UNIT 5
Electrochemistry
 Outline the development of oxidation and reduction reactions.
 Determine the oxidation numbers for atoms in compounds and
ions.
 Identify reactions as redox or non redox.
 Balance oxidation-reduction reactions using redox methods.
 Research practical applications of redox reactions.
 Develop an activity series experimentally.
 Predict the spontaneity of reactions using an activity series.
 Outline the historical development of voltaic (galvanic) cells.
 Explain the operation of a voltaic (galvanic) cell.
 Construct a functioning voltaic (galvanic) cell and measure its
potential.
 Define standard electrode potential.
 Calculate standard cell potentials given standard electrode
potentials.
 Predict the spontaneity of reactions using standard electrode
potentials.
 Compare and contrast voltaic (galvanic) and electrolytic cells.
 Explain the operation of an electrolytic cell.
 Describe practical uses of electrolytic cells.
 Using Faraday’s law, solve problems related to electrolytic
cells.
Oxidation and Reduction
Oxidation reduction reactions and the loss
and gain of electrons have been studied
since the early 1800’s. It is oxidation
reduction reactions and the movement of
electrons that has focused the attention of
scientists to help solve our energy crisis and
our struggle against global warming and the
health of our planet. Many scientists believe
that hydrogen cells and fuel cells are the
way of the future. If you are to make
environmentally sound choices for the
future, you should at least understand the
current electrochemical technologies you
will be using.
Demonstration:
A piece of solid copper immersed in a
colourless silver nitrate solution.
Draw a picture of the copper and silver
nitrate solution at the moment the copper is
immersed in the solution and 2-3 minutes
later.
Time = 0 min.
Time = 3 min
What evidence is there that a reaction is
taking place?
Can you explain what is happening?
Write an equation that symbolizes the
reaction that is taking place.
What does the equation tell you?
Oxidation and Reduction
Using your textbook, please find definitions
for the following terms:
Oxidation
Reduction
Reducing agent
Oxidizing agent
In your textbook, find three examples of an
oxidation-reduction reactions or redox
reactions.
Oxidation and Reduction
Examples of an oxidation-reduction
reactions or redox reactions.
Reactions in batteries
Burning of wood
Corrosion of metal
Ripening of fruit
Combustion of gasoline
Oxidation: process by which electrons are
removed from an atom or ion.
Reduction: process by which any atom or
ion gains electrons.
Reducing agent: causes the reduction of
another species
Oxidizing agent: causes the oxidation of
another species
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Oxidation and Reduction
How do we determine what is being
oxidized and what is being reduced?
Cu(s) + 2AgNO3  2Ag(s) + Cu(NO3)2
Step 1: Write out the reaction in ion form.
Cu(s) + 2Ag+ + 2NO3-  2Ag(s) + Cu2+ +
2NO3Step 2: Look at the change in ion charge as
a function of electrons.
Cu(s)  Cu2+ + 2ethese are
half-reactions 2Ag+ + 2e-  2Ag(s)
*the charges need to balance on each side*
*add enough electrons to balance the
charges*
Step 3: Looking at just the reactants, ask
yourself which one has lost electrons.
That will be the one that is oxidized.
The other reactant will then be the one that
has been reduced.
Cu(s)  Cu2+ + 2e- (oxidation)
2Ag+ + 2e-  2Ag(s) (reduction)
Example:
Mg(s) + Cl2(g)  MgCl2(aq)
Mg(s)  Mg2+ + 2e- (oxidation) (reducing
agent)
Cl2(g) + 2e-  2Cl- (reduction) (oxidizing
agent)
Half-Reaction Practice Problems
1. Write balanced half-reactions from the net ionic equations below. State
which atoms are being oxidized and reduced.
a. Al(s) + Fe3+(aq)  Al3+(aq) + Fe(s)
b. Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s)
c. Cd(s) + 2Ag+(aq)  Cd2+(aq) + 2Ag(s)
2. Write balanced half-reactions for each of the following reactions. Hint:
write the NET ionic equation first. Identify the oxidizing and reducing
agents as well as which atoms are being oxidized and reduced.
a. Sn(s) + PbCl2(aq)  SnCl2(aq) + Pb(s)
b. Au(NO3)3(aq) + 3Ag(s)  3AgNO3(aq) + Au(s)
c. 3Zn(s) + Fe2(SO4)3(aq)  3ZnSO4(aq) + 2Fe(s)
Oxidation Numbers
In complex reactions it is not always
obvious what is being reduced or oxidized.
Chemists have created a set of rules to
allow us to determine more easily the
oxidation number of a given element within
a compound or complex ion.
Oxidation number:
the charge an atom
has, or appears to have, when the electrons
of the compound are counted in accordance
with a set of rules.
In a sense, assigning oxidation numbers is
arbitrary, but it is useful because it allows
you to follow the exchange of electrons that
occurs in redox reactions.
Note: oxidation numbers are written as +2
while ionic charge is written as 2+
Oxidation Number Rules
When determining oxidation numbers, the rules must be followed in the order
given.
1. The oxidation number of free elements is zero.
2. The oxidation number of a monatomic ion is equal to the charge on the
ion.
3. Fluorine always has an oxidation number of -1 when it is bonded to
another element.
4. The oxidation number of hydrogen in most compounds is +1. EXCEPT:
When hydrogen is bonded to metals in binary compounds it has an
oxidation number of -1.
5. The oxidation number of oxygen in compounds is always -2.
EXCEPT: The oxygen has an oxidation number of -1 in peroxide
compounds
6. The oxidation number of halogens is always -1.
7. The metals of groups 1A and 2A and aluminium form compounds in which
the metal atom always has a positive oxidation number equal to the
number of its valence electrons.
8. The sum of the oxidation numbers in a neutral compound is zero.
9. The sum of the oxidation numbers of the atoms in a polyatomic ion is
equal to the charge on the ion.
Oxidation Numbers Rules
Example: Determine the oxidation number
of the elements written in bold.
HNO3
Na3PO4
Cr2O72-
Try the following on your own:
V2O5
H2CO3
Ra(NO2)2