and that
also balanced the oxygen.
C2O4
2−  2 CO2
• To balance the charge, add 2 electrons to the right side
C2O4
2−  2 CO2 + 2 e−
#2.) write out the half reactions and balance each one

11. Electrochemistry
Reduction ½ rxn
MnO4
−  Mn2+
• Manganese is balanced.
• To balance the oxygen, add 4 waters to the right side.
MnO4
−  Mn2+ + 4 H2O
• To balance the hydrogen, we add 8 H+ to the left side.
8 H+ + MnO4
−  Mn2+ + 4 H2O
• To balance the charge, we add 5 e− to the left side.
5 e− + 8 H+ + MnO4
−  Mn2+ + 4 H2O

12. Electrochemistry
Add the Half-Reactions
Balance for both mass and charge
C2O4
2−  2 CO2 + 2 e−
5 e− + 8 H+ + MnO4
−  Mn2+ + 4 H2O
Balance charge, by multiplying the first reaction by 5 and the
second by 2 and then add them together.
10 e− + 16 H+ + 2 MnO4
− + 5 C2O4
2− 
2 Mn2+ + 8 H2O + 10 CO2 +10 e−
Subtract anything that appears on both sides of the equation.
In this case it is the electrons.
16 H+ + 2 MnO4
− + 5 C2O4
2− 2 Mn2+ + 8 H2O + 10 CO2

13. Electrochemistry
Electrochemistry
• The study of chemical reactions that produce
electricity, and chemical reactions that occur due
to electricity.
• There are many processes that take place
because of electrochemical rxns:
Electroplating
Electrolysis of water
Production of aluminum metal
Storage of electricity in batteries

14. Electrochemistry
Direct Electron Transfer
• In (redox) reactions,
electrons are
transferred and
energy is released.
• Cu gains e- and Zn
loses e-
• This type of transfer
doesn’t allow for any
useful work to be
done by the
electrons.

15. Electrochemistry
• We can use the energy
from the transfer of
electrons to do work if
we make the electrons
flow through an
external device.
• It provides power to do
work.
• We call such a setup :
Voltaic cell
Battery
Electrochemical cell
Galvanic cell

16. Electrochemistry
Typical
Voltaic
Cells
• The oxidation occurs at the ANODE.
• The reduction occurs at the CATHODE.
• Electrons flow through the wire!

17. Electrochemistry
VOLTAIC CELLS
• Once even one electron flows from the anode to the
cathode, the charges(ions) in each beaker would not
be balanced and the flow of electrons would stop.
• Therefore, we use a salt bridge, usually a U-shaped
tube that contains a salt solution, to keep the
charges balanced. Only ions flow through the salt
bridge!
Cations move toward the CATHODE.
Anions move toward the ANODE

18. Electrochemistry
• 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.
• The ANODE will
DECREASE IN
SIZE

19. Electrochemistry
• As the electrons
reach the cathode,
cations in the
cathode compartment
are attracted to the
now negative
cathode.
• The metal is
deposited on the
cathode.
• THE CATHODE
WILL INCREASE IN
SIZE.

20. Electrochemistry
CELL POTENTIAL
• The potential difference between the anode and
cathode in a cell is called the cell potential, and is
designated Ecell.
• It is a measure of the driving force of the redox
reaction taking place in the cell.
• Cell potential is measured in volts (V).
1 V = 1
J
C

21. Electrochemistry
Standard Reduction Potentials
The species
with the more
positive (V)
will be
reduced.
The species
lowest on the
table, and the
more negative
(V) will be
oxidized.

22. Electrochemistry
Standard Cell Potentials
The cell potential at standard conditions can
be found by looking up the reduction potentials
for each half reaction on a table.
The standard cell potentials (net E0) of all galvanic cells
have positive values and happen spontaneously.
The oxidation half reaction will be reversed from how it
appears on the table, and the value of it’s voltage will
have a reversed sign too.

23. Electrochemistry
Cell Potentials
= +0.34 V +0.76 V
= +1.10 V
Ean = +0.76 V
Ecat = +0.34 V

24. Cell Notation

25. Electrochemistry
Special Electrochemical Cell
• A metal–air electrochemical cell is a portable
electrochemical cell that uses an anode made from pure
metal and an external porous cathode that lets in oxygen
from the air, typically with an aqueous electrolyte paste.
• The big advantage here
is that you only need to
store the metal (anode)
in the battery because
the oxygen comes from
the air.
• You get a battery that is
compact and lasts
longer.

26. Electrochemistry
Metal-Air ElectrochemicalCell
• It works by oxidation of
the metal at the anode
and reduction of oxygen
at the cathode to create a
current flow of electrons.
• The anode is made from
a pure metal such as
Lithium and the external
cathode is made of the
ambient air.
• An aqueous electrolyte
exists between the them.

27. Electrochemistry
Electrolytic Cells

28. Electrochemistry
Electrolytic Cells
• An electrolytic cell is an
electrochemical cell that
drives a non-spontaneous
redox reaction through an
electric current supplied by
an external source.
• They are often used to
decompose chemical
compounds, in a process
called electrolysis—the Greek
word lysis means to break up.

29. Electrochemistry
• Electrolysis of water is the
decomposition of water into
oxygen and hydrogen gas due
to the passage of an electric
current.
• The reaction has a standard
potential of −1.23 V, meaning it
ideally requires a potential
difference of 1.23 volts to
split water.
• In a non-spontaneous reaction
the E0
net is always less than
zero.

30. Electrochemistry
Electroplating
Electroplating is a procedure
that uses electrolysis to apply a
thin layer of a metal over the
surface of another metal.

31. Electrochemistry
• In electroplating, the anode is made up of the
metal you want to use to coat the surface of
another metal.
• A salt solution containing the metal that makes
up the anode is used. During electrolysis, the
anode metal is oxidized and goes into solution
as positive ions.
• These positive ions are then reduced on the
surface of the cathode (the metal you are
coating)

32. Electrochemistry
Quantitative Electrochemistry
ELECTRIC CURRENT
• Electric Current is the rate of flow of Electric charges
• The SI unit of Electric Current is Ampere (A).
I = Q/T
• Where,
• I = Electric current in amperes,
• Q = Amount of charge in coulombs,
• T = Time in seconds.

33. Electrochemistry
Thermochemistry and
Electrochemistry Connection
∆G Gibb’s Free Energy
n number of electrons transferred
F Faraday constant
(96,500 C/mol or J/C)
E potential difference (V)
 E is (+) then rxn is spontaneous
 E is(-) then rxn is not spontaneous

34. Electrochemistry
Practice Problem:
If the standard cell potential at 298 K is 1.10 V for the
following reaction Zn(s) + Cu2+
(aq)  Zn2+
(aq) + Cu(s), then what
is the change in Gibbs Free Energy?

35. Electrochemistry
Gibbs Free Energy and Equilibrium
ΔG° = - RT lnK
The equilibrium constant for the reaction below is 1.8 × 1019 at 298K.
What is the value of the standard cell potential for this reaction?
Ni(s) + Hg2Cl2(s)  2Hg(l) + 2Cl-
(aq) + Ni2+
(aq)
ΔG = -RT ln K
= - (8.314 J K-1 mol-1) (298 K) ln (1.8 × 1019)
= - 109847.8 J mol-1
= - 1.098 × 105 J mol-1
E= -ΔG°
n F
= -(-1.098 × 105 J mol-1)
(2 mol) (96500 C)
= 0.57 J/C
They can be used to together to
find E(V)

36. Electrochemistry
This one is not
on the equation
sheet but it can
be derived by
setting the
other two
equations
equal to one
another and
simplifying.
Relationships between ∆G
◦, K , E
◦ cell

37. Electrochemistry