1. Topic 9
Oxidation and Reduction

2. Topic 9
Oxidation and Reduction (Redox)
Oxidation and Reduction
OIL RIG/LEO goes GER
Oxidation Numbers
know the rules for assigning oxidation numbers
Recognizing Redox Reactions
a. Identifying Oxidation and Reduction
b. Identifying Oxidizing and Reducing Agents
Half Reactions
a. Be able to write and balance ½ reactions
Electrochemical Cells
a. Galvanvic Cells
b. Electrolytic Cells
c. Be able to draw and identify all parts

3. We have done types of reactions:
1. Direct Combination/Synthesis
2. Analysis/Decomposition
3. Single Replacement Reactions
4. Double Replacement Reactions
5. Combustion
The last type of reaction to be studied is Oxidation-Reduction
It is commonly referred to as redox since oxidation and reduction
must occur simultaneously
Oxidation:
Loss of electrons by an atom or ion. The oxidation
number of the atom or ion INCREASES
Reduction:
Gain of electrons by an atom or ion. The oxidation
number of the atom or ion DECREASES

4. We remember this by:
Oxidation Is Losing Electron
OIL RIG
Reduction Is Gaining Electrons or
Losing Electrons is Oxidation
LEO GER
Gaining Electrons is Reduction

5. Rules for Assigning Oxidation Numbers
Used to identify the path of electrons in
redox reactions
1. Each uncombined element has an oxidation number of zero
Ex.: 2Na + Cl2  2NaCl
Monatomic ions have an oxidation number equal to the ionic
charge +2 -1
Ex.: Mg2+
Cl1- -2
O2-

6. 3. The metals of Group 1 always have an oxidation number of +1
*Hydrogen is not a metal so it may have an oxidation number of +1 or -1
The metals of Group 2 always have an oxidation number of +2
4. The element with the lower EN value is + and usually written
first
The element with the higher EN value is – and usually written last
Oxygen is usually -2 except when it is combined with fluorine.
when with fluorine, it is +2
When oxygen is with peroxide ion (O22-), oxygen is -1

7. 6. The sum of the oxidation numbers in all compounds must be zero+1 -1 +2 -1 +2 +6 -2
NaCl
CaF2
MgSO4
The sum of the oxidation numbers in all ions must be equal
to the charge +7 -2 +6 -2 +4 -2
ClO41-
SO42-
CO32-
Exceptions: -2 +1 -3 +1 -3 +1
OH1-
NH3
NH41+ +1 -1 +2 -1 +1 -1 +2 -1
BaO2
H2O2
Na2O2
OF2

8. When looking at reactions, you must assign oxidation numbers to
all the elements. Depending on what happens to the oxidation
numbers of the elements will indicate what is oxidized and what
is reduced +1 -1
Na + Cl2  NaCl
Reduction:
Gain of electrons
Oxidation number decreases
Cl2 went from 0 to -1, therefore, it was reduced
Oxidation:
Loss of electrons
Oxidation number increases
Na went from 1 to +1, therefore it was oxidized

9. The substance oxidized+1 -1
Oxidizing agent
Na + Cl2  NaCl
oxid
red
Reducing agent
Oxidizing agent-
The substance reduced
Reducing agent-
The substance oxidized
Agents can only be found on the reactant (left) side of the
equation. They are never found on the product side.

10. +1 -1
Na + Cl2  NaCl
Half-Reactions
Reduction ½ reaction: -1
Cl2 +  Cl + e- 2 2
Oxidation ½ reaction: +1 +1 2 (
Na + e  Na ) =
2Na + 2e-  2Na
___________________
2Na + Cl2 2NaCl

11. To quickly recognize redox reactions, there must be a change in oxidation numbers of the elements
If an uncombined element appears on one side of an equation and
it is in a compound on the other side, the reaction must be a redox
reaction
Zn + 2HCl  H2 + ZnCl2
If you recognize a reaction as a double replacement reaction,
it is never a redox reaction
NaCl + HOH  NaOH + HCl

12. +1 +2
Red. agent
Cu + Ag1+  Cu2+ + Ag
oxid
Oxid. agent
red
1. Assign oxidation numbers to each element
2. Identify what is oxidized and what is reduced
3. Identify oxidizing and reducing agents
4. Write ½ reactions +1 +1
Red. ½ rx: Ag e-  Ag ( ) 2 =
2Ag + 2e-  2Ag +2
Oxid. ½ rx: Cu  Cu e- 2
5. Balance by inspection +2 +1
2Ag + Cu  Cu + 2Ag

13. mass, energy and charge must be conserved
Remember…in any chemical reaction…
mass, energy and charge must be conserved
Therefore…when writing ½ reactions, you must balance the
number of atoms/ions and the number of electrons

14. Electrochemical Cells
Involves a chemical reaction and a flow of electrons
2 types
1. Voltaic Cell/Galvanic Cell
a. Spontaneous reaction
b. Chemical energy is converted into electrical energy
2. Electrolytic Cell
a. Nonspontaneous reaction
b. Electrical energy converted into chemical energy

15. Voltaic Cell/Galvanic Cell
The lower metal
on Table J
will be reduced
in a voltaic cell
Voltaic Cell/Galvanic Cell e- e-
+ions
Electrons move
through the wire
toward the cathode
-ions
Ions move through
the salt bridge
+ ions to cathode
- Ions to anode
cathode
anode
Gains mass
Loses mass + -

16. A meter that indicates how much electricity is being generated
Voltmeter-
A meter that indicates how much electricity is being
generated
Not a power source
Wire-
Electrons move through the wire to the cathode
Salt Bridge-
Ions move through the salt bridge
+ ions to the cathode
- ions to the anode
Helps maintain electrical neutrality
If the cell potential is zero (on the meter),
the reaction is at equilibrium

17. How do we remember?
Regardless of what cell we are using:
We have AN ANOREXIC OX and a BIG RED CAT
AN ANOREXIC OX
Anode loses mass oxidation
BIG RED CAT
Gains mass reduction cathode

18. Voltaic Cell
Electrolytic Cell
Lower metal is oxidized
Lower metal is reduced + +

19. Electrolytic Cell
Electrical energy is converted into chemical energy
Nonspontaneous reaction e- e-
oxidation
reduction + -
Loses mass
Gains mass

20. Electroplating
Is the deposition of a thin layer of a metal on an object in an
electrolytic cell
The object to be plated must be the cathode
oxidation
anode
cathode +
reduction -

21. Similarities between voltaic and electrolytic cells
1. Both use redox reactions
2. The anode is the site of oxidation
3. The cathode is the site of reduction
4. The electrons flow through the wire from the anode to the cathode
5. Positive ions flow to the cathode
6. Negative ions flow to the anode

22. Differences between voltaic and electrolytic cells
The redox reaction in a voltaic cell is spontaneous,
but it is nonspontaneous in an electrolytic cell
In an electrolytic cell, the anode is positive and the
cathode is negative.
In a voltaic cell the anode is negative and the cathode is
positive.
The voltaic cell has a voltmeter and the electrolytic cell
has a battery (power source).
The voltaic cell has a salt bridge and the electrolytic
cell does not.