1. Assigning Oxidation Numbers
RULES
RULES
Examples
Examples
2Na + Cl2  2NaCl
Na = 0 or written Na0
Cl2 = 0 or written Cl20
2Na + Cl2  2NaCl
1. Each Uncombined Element has an oxidation number = 0
Na = 0 or written Na0
Cl2 = 0 or written Cl20
Monatomic ions have an oxidation number equal to the ionic charge.
Al3+ then Al = +3
In KCl: K=+1
In CaCl2: Ca=+2
In MgO; Mg=+2
In Li2O: Li=+1
2. The metals of Group 1 always have an oxidation number of +1 and the metals of Group 2 always have an oxidation number of +2.
Group 1 = +1
Group 13 = +3
Group 2 = +2
In HF: F = -1
In O2F: F = -1
In CaCl2: Cl = -1
In HClO; oxygen is more electronegative than Cl  Cl=+1
3. Fluorine is ALWAYS = -1
The other halogens are also = -1 when they are the most electronegative element in the compound

2. Assigning Oxidation Numbers
RULES
RULES
Examples
Examples
In HF; H=+1
In H2SO4; H=+1
In CaH2: H=-1
In LiH: H=-1
In HF; H=+1
In H2SO4; H=+1
In CaH2: H=-1
In LiH: H=-1
4. Hydrogen is +1 in a compound EXCEPT when it is combined with a METAL. When combined with a METAL, Hydrogen = -1
In H2O: O=-2
In OF2: O=+2
In Na2O2 (sodium peroxide): O=-1
In H2O: O=-2
In OF2: O=+2
In Na2O2 (sodium peroxide): O=-1
5. Oxygen is -2
EXCEPT when with Fluorine, then Oxygen = +2
EXCEPT in peroxide ion (O22-) , then oxygen = -1
6. Most elements will be the charge on the periodic table.
Zn = +2
The sum of the oxidation numbers in all compounds must be ZERO
7a. The sum of the oxidation numbers in polyatiomic ions must equal the charge on the ion
In NaCl: (+1) + (-1) = 0
In CaCl2: (+2) + 2(-1) = 0
In Al2(SO4)3: 2(+3)+3(-2) =0
In SO42-: S + 4(-2) = -2
S must = +6

3. OXidation-REDuction Reactions (also called REDOX)
WHAT ARE….
OXidation-REDuction Reactions (also called REDOX)
OXIDATION
REDUCTION
What Happens to the electrons?
Electrons are LOST
(LEO = loss of electrons=oxidation)
Electrons are GAINED
GER: Gain of Electrons=Reduction
LEO the lion goes GER
What Happens to the charge?
Charge goes up
(The oxidation number on the
atom increases)
Charge goes down
(the oxidation number on the
Atom decreases)
Write the reaction
Include:
Balanced # of atoms
Balanced charge
Ex: Mg+Cl2MgCl2
Oxidation Half reaction is:
Mg  Mg2+ + 2e-
Ex: Mg+Cl2MgCl2
Reduction Half reaction is:
Cl2 + 2e-  2Cl-
Ex: Hg I-  Hg + I2
Reduction Half reaction is:
Hg2+ + 2e-  Hg
Ex: Hg I-  Hg + I2
Oxidation Half reaction is:
2I-  I2 + 2e-
What happens to the electrons if you add the two half reactions?
NOTE: e- are on right side of arrow
NOTE: e- are on left side of arrow

4. OXidation-REDuction Reactions (also called REDOX)
WHAT ARE….
OXidation-REDuction Reactions (also called REDOX)
OXIDATION
REDUCTION
Also called the…
Reducing agent
(the thing that gets oxidized
is doing the reducing so..)
Oxidizing Agent
(the thing that gets reduced
is doing the oxidizing so..)
ELECTRONS LOST
ELECTRONS GAINED
EVERYTHING IS ELECTRICALLY NEUTRAL!! ALL CHARGES MUST BALANCE (AND CANCEL)!
Identifying redox reactions
Assign oxidation numbers to both sides of the equation
Look to see if oxidation numbers change. If they do = redox (if not then not)
HINTS: Single replacement reactions are REDOX
Double replacement are NOT redox

5. How to Balance a Redox reaction:
Assign oxidation #’s to all of the atoms.
Identify which are oxidized & which are reduced.
Use one bracketing line to to connect the atoms that undergo oxidation & another to connect those that undergo reduction.
Make the total increase in oxidation # equal to the total decrease in oxidation # by using the appropriate coefficient.
Make sure that the equation is balanced for both atoms & charges.
Ex.
Fe2O3(s) + CO(g) Fe(s) + CO2(g)
-3 reduction
2 x (-3) = -6 3 2
+2 oxidation
3 x (+2) = +6

6. IS an application of REDOX reactions
ELECTROCHEMISTY
IS an application of REDOX reactions
Energy is produced or applied from REDOX reactions
Two different ways
ELECTROCHEMICAL
ELECTOLYTIC
Electricity FORCES a chemical change
(apply the electricity)
Chemicals PRODUCE electricity
(I.e., metals & solutions)
HALF CELL /Half Reaction:
»Show only 1/2 half of the reaction
Either oxidation rxn
OR reduction rxn
CELL POTENTIALS:
»Voltages that are produced by each cell
If E0 = 0.00, have DEAD BATTERY (equilibrium)
» BOTH mass, energy and charge are conserved!! (equal on both sides)
Strip of metal
Solution that
Contains that
Metal ions Al
Al3+

7. ELECTROCHEMISTY Spontaneous Reactions: USE TABLE J TO PREDICT
The neutral metal must be higher than the metal ion with which it’s reacting
Zn + Cu2+  Cu + Zn2+ On table J, Zn is higher than Cu = spontaneous
Cu + Mg2+  Cu2+ + Mg: NOT SPONTANEOUS, Cu not higher than Mg NEEDS BATTERY!
Electrons naturally flow from higher metal to lower metal

8. ELECTROCHEMICAL CELL (aka: Voltaic cells or Galvonic cells)
Chemical rxn produces electricity
SPONTANEOUS: -G
Chemistry electricity
+ volts and +E0
i.e. making a battery
How do we tell if it is spontaneous?
Salt Bridge
(U-tube)
Salt Bridge: ions migrate in both directions
Metals as electrodes Zn, Cu
Electrons flow from higher metal to lower metal on TABLE J
Zn higher than Cu
CATHODE: reduction occurs (Cu) Positive (+)
ANODE: oxidation occurs (Zn) Negative (-)
Zn0 gives up electrons
Zn mass
Ion Concentration
Cu+2 gains electrons
Mass of Cu
Ion Concentration
AN OX
Anode = oxidation
Decrease mass /Decrease e-
RED CAT
Reduction = cathode
Increase mass/increase e-

9. “Plating” or coating an
ELECTROLYTIC CELL
(NEEDS A BATTERY)
Apply electrical energy to make a chemical rxn happen.
Need a battery
NOT SPONTANEOUS!!
+G (-) volts
TWO MAJOR TYPES
Electrolysis
Separating a compound
Into its elements
Electroplating
“Plating” or coating an
Element onto an object
EXAMPLE OF THE TWO MAJOR TYPES:
BRINE SOLUTION (NaCl):
2NaCl  2Na + Cl2(g)
Battery - +
 Cl-
H+
OH-
Na+
H2(g)
Cl2(g)
Anode
Ox.
2Cl-  Cl2 + 2e-
Cathode
Red.
2H+ +2e- H2
e-
e-
ELECTROLYSIS: +1 -2
2H2O  2H O2
Anode
Oxidation
Cathode
Red.
2H+ +2e- H2
O-2 O2(g) + 2e-
e-
e- 

10. FOR ALL ELECTROLYTIC CELLS ELECTROPLATING (2nd type) Reduction
(NEEDS A BATTERY)
FOR ALL ELECTROLYTIC CELLS
ELECTROPLATING (2nd type)
Reduction
@ Cathode
RED CAT
Neg. term (-)
AN OX
Oxidation
@ Anode
Pos. term. (+)
Object to be plated-
Always cathode (reduction)
Always attached to negative term.
Metal used to plate-
Always anode (oxidation)
Always attached to
positive term.

11. Similarities & Differences:
Cu + Pb+2 ==> not spontaneous.
Zn+ Ag+ ==> Spontaneous!
Electrochemical/
Voltaic Cell
Electrolytic Cell
Table J
Metal is Higher than the ion (compound)
Summary:
What’s going on?
Chemicals producing electricity- EXO
(IS Spontaneous)
+ volts
Electricity forcing a chemical reaction to occur Endo
(NOT spontaneous)

12. IS a Battery! NEED a Battery! Anode Cathode Cathode + Attached Anode
Charge:
Positive or Negative?
Anode
+ Attached
To +
Cathode
Attached
to -
Cathode
+ lower on Table J
Anode
- Higher on Table J
Reduction
Oxidation
Oxidation
Reduction
An Ox & Red Cat
IS a battery or NEEDS a battery?
IS a Battery!
NEED a Battery!

13. e- flow anode cathode
Salt bridge- allows IONS to migrate in both directions
e- flow out- in + opposites attract
object to be plated is ALWAYS attached to the negative!!!

14. Electrolytic Electrochemical • NOT Spontaneous • Also known as Voltaic
• Needs a battery to force e- to flow in wanted direction
• Spontaneous-produces electricity
Both involve the transfer and movement of e- to make something happen
“LEO” & “GER”
• Anode is connected to (+) terminal of battery and is NOW (+) electrode
• Makes a battery
• Electron flow predicted by Table J. --From higher to lower metal.
• Cathode is connected to (-) terminal of battery and is now (-) electrode
• Oxidation occurs at the anode “An Ox”
• Anode is the (-) terminal
Used for :
• Electrolysis-separating compounds into elements
• Electroplating –
coating items
• Reduction occurs at the cathode “Red Cat”
• Cathode is the (+) terminal