Here's a refresher on the rules for Redox reactions (see number 7 on Getting Started worksheet).
1. Each atom in a pure element has an oxidation number of zero. Example: Mg(solid), O2(diatomic oxygen)
2. For ions consisting of a single atom, the oxidation number is equal to the charge on the ion. Example: Ca2+, Cl-
3. Fluorine is always -1 in compounds with other elements.
4. The oxidation number of H is +1 and of O is -2 in most compounds.
5. Cl, Br, and I are always -1 in compounds (except when combined with oxygen and fluorine.)
6. The algebraic sum of the oxidation numbers in a neutral compound must be zero; in a polyatomic ion, the sum must be equal to the ion charge. Example: NaCl (Na is +1, Cl is -1), OCl- (O is +2, so Cl is -3 to get overall charge of -1).
Group one metals are +1, group two metals are +2.
Transition metals can have different oxidation numbers, no rule of thumb. Example: Fe +2, Fe +3
There may be exceptions to some of these rules, but for our class, this is what you need to know. We will get back to Redox reactions when we cover Electrochemistry later this semester.
Here's a link for some more practice:
Oxidation Numbers
