Oxidation-Reduction (Redox) Reactions

Oxidation and reduction reactions are crucial in chemistry, with applications ranging from batteries to corrosion prevention.
Originally, oxidation meant combining with oxygen, and reduction meant losing oxygen; now, the definitions are broader and interdependent.

Oxidation is defined as the loss of electrons.
Reduction is defined as the gain of electrons.
Redox reactions involve both reduction and oxidation, which must occur together.

Oxidation numbers track electron loss or gain by atoms in a reaction.
Oxidation is the loss of electrons and an increase in oxidation number.
Reduction is the gain of electrons and a decrease in oxidation number.
Mnemonic: LEO says GER (Loss of Electrons is Oxidation; Gain of Electrons is Reduction).
Rules for assigning oxidation numbers:
1. Uncombined elements have an oxidation number of zero.
2. Monatomic ions have an oxidation number equal to their ionic charge.
3. Group 1 metals are always +1, and Group 2 metals are always +2 in compounds.
4. Fluorine is always -1 in compounds; other halogens are -1 when most electronegative.
5. Hydrogen is +1 in compounds, unless with a metal, where it is -1.
6. Oxygen is usually -2 in compounds, but +2 when combined with fluorine.
7. The sum of oxidation numbers in a compound is zero.
8. The sum of oxidation numbers in polyatomic ions equals the ion's charge.

To identify redox reactions, assign oxidation numbers to atoms on both sides of the equation.
A change in oxidation number indicates a redox reaction.
If an uncombined element appears on one side and is in a compound on the other, it's a redox reaction.
Double replacement reactions are not redox reactions.
The atom with an increased oxidation number is oxidized; the atom with a decreased oxidation number is reduced.

Half-reactions show either the oxidation or reduction portion of a redox reaction, including electrons gained or lost.
Reduction half-reactions show an atom or ion gaining electrons, decreasing its oxidation number.
Oxidation half-reactions show an atom or ion losing electrons, increasing its oxidation number.
Half-reactions must conserve mass and charge.

Electrochemical cells involve chemical reactions and electron flow.
Voltaic cells use spontaneous chemical reactions to produce electron flow.
Electrolytic cells require an electric current to force nonspontaneous reactions.
Electrodes are surfaces where oxidation or reduction occurs.
The anode is where oxidation occurs.
The cathode is where reduction occurs.

In voltaic cells, chemical energy spontaneously converts to electrical energy.
A salt bridge connects two containers, allowing ion flow to complete the circuit.
Electrons flow from the anode to the cathode through a wire.
The Activity series can identify the anode (higher on the chart, oxidized) and cathode (lower on the chart, reduced).
Remember RED CAT (Reduction occurs at the CAThode) and AN OX (ANode is the site of Oxidation).
$Eº<em>{cell} = Eº</em>{reduction} - Eº_{oxidation}$

Electrolytic cells use electricity to force nonspontaneous chemical reactions (electrolysis).
Electrolysis can obtain active elements from fused salts or electroplate metals.

Voltaic cells have spontaneous redox reactions; electrolytic cells have nonspontaneous reactions.
In voltaic cells, the anode is negative, and the cathode is positive.
In electrolytic cells, the anode is positive, and the cathode is negative.
Both use redox reactions, with oxidation at the anode and reduction at the cathode, and electrons flow from anode to cathode.