Outline:
3/30/07
Pick up Exam 2 – from me
 Pick up CAPA 17 - outside
 Seminar today @ 4pm

Today: Start Chapter 19
 Redox reactions
 Balancing redox reactions
Exam 2

Hard exam?
Exam 2 (121-03)
Avg. = 76.0%
10
9
8
# students
7
6
5
4
3
2
1
0
5
15
25
35
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65
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85
95
Score
Well done! 99, 99, 94, 92, 92, 89, 89, 88, 86, 86…
19.1 Recognizing Redox Reactions
Definition: Any reaction which
involves a change of oxidation states
for some of the species involved.
FeO(s) + CO (g)  Fe (s) + CO2(g)
Iron Oxide loses oxygen and is reduced.
 CO gains an oxygen and is oxidized.

What is actually happening: atoms are
gaining or losing electrons
Definitions!
Reducing Agent (RA): agent that brings
about reduction, RA is oxidized
 Oxidizing Agent (OA): agent that brings
about oxidation, OA is reduced

RA: loses electrons (is oxidized)
OA: gains electrons (is reduced)
LEO the lion goes GER
If one substance is oxidized, another
substance must be reduced. Thus the
reactions are called REDOX.
Oxidation Numbers: the electric
charge an atom has, or appears to have
1.
2.
3.
4.
Treat polyatomic ions separately.
Sum of all the oxidation numbers must
equal the charge of the species.
Hydrogen usually has an oxidation
number of +1.
The most electronegative atom in a
species has a negative oxidation number
equal to the number of electrons needed
to complete its valence octet.
Oxidation Number Hints
5. A pure element has an oxidation number of
zero [Mg(s), Cu(s), etc...].
6. For monatomic ions, the oxidation number
is equal to the charge on the ion (Ca2+ has a
oxidation number of 2+).
7. The oxidation number of H is +1 and O is –
2 in most compounds.
Exceptions:
 When H is bonded to a metal, it becomes a
–1. For example, in CaH2, Ca2+ and H-1).
 Oxygen can have a –1 charge only when it
is called a peroxide (H2O2)…the rest of the
time it is -2 (except as O2 = 0).
Example #1
 SO4 2-
Overall charge = -2
 Assign oxidation number for most
electronegative element first
 Oxidation number for Oxygen
 6-8= -2
 Now find Oxidation number for S
 Oxidation
number for sulfur = +6
Example 2
Assign oxidation numbers to all atoms
in these two reactions involving iron:
FeO(s) + CO (g)  Fe (l) +CO2(g)
FeCO3 (s)  FeO(s) + CO2(g)
19.2 Balancing Redox Reactions
You must balance electrons as well as atoms.
 The number of electrons transferred in a
reaction may not be obvious.
 Oxidation and reduction always happens in
pairs.
 If electrons are being lost by one species,
then they must be gained by another.
 So, we separate the reaction into oxidation
and reduction using half-reactions.

How to balance redox rxns:
Identify charge states; is it a redox rxn?
 Separate into half reactions LEO-GER
 Balance each half-reaction for mass
(adding: H2O, H+, OH-)
 Balance each half-reaction for charge e Equate charge
 Add half-reactions together

Example
Separate the following unbalanced redox
processes into half-reactions. All occur in
aqueous solution:
a.
b.
c.
Cr2O72-(aq) + Fe2+(aq)  Cr2+(aq) + Fe3+(aq)
H2O2(aq) + SO3- (aq)  SO42-(aq)
BrO2-(aq)  BrO3-(aq) + Br-(aq)