UNIT 10: REDOX
CDO CP CHEMISTRY
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OXIDATION AND REDUCTION
• Oxidation – Reduction (REDOX) Reactions: The
chemical changes that occur when electrons
are transferred between reactants
• More Fundamental definition
• Oxidation – loss of electrons (LEO)
• Reduction – gain of electrons (GER)
• In an oxidation reduction reaction both occur!
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OXIDATION NUMBERS
• Assign oxidation numbers to keep track of
electrons.
There are two parts to oxidation numbers:
• Sign - + means electrons are lost, - means electrons
are gained
• Value – number of electrons gained or lost
ASSIGNING OXIDATION NUMBERS
1. The sum of the oxidation numbers
in a neutral compound is 0.
2. The sum of the oxidation numbers
in a polyatomic ion is the charge on
the ion.
ASSIGNING OXIDATION NUMBERS
3. Elements in their elemental form
have an oxidation number of 0.
0
𝟎
Fe Cl2
4. The oxidation number of simple
ions is the same as its charge.
+1
+2
-3
Na +
Cu 2+
N3-
ASSIGNING OXIDATION NUMBERS
5. The oxidation number for
Fluorine is -1 .
6. Group 1 metals have an
oxidation number of +1
Group 2 metals are +2
Al, Sc, and Y are +3
ASSIGNING OXIDATION NUMBERS
7. Hydrogen has an oxidation
number of +1 when combined
with a nonmetal and a -1
when combined with a metal.
ASSIGNING OXIDATION NUMBERS
8. Oxygen has an oxidation
number of -2 unless
• It is combined with F, then it is
+2
• It is in a peroxide, then it is -1
ASSIGNING OXIDATION NUMBERS
9. The usual oxidation number is the
same as the charge of the most
common ion
Element
Exceptions
Group 7A
Usual
Oxidation
Number
-1
Group 6A
-2
When combined with O
or F it is positive
Group 5A
-3
When combined with O
or F it is positive
When combined with O
or F it is positive
EXAMPLE 1: CALCULATING OXIDATION
NUMBERS
•
•
•
•
•
SO2
SO42Cr2O72H 2 O2
HNO3
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REDOX REACTIONS
• A redox reaction is one that has changes in
oxidation number
• Reactions that are always redox
• Combustion – element O2 becoming a compound
• Synthesis – any element reacting with another to produce a
compound
• Ions changing charge
• Ions changing the number of oxygens
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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.
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.
OXIDATION AND REDUCTION
• What is reduced is the oxidizing agent.
• H+ oxidizes Zn by taking electrons from it.
• What is oxidized is the reducing agent.
• Zn reduces H+ by giving it electrons.
EXAMPLE 2
• Indicate which of the reactants is reducing and
which is oxidizing
• 2Ce4+(aq) + Sn2+(aq)  2Ce3+(aq) + Sn4+(aq)
• 8H+(aq) + MnO4-1(aq) + 5Fe2+(aq)  5Fe3+(aq) +
Mn2+(aq) + 4H2O(l)
BALANCING REDOX EQUATIONS
• Perhaps the easiest way to balance the equation
of an oxidation-reduction reaction is via the halfreaction method.
Balancing Redox Equations
oThis involves treating (on paper only) the oxidation
and reduction as two separate processes,
balancing these half reactions, and then
combining them to attain the balanced equation
for the overall reaction.
THE HALF-REACTION METHOD
(ACID)
•
•
Combine like compounds and write each half
reaction
Balance each half-reaction.
a.
b.
c.
d.
e.
Balance elements other than H and O.
Balance O by adding H2O.
Balance H by adding H+.
Balance charge by adding electrons.
Multiply the half-reactions by integers so that the
electrons gained and lost are the same.
THE HALF-REACTION METHOD
f.
g.
h.
Add the half-reactions, subtracting things that
appear on both sides.
Make sure the equation is balanced
according to mass.
Make sure the equation is balanced
according to charge.
THE HALF-REACTION METHOD
Consider the reaction between MnO4− and C2O42−
(acidic solution)
MnO4− (aq) + C2O42− (aq)  Mn2+ (aq) + CO2 (aq)
OXIDATION HALF-REACTION
C2O42−  CO2
Reduction Half-Reaction
MnO4−  Mn2+
COMBINING THE HALF-REACTIONS
Wolpa/Christman AP Chemistry
EXAMPLE
• Zn(s) + VO2+(aq) --> V2+(aq) + Zn2+(aq)
Wolpa/Christman AP Chemistry
EXAMPLE: BALANCING HALF
EQUATIONS
• 2 Br- + Cl2  2Cl- + Br2
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VOLTAIC CELLS
• A voltaic cell allows
a chemical reaction
to produce
electricity based on
a difference in
reactivity
• The greater the
difference in
reactivity the more
voltage the cell can
make
VOLTAIC CELLS
• In a cell the two half
reactions are kept
separate so the electrons
are forced to travel
through a circuit where
they can do work
• The oxidation occurs at
the anode.
• The reduction occurs at
the cathode.
RED CAT AN OX
VOLTAIC CELLS
• Once even one
electron flows from
the anode to the
cathode, the
charges in each
beaker would not
be balanced and
the flow of electrons
would stop.
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VOLTAIC CELLS
• Therefore, we use a
salt bridge, usually a Ushaped tube that
contains a salt
solution, to keep the
charges balanced
and completes the
circuit
• Cations move toward
the cathode.
• Anions move toward the
anode.
VOLTAIC CELLS
• 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.
VOLTAIC CELLS
• As the electrons
reach the
cathode, cations
in the cathode
are attracted to
the now negative
cathode.
• The electrons are
taken by the
cation, and the
neutral metal is
deposited on the
cathode.
VOLTAIC CELL
• Draw a voltaic cell and label the anode, cathode, ion
channel, voltmeter, flow of electrons, flow of ions.
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CELL NOTATION
Zn/Zn2+ ll Cu2+/Cu
o
o
o
anode reaction is shown at left
salt bridge is indicated by ll
cathode reaction is shown at right
CELL POTENTIAL
The cell potential for a voltaic cell can be
calculate using the following equation:
o
E
cell
=
o
E
cathode –
o
E
anode
Eocell is positive for spontaneous reactions
and negative for nonspontaneous
reactions
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CELL POTENTIAL
The substance with the larger potential is
the anode, and the substance with smaller
is the cathode
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CELL POTENTIAL
Calculate the standard potential for a
voltaic cell with the following reaction:
Al(s) + NO3- + 4H+  Al3+ + NO + H2O
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