Download Oxidation Number Rules

Oxidation Number Rules
1. The oxidation number for an atom in its elemental form is always zero.
o A substance is elemental if both of the following are true:
 only one kind of atom is present
 charge = 0
o Examples:
 S8: The oxidation number of S = 0
 Fe: The oxidation number of Fe = 0
2. The oxidation number of a monoatomic ion = charge of the monatomic
ion.
o Examples:
2 Oxidation number of S is -2.
3+
 Oxidation number of Al is +3.
The oxidation number of all Group 1A metals = +1 (unless elemental).
The oxidation number of all Group 2A metals = +2 (unless elemental).
3. Oxygen (O) has two possible oxidation numbers:
2o -1 in peroxides (O2 )....pretty uncommon
o -2 in all other compounds...most common
4. Hydrogen (H) has two possible oxidation numbers:
o +1 when bonded to a nonmetal
o -1 when bonded to a metal
5. The oxidation number of fluorine (F) is always -1. All other halogens
have a -1 oxidation number in compounds, except when with oxygen or
other halogens where their oxidation numbers can be positive.
6. The sum of the oxidation numbers of all atoms (or ions) in a neutral
compound = 0. The sum of the oxidation numbers of all atoms in a
polyatomic ion = char ge on the polyatomic ion.
Solubility Rules for Ionic Compounds (Dissociates 100%)
1
All compounds containing alkali metal cations and the ammonium ion are soluble.
2
All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble.
3
All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+,
or Hg22+.
4
All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or Ba2+.
5
All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and
Ba2+.
6
All other compounds containing PO43-, S2-, CO32-, CrO42-, SO32- and most other
anions are insoluble except those that also contain alkali metals or NH4+.
Strong Acids (Ionizes 100%)
HCl, HBr, HI, HClO4, HNO3, H2SO4
Strong Bases (Dissociates 100%)
NaOH, KOH, LiOH, Ba(OH)2, Ca(OH)2, Sr(OH)2
Net ionic equations:
soluble salts (aq), strong acids/bases break up into ions
g, l, insoluble salts (s), weak acid/bases do not break up into ions
Rules to balance redox equations
1.) Convert to net ionic form if equation is originally in molecular form (eliminate
spectator ions).
2.) Write half reactions.
3.) Balance atoms using H+ / OH- / H2O as needed:
acidic: H+ / H2O put water on side that needs O
basic: OH- / H2O put water on side that needs H but if there is no H involved then
put OH- on the side that needs the O in a 2:1 ratio 2OH- / H2O balance O with OH, double OH, add 1/2 water to
other side.
4.) Balance charges for half rxn using e-.
5.) Balance transfer/accept number of electron in whole reaction.
6.) Convert equation back to molecular form if necessary (re-apply spectator ions).