a. An atom in its elemental form has an oxidation number of 0.

1. Na, N2, Cu, O3 all = 0

b. Any monoatomic ion has the oxidation number equal to the charge on the ion.

1. Fe3+ = +3

c. Oxygen in a compound has an oxidation number of -2 (exceptions are peroxides, O2
2- =-1).

1. CO2, SO3
2. Exception: Na2O2 (Na= +1 and O = -1)

d. Hydrogen in a compound has an oxidation number of +1 (exceptions are hydrides where H = -1).

1. NH3, HCl
2. Exception: LiH

e. In binary compounds, the more electronegative atom is given the oxidation state of its species (ie.
the common ion of its group). Fluorine is always -1.
f. Atoms in groups 1 and 2 are +1 and +2, respectively.
g. The sum of the oxidation state must be equal to the overall charge of the species.

Example Redox Reactions

Synthesis
2 Na (s) + Cl2 (g)  2NaCl (s)

Na (s) = 0 changed to Na in NaCl = +1 : Na lost an e-
Cl2 (g) = 0 changed to Cl in NaCl = -1 : Cl gained an e-
Decomposition

CaO (s)  Ca (s) + O2 (g)

Ca in CaO (s) = +2 changed to Ca (s) = 0 : Ca gained 2 e-
O in CaO (s) = -2 changed to O2 (g) = 0 : O lost 2 e-
Single Replacement (oxidation of metals by acid and salts)

Zn (s) + HCl (aq)  H2 (g) + ZnCl2 (aq)
Al (s) + CuCl2 (aq)  Cu (s) + AlCl3 (aq)
Use an activity series table to determine if the reaction will occur. If the elemental metal is higher or more
reactive than the one it is replacing then a reaction will occur. The above reactions will occur, but the
reverse of either will not.
Combustion of hydrocarbons:
CH4 + O2  CO2 + H2O
The O in O2 =0, and the O in both CO2 and H2O has a -2 oxidation number (reduction occurs; gain 2 e-)
The C in CH4 = -4, while the oxidation number of C in CO2 is +4 (oxidation occurs; lost 8e-)