Oxidation/reduction reactions are those types of reactions where an electron is transferred from one species to another.
An oxidation is a reaction where a reactant loses one or more electrons; this reactant is oxidized.
A reduction is a reaction where a reactant gains one or more electrons; this reactant is reduced.
Oxidations and reductions always occur together.
Other terminology:
Oxidizing agent or oxidant: a reactant that causes something else to be oxidized (it does this by being reduced).
Reducing agent or reductant: a reactant that causes something else to be reduced (it does this by being oxidized).
In order to keep track of electrons and how they move from atom to atom in a reaction, we use the idea of oxidation numbers.
An Oxidation Number is an artificial bookkeeping device that represents the charge on an atom in a compound. Formally, this is found by writing the Lewis Dot structure and assigning all the electrons in each shared bond to the more electronegative atom and then calculating the formal charge.
More practically, there are a few simple rules that can be used to rapidly assign oxidation numbers. (This is review from CHM 101.)
An atom in its elemental state has an oxidation number of 0.
The oxidation number of a monoatomic ion equals its charge.
Group 1 elements nearly always have oxidation numbers of +1.
Group 2 elements nearly always have oxidation numbers of +2.
F always has an oxidation number of –1 in compounds.
Group 17 (halogens) have an oxidation number of –1 when they are the most electronegative atom in the compound.
In most compounds, oxygen has an oxidation number of –2. The common exception to this is peroxides, where O has an oxidation number of –1.
Hydrogen has an oxidation number of +1 (acids) if it is less electronegative than the atom to which it is bonded and an oxidation number of –1 (hydrides) when it is more electronegative than the atom to which it is bonded.
The sum of all oxidation numbers in a compound or ion must equal the total charge on the compound or ion.
These Rules Must Be Memorized.
We can use oxidation numbers to identify oxidation and reduction reactions:
oxidation: the oxidation number of an atom increases from reactants to products.
reduction: the oxidation number of an atom decreases from reactants to products.
Usually, this cannot be done by inspection. Fortunately, the half–reaction method
is a systematic way to balance any oxidation/reduction reaction.
Oxidation/reduction chemistry is often pH sensitive, so different reactivity occurs under acid and base conditions.
Separate the two half–reactions, choosing only the species that have a change in oxidation number. Ignore spectator ions, H+(aq), or H2O(l) if they are not oxidized or reduced.
For each half-reaction:
Balance the mass of all atoms except H and O by inspection.
Balance the O mass by adding the correct number of moles of H2O(l) to the oxygen deficient side.
Balance the H mass by adding the correct number of moles of H+(aq) to the hydrogen deficient side.
Balance the charge by adding the correct number of moles of e– to the side with excess total positive charge.
Multiply each half-reaction by the appropriate integer so that both half-reactions have the same number of electrons.
Add the two half-reactions together and eliminate common species (e– always; H+(aq), H2O(l) sometimes).
Make sure that the smallest set of integer coefficients is being used; if necessary, divide all stoichiometric coefficients by the highest common factor.
Add in any counterions or spectator ions, if required, and balance by inspection.
Check for both mass and charge balance.
Balance as if in acid.
Identify the side of the equation with H+(aq) and the stoichiometic coefficient, n.
If H+(aq) is a reactant, write a second equation: n H2O(l) → n H+(aq) + n OH–(aq)
If H+(aq) is a product, write a second equation: n H+(aq) + n OH–(aq) → n H2O(l)
Add the two equations together and eliminate common species (H+(aq) always; H2O(l) sometimes).
Check for both mass and charge balance.
Complete and balance:
1. Mg(s) + NO3–(aq) → Mg2+(aq) + NO2(g) (pH = 1)
2. Cr2O72–(aq) + SO2(g) → Cr(OH)3(s) + SO42–(aq)
3. ClO2–(aq) → ClO–(aq) + ClO4–(aq) (pH = 14)
4. BrO3–(aq) + [Cr(OH)4]–(aq) → CrO42–(aq) + Br–(aq)