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: . Oxidation number of S2- is -2. Oxidation number of Al3+ 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: 2- o -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+, 2+ or Hg2 . 2+ 2+ 2+ 2+ 2+ 4 All sulfates are soluble except those containing Hg2 , Pb , Sr , Ca , or Ba . 5 All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and Ba2+. 3- 2- 2- 2- 2- 6 All other compounds containing PO4 , S , CO3 , CrO4 , SO3 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). .