Unit 6: Comprehensive Summary

A redox reaction involves the transfer of electrons.

• Oxidation is the loss of electrons, gain of oxygen, loss of hydrogen, or an increase in oxidation number.
• Reduction is the gain of electrons, loss of oxygen, gain of hydrogen, or a decrease in oxidation number.

Oxidation numbers are used to track electron transfer. An oxidising agent (oxidant) accepts electrons and gets reduced. A reducing agent (reductant) donates electrons and gets oxidised.

Half-equations are used to show the oxidation and reduction processes separately. They can be combined to form a full, balanced redox equation by ensuring the number of electrons lost equals the number gained.

A galvanic cell (voltaic cell) converts chemical energy into electrical energy via a spontaneous redox reaction. It consists of two half-cells connected by a wire and a salt bridge.

• The anode is the negative electrode where oxidation occurs.
• The cathode is the positive electrode where reduction occurs.

The Standard Electrode Potential ($E^\circ$) measures the tendency of a half-cell to be reduced, relative to the Standard Hydrogen Electrode (SHE) which is assigned a value of 0.00 V. A more positive $E^\circ$ indicates a stronger oxidising agent, while a more negative $E^\circ$ indicates a stronger reducing agent.

The cell potential is calculated as $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$.

Primary cells are non-rechargeable (e.g., dry cells), while secondary cells are rechargeable (e.g., lead-acid batteries). Fuel cells run continuously by being supplied with external reactants.

Electrolysis converts electrical energy into chemical energy, using an external power source to drive a non-spontaneous redox reaction in an electrolytic cell.

• The anode is the positive electrode where oxidation occurs.
• The cathode is the negative electrode where reduction occurs.

In the electrolysis of molten salts, the metal is formed at the cathode and the non-metal at the anode. In aqueous solutions, water may be oxidised or reduced in preference to the ions from the salt, depending on reactivity and concentration.

Applications include metal extraction (e.g., aluminium), purification, and electroplating. Faraday's constant (96,500 C/mol) relates the amount of charge passed to the moles of substance produced.

Rusting is the electrochemical corrosion of iron, requiring both oxygen and water. Iron is oxidised at the anodic region, and oxygen is reduced at the cathodic region. The presence of electrolytes (salts, acids) accelerates the process.

Prevention methods include:

• Barrier Methods: Painting, oiling, or plastic coating to physically block out air and water.
• Sacrificial Protection: Connecting iron to a more reactive metal (e.g., zinc, magnesium). The more reactive metal acts as a sacrificial anode and corrodes instead of the iron. Galvanising is the process of coating iron with zinc, providing both barrier and sacrificial protection.