Redox Reactions

Oxidation numbers, also known as oxidation states, are used to track electron transfer in chemical reactions.
The overall charge of a neutral compound is zero.
The oxidation number of an atom in its elemental form is zero. Examples include $O<em>2$ , $Na$ , and $H</em>2$ .

A redox reaction involves both reduction and oxidation processes.
If the oxidation number of a substance increases, it is oxidized. The substance loses electrons.
If the oxidation number of a substance decreases, it is reduced. The substance gains electrons.
If the oxidation number of an element does not change during a reaction, it is neither oxidized nor reduced.

Oxidation is the process where a substance loses electrons, resulting in an increase in its oxidation number.
A substance being oxidized is a reducing agent because it causes another substance to be reduced.

Reduction is the process where a substance gains electrons, resulting in a decrease in its oxidation number.
A substance being reduced is an oxidizing agent because it causes another substance to be oxidized.

Metals higher in the activity series are more easily oxidized.
Metals lower in the activity series are more easily reduced.
The metal that needs to be oxidized must be higher in the series than the metal being reduced.
The reading on a voltmeter will be lower than $2. 3$ , but higher than $z jY$ .

Example: $Cl$ is reduced, and something else (Mnt) is oxidized. $Mnt Mnt$ gains $11 t I 1$
$2H^+ + 2e^- \rightarrow H<em>2$ is an example of reduction. $H</em>2$ is being reduced.
Oxidation number: $2H^+ \text{is reduced to } H_2$
The total number of electrons lost must equal the total number of electrons gained in all redox reactions.
Number of electrons lost = Number of electrons gained.
Number of electrons for element oxidized = number of electrons for element reduced.