oxidation-reduction reaction: a reaction in which one or more electrons is transferred from one atom to another.
Originally, oxidation meant that an atom was combined with oxygen, and was therefore “oxidized”. For example:
2 Cu + O2 → 2 CuO
If we split this reaction into two “half-reactions,” we would have:
2 Cu0 → 2 Cu+2 + 2 e−
O20 + 2 e− → 2 O−2
In the first half reaction, copper (which was “oxidized”) lost electrons. In the other half reaction, oxygen gained electrons.
oxidation: the loss of one or more electrons by an atom in a chemical reaction
reduction: the gain of one or more electrons by an atom in a chemical reaction.
Stupid Mnemonics: There are two popular mnemonics for remembering oxidation and reduction, one “Democratic” and one “Republican”.
LEO the lion says ‘GER’ (“Democratic” mnemonic involving endangered species): LEO stands for “Loss of Electrons is Oxidation” and GER stands for “Gain of Electrons is Reduction”
OIL RIG (“Republican” mnemonic involving oil companies): OIL stands for “Oxidation Involves Loss (of electrons)”, and RIG stands for “Reduction Involves Gain (of electrons.”
In a redox reaction, at least one element is oxidized, and at least one element is reduced. An element cannot be oxidized in a chemical reaction unless some other element is reduced, and vice-versa.
oxidizing agent (or “oxidant”): the compound or ion that causes something else to be oxidized. (The oxidizing agent gets reduced in the process.)
reducing agent (or “reductant”): the compound or ion that causes something else to be reduced. (The reducing agent gets oxidized in the process.)
oxidation number: the charge that an atom would have in a compound if all bonds were completely ionic.
· The oxidation number of an element is 0. (Even if it’s diatomic.)
· The oxidation numbers of the elements in a compound add up to 0
· The oxidation numbers of the elements in a polyatomic ion add up to the charge of the polyatomic ion.
· In a compound:
o The most electronegative element (the last one in the formula) has a negative oxidation number equal to the number of electrons it needs to fill its valent shell.
o Fluorine is always −1.
o Oxygen is always −2 except in OF2.
o Hydrogen is always +1 except in metal hydrides.
o Alkali metals are always +1.
o Alkaline Earth metals are always +2.
o Al is always +3, Zn is always +2, and Ag is always +1.
o Calculate other elements from the above.
To fully balance a redox reaction, you must balance:
· Atoms (as you would in a regular equation)
· Electrons
· Total charge
Often, redox reactions are shown and balanced as net ionic equations. In this case, balancing them is a simple matter of making sure that the same number of electrons that are produced by the oxidation half-reaction are consumed by the reduction half-reaction.
For example, consider the unbalanced net ionic equation:
Al0 (s) + Fe2+ (aq) → Al3+ (aq) + Fe0 (s)
In this reaction, Al is oxidized from Al0 to Al+3, and Fe is reduced from Fe+2 to Fe0.
The two half-reactions are:
Al0 → Al3+ + 3 e−
Fe2+ + 2 e− → Fe0
To balance the electrons, we need to multiply the first half-reaction by 2, and the second one by 3, giving:
2 Al0 → 2 Al3+ + 6 e−
3 Fe2+ + 6 e− → 3 Fe0
If we combine these and cancel the electrons (because we have the same number on both sides), we get the balanced net ionic equation:
2 Al0 (s) + 3 Fe2+ (aq) → 2 Al3+ (aq) + 3 Fe0 (s)
If we add the spectator ion Cl− to balance the charges on each ion, we would end up with the balanced single displacement reaction:
2 Al (s) + 3 FeCl2 (aq) → 2 AlCl3 (aq) + 3 Fe (s)