**Oxidation-Reduction Reactions**: electron transfer reactions; one substance loses electrons while the other gains electrons; charges will change; typically exothermic and gas producing.

- Oxidation: process of
*losing electrons*; results in the charge of atom*increasing* - Reduction: process of
*gaining electrons*; results in the charge of the atom*decreasing*

The A, B, C’s of Redox Reactions:

1. Redox reactions can take the form of composition, decomposition, exchange/displacement.

2. They typically involve a reaction between two metals (and one of their respective salts) or a metal and hydrogen.

3. Combustion reactions are typically redox reactions.

4. For each element involved in a redox reaction, you should determine its charge, or oxidation number, before and after the reaction.

Oxidation Number: same as charge on the atom; represents the number of electrons lost or gained; an atom can have more than one form or oxidation number (ie Cl^{+7}, Cl^{+5}, Cl^{-1})

5. Identify which element’s charge increased. That atom has been oxidized. The atom which has been reduced will see a decrease in its charge.

- Oxidized element: charge increases; also called the reducing agent
- Reduced element: charge decreases; also called the oxidizing agent

6. The changes in charge or oxidation number are equal to the number of electrons lost/gained.

NOTE: Since electrons are matter, they must be conserved; therefore, in any redox reaction, the number of electrons lost must equal the number of electrons gained. We achieve this by ensuring the chemical equation is always balanced.

**Guidelines for Determining Oxidation Number:**

- Atoms in their original form and on their own have an oxidation number of zero.
- Ex. O
_{2}, Zn, Cl_{2}, S_{8}, Cu, H_{2}

- Ex. O
- Monoatomic ions have an oxidation number equal to their charges. If an ion already has a charge listed, that is its oxidation number.
- Ex. Na
^{+}= +1 - Ex: Al
^{3+}= +3

- Ex. Na
- Hydrogen has an oxidation number of +1 (except when bonded to a metal in which case it has a -1 charge).
- Oxygen has oxidation number of -2 (except in peroxides H
_{2}0_{2}where its oxidation number is -1) - Fluorine is always -1. Group I metals are +1 and group II metals are +2
- If a polyatomic ion appears in the
*same*form on both sides of reaction, then their oxidation numbers do not change. - All other atoms’ oxidation numbers must be calculated using the known oxidation numbers of the other atoms to solve for the unknown atom.
- The sum of all oxidation numbers in a compound should equal the net charge on the compound. Neutral atoms should have charges adding up to zero.

NH_{4}** ^{+}** : all oxidation numbers should add up to

**+1**

H_{2}O: all oxidation numbers should add up to zero. Water is neutral.

Ex. NH_{4}^{+}

Using Rule #3 above, “H” is +1. Therefore, we solve for the oxidation number of “N,” the unknown. The sum of 1 N and 4 H atoms should equal +1, the charge on ammonium.

Ex. ClF_{4}

Using Rule #5, we known that “F” is -1. Therefore, we solve for the oxidation number of “Cl.” Since ClF_{4} is neutral, the sum of 1 Cl and 4 F atoms should equal zero.

HINT: Think of these problem like algebraic equations, where the unknown element is “x.”