Redox Reactions and Oxidation States
Each element in a chemical equation has an oxidation state, and you will have to assign these oxidation states to each element in order to determine the correct redox reaction stoichiometry. These states can be fairly easily determined just by looking at the periodic table.
To start, I’d recommend looking at the elements in rows 1 and 7. These atoms will have the least flexibility, and will in almost every circumstance have an oxidation state of +1 and -1 respectively. As you move towards the transition metals, elements will gradually gain more flexibility. For example, oxygen is almost always -2, but at times can have an oxidation state of -1 (think H2O – Each hydrogen has a +1, this cannot change, so the oxygen must have a -2 to cancel out the charges. However, in peroxides, oxygen is forced to have a -1, H2O2, because it has more flexibility than hydrogen). The transition metals should be the last elements of the equation you need to determine. These metals will have a positive value, but it can be anything from copper with a +1 or +2 state to manganese with +2, +3, +4, +6, or +7.
When I determine oxidation states, these are the steps I take, in order:
All free elements and ions get their charge
Free, neutral elements get a state of 0
Free, ionic elements get a state that is equivalent to their charge
All alkali metals get +1, all halogens get -1
All alkaline earth metals get +2
Chalogens (column 6) get a -2 unless required to be something else
Transition metals get the difference to make a polyatomic compound neutral
Redox reactions are the combination of a reduction reaction and an oxidation reaction. The reduction portion involves a transition metal to be reduced, or gain an electron, while the oxidation portion involves a transition metal to be oxidized, or lose an electron. This gain or loss of electrons changes the bonding properties of the metal, and will often convert between its neutral elemental form and its cationic form that bonds with an anion.
In redox reactions, you must first assign each element an oxidation state. Then, all you’re doing is switching the non-transition metal ion to the other metal, and rebalancing the stoichiometry. On each side of the equation you’ll have an ionic metal and a solid metal, the difference being that on either side, the state of the metal is switched. However, the oxidation state for the transition metal will change. This is the whole purpose of the redox reaction. One metal is getting reduced and the other is being oxidized, which leads us to the next question you may see on an exam: identify which gets reduced and which gets oxidized.
In class they probably taught you OIL RIG
Or there’s also LEO the lion says GER:
They both work, you just need to know one of them. Also, don’t get tripped up with the negative charge of the electron. By gaining an electron, you are gaining a negative charge, so the charge of the element is decreased by one. The opposite is also true, losing an electron causes the charge of the element to increase by one. Make sure you are counting the electrons being gained or lost and not the charge.
Finally, the oxidizing agent is the component that does the oxidizing. Therefore, the oxidizing agent, or oxidizer, gets reduced. If they ask about the reducing agent, this element will reduce the other, so it itself must be oxidized.
While you’re going through these problems it’s easy to confuse yourself. They’re usually the last few questions on the test, but don’t let yourself get lazy. Double and triple check the oxidation states and the stoichiometry, write it down if you have to. The simpler questions can be easy to miss.