Basic Principles and Terms of Electrolysis for the ESAT

Introduction to Electrolysis

Electrolysis is a chemical process where electrical energy is used to drive a non-spontaneous chemical reaction. This typically involves the decomposition of a compound into its constituent elements. Unlike spontaneous chemical reactions that release energy, electrolysis requires an external power source, such as a battery or a direct current supply, to force electrons to move through a system.

Understanding the Electrolyte

An electrolyte is a substance that contains free moving ions and can therefore conduct electricity. For electrolysis to occur, the compound must be in a state where its ions are mobile. In a solid ionic lattice, the ions are held in fixed positions by strong electrostatic attractions and cannot move to carry a charge.

As shown in the lattice structure above, the ions are locked in place. To create an electrolyte, the ionic compound must be either melted to form a molten liquid or dissolved in water to form an aqueous solution. Once the lattice is broken, the ions are free to migrate toward the electrodes when a potential difference is applied. Covalent compounds, such as ethanol or glucose, do not contain ions and therefore cannot act as electrolytes.

The Electrodes: Anode and Cathode

The electrodes are solid conductors, often made of metal or graphite, that are placed into the electrolyte to complete the circuit. There are two types of electrodes that you must distinguish between for the ESAT:

1. The Anode: This is the positive electrode. It is connected to the positive terminal of the power supply. During the process, it attracts negatively charged ions.
2. The Cathode: This is the negative electrode. It is connected to the negative terminal of the power supply. During the process, it attracts positively charged ions.

A useful mnemonic to remember the charges is PANIC: Positive Anode, Negative Is Cathode.

Ion Migration and Redox Processes

When the power supply is turned on, ions in the electrolyte move toward the electrode with the opposite charge:

• Cations are positively charged ions. They are attracted to the negative cathode. At the cathode, they gain electrons, a process known as reduction (as taught in section C5.2).
• Anions are negatively charged ions. They are attracted to the positive anode. At the anode, they lose electrons, a process known as oxidation.

Half-Equations in Electrolysis

To describe the reactions at each electrode, we use half-equations to show the movement of electrons. As seen in section C3.4, we can look at the electrolysis of molten sodium chloride as a primary example.

Example: Electrolysis of Molten Sodium Chloride ($NaCl$):

In molten sodium chloride, the ions present are $Na^{+}$ and $Cl^{-}$.

1. At the Cathode (Negative Electrode): Positive sodium ions are attracted here. Each $Na^{+}$ ion gains one electron to become a neutral sodium atom. This is a reduction reaction: $Na^{+} + e^{-} \rightarrow Na$

2. At the Anode (Positive Electrode): Negative chloride ions are attracted here. Two $Cl^{-}$ ions each lose an electron to form a neutral chlorine molecule ($Cl_{2}$). This is an oxidation reaction: $2Cl^{-} \rightarrow Cl_{2} + 2e^{-}$

Example: Electrolysis of Molten Aluminium Oxide ($Al_{2}O_{3}$):

1. At the Cathode: Aluminium ions gain three electrons each: $Al^{3+} + 3e^{-} \rightarrow Al$

2. At the Anode: Oxide ions lose electrons to form oxygen gas: $2O^{2-} \rightarrow O_{2} + 4e^{-}$