Oxidation number for elements by itself
0
For monatomic ions:
Their oxidation number is equal to their charge
All group 1 atoms (alkali metals) have an oxidation number of
+1
Group 2 atoms (alkali earth metals) in compounds have an oxidation number of
+2
Halogens in compounds have an oxidation number of
-1
Aluminum has an oxidation number of
+3
Hydrogen’s oxidation number is usually
+1
Exception for oxidation # for hydrogen
When combined with a metal it is -1
Oxidation number for oxygen
-2
All atoms are neutral so oxidation #s must add to
0
Trick to remember oxidation and reduction
O - oxidation
I - is
L - losing electrons
R - reduction
I - is
G - gaining electrons
Oxidizing agent
Allows element to be oxidized
-BEING REDUCED
Reducing agent
Allows an element to be reduced
-BEING OXIDIZED
Metals (and non metals) that are more active will
Replace those that are below them on the table, in the compound
Thus, reaction occurs (its spontaneous)
Oxidation half reaction:
Electrons lost must be shown on product side of equation
Reduction half reaction:
Electrons gained must be shown on reactant side of equation
Electrodes
Metal bars
Salt Bridge
Keeps solutions neutral
Half Cells
Have solutions of dissolved ions
Voltaic Cells
- Spontaneous Reaction
- Electrons flow from the anode to the cathode
- Uses chemical energy to create electrical energy (batteries!)
Anode
- Negatively charged electrode
- More active metal
- Metal atoms are oxidized (lose electrons)
- The electrode loses mass as atoms become ions and dissolve into the solution
Cathode
- Positively charged electrode
- Less reactive metal
- Metal ions from the solution get reduced (gain electrons)
- The electrode gains mass as ions from the solution become atoms attached to the electrode
Metals on Table J
- Higher up → more easily oxidized
- Lower down → least easily oxidized
Non-metals on Table J
- Higher up → most easily reduced
- Lower down → least easily reduced
