Electrochemistry : Oxidation and Reduction

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Electrochemistry Oxidation and Reduction
Electrochemical Reaction - Chemical reaction that
involves the flow of electrons. Redox Reaction
(oxidation-reduction reaction) - A reaction in
which at least one atom
changes in oxidation state.
Reduction - Any process in which the oxidation
number of an atom
decreases (becomes more negative). Oxidation -
Any process in which the oxidation number of an
atom increases (becomes
more positive).
Oxidation Number - The charge that an atom would
have if the compound in
which it were found were ionic. (Next page is a
refresher on How to.)
To help remember oxidation and reduction,
remember the following OILRIG Oxidation
Is Loss Reduction Is Gain
Types of Redox Reactions Corrosion - A type of
redox reaction in which a metal is destroyed.
4 Fe(s) 3 O2(g) 2 Fe2O3 3
H2O Metathesis Reaction - A reaction in which
atoms are interchanged and
there is no change in oxidation number.
Disproportionation Reaction - A reaction in which
a single reactant undergoes
both oxidation and reduction.

Disproportionation
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Before we balance a Redox equation lets first
refresh our memory on how to calculate oxidation
numbers.
Oxidation Number - The charge that an atom would
have if the compound in which it were found were
ionic.
The rules 1) The sum of the oxidation
numbers of the atoms in a molecule must be equal
to the overall charge on the molecule.
2) To assign a number to a transition
metal ion (not listed in the table below) start
with the overall charge, add the total number of
negative charges for oxygen (if there were four
as in the case of MnO4- then you would add 8
for a total of 7 for Mn), continue until all
other species listed in the table below are
considered (subtract if it is a positive value.)
The result is the oxidation number of the
transition metal ion. 3)The most
electronegative element will have a negative
oxidation number.
Assigning Oxidation Numbers Category

Oxidation
Example 1) Neutral substances
containing only a single element 0 N2, He
2) Monatomic ions
same as the charge Na 1
3) Hydrogen combined with a nonmetal 1
HBr, CH4, OH- 4) Hydrogen combined with
a metal -1 NaH, CaH2 5) Metals in
Group IA 1 Li3N, Na2S 6) Metals in
Group IIA 2 Mg3N2 7) Oxygen
-2 H2O, NO (Exceptions H2O2, O22-) -1
8) Halogens -1 AlF3, HCl
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Balancing Redox equations using the Oxidation
number method (Basic
solution is demonstrated)
1. What are the reduction and oxidation pairs?
a) Al(s) and Al(OH)-4(aq) (oxidized)
b) ? and H2(g) (reduced)
Hint 1. The reaction is taking place in a
basic, aqueous media.
2. Look for a reduction potential for H2(g) in
a table.
2. Calculate the Oxidation numbers
and transfer to the redox partner
3. Mass Balance
4. Charge Balance
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Balancing Redox equations using the Half Reaction
method (Basic
solution is demonstrated)
1. What are the reduction and oxidation pairs?
2. Mass Balance both equations
(add H2O to balance extra oxygens
then add extra H to balance extra hydrogens from
the added H2O)
3. Charge balance both equations (add extra e-)
4. Cancel any common terms.
5. Is the reaction taking place in a basic
solution?

Are there are any
H left? add OH- to both sides.
H and OH- will make H2O on one side.
6. Add the two half reactions and cancel any
extra water.
H2O
H
2e-
H2O
OH-
2 H
2e-
OH-
H2O 2e-
OCl-(aq)
H2O
Pb(OH)-3(aq)
PbO2(s)
Cl-(aq)
2H2O
OH-