Oxidation states

      The first step in any electrochem problem is determining the oxidation numbers of the reactants and products. When an element - typically a metal - loses an electron, it becomes a cation, and has a positive charge. On the other hand, when an element - typically a nonmetal - gains an electron, it becomes an anion, and has a negative charge.

      Ex: Chlorine (nonmetal) gains an electron: Cl --> Cl-

      Ex: Sodium (metal) loses an electron: Na --> Na+

      *Free elements have an oxidation state of zero. This is typical of noble gases because they initially have a full valence shell, and don't have to gain or lose electrons to become stable (Neon --> Ne)

      This process works for all elements of the periodic and an element can gain or lose multiple charges, not just one electron. Specifically for transition metals, they have d-orbitals, which means that in order to become stable and get a full valence shell, they can gain or lose multiple electrons.

      Ex: Iron (transition metal) loses two or three electrons: Fe --> Fe+2, Fe+3

      *An easy tip is to remember the following rules, which is also seen in the periodic table to the left. There are more exceptions, but these are the most common and useful rules:

- All alkali metals are always +1 (except Hydrogen, which can be -1)

- All alkaline earth metals are always +2

- Transition metals typically have multiple charges (with a few exceptions)

- All halogens always have a -1 charge (some have other positive charges as well)

- Etc.