Electrolysis - Chemistry - Pearson IGCSE

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Welcome to the World of Electrolysis!

Hello future Chemists! This chapter, Electrolysis, is one of the most practical and fascinating topics in Chemistry. It explains how we can use electricity to force chemical reactions to happen—reactions that wouldn't normally occur on their own!

We will learn the essential definitions, how the setup works, and the crucial rules that determine which products form, especially in industrial processes like extracting metals and making household chemicals.

Key Definitions You Must Know

Before we dive into the process, let's nail down the vocabulary. These are the building blocks of the entire chapter:

Electrolysis: The decomposition (breaking down) of a compound using electricity. (Electro = electricity, lysis = breaking.)
Electrolyte: The liquid (molten ionic compound or aqueous solution) that conducts electricity because it contains free-moving ions.
Electrodes: Solid conductors (usually metal or graphite) dipped into the electrolyte. They carry the current into and out of the electrolyte.
Cations: Positively charged ions (\(Na^+\), \(H^+\), \(Mg^{2+}\)). They move towards the negative electrode.
Anions: Negatively charged ions (\(Cl^-\), \(OH^-\), \(SO_4^{2-}\)). They move towards the positive electrode.

Quick Review: Conduction

Remember, electricity only flows if charged particles can move:

Metals: Conduct electricity because of moving free electrons.
Electrolytes: Conduct electricity because of moving free ions.

The Setup: The Electrolytic Cell

Electrolysis happens in a vessel called an electrolytic cell. It requires two electrodes connected to a power source (like a battery).

Understanding the Electrodes

The battery dictates the charge of the electrodes, which in turn attracts the oppositely charged ions:

The Cathode: This is the negative electrode (\(–\)) because it is connected to the negative terminal of the power supply.

It attracts Cations (\(+\)).
The Anode: This is the positive electrode (\(+\)) because it is connected to the positive terminal of the power supply.

It attracts Anions (\(–\)).

Memory Aid: P.A.N.I.C.

Positive Anode Negative Is Cathode. (This helps you remember the charges in an electrolytic cell!)

What Happens at the Electrodes? (Redox)

When ions arrive at the electrodes, they either gain or lose electrons. This is a Redox reaction (Reduction-Oxidation).

1. At the Cathode (Reduction):

Cations (\(+\)) arrive. They need electrons to become neutral atoms.
Reduction always happens at the Cathode (Red Cat).
Ions gain electrons: \(\text{Metal}^+ + e^- \rightarrow \text{Metal}\) (e.g., \(\text{Na}^+ + e^- \rightarrow \text{Na}\))

2. At the Anode (Oxidation):

Anions (\(–\)) arrive. They have extra electrons and must lose them to become neutral atoms or molecules.
Oxidation always happens at the Anode (Ox An).
Ions lose electrons: \(\text{Non-metal}^- \rightarrow \text{Non-metal} + e^-\) (e.g., \(2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-\))

Key Takeaway

Electrolysis is a redox process driven by electricity. At the cathode, metal/hydrogen is usually formed via reduction (gain of electrons). At the anode, non-metals (like halogens or oxygen) are usually formed via oxidation (loss of electrons).

Case 1: Electrolysis of Molten Compounds

This is the simplest type of electrolysis because there are only two ions present (the metal cation and the non-metal anion).

Imagine molten lead bromide, \(\text{PbBr}_2\). The only ions present are \(\text{Pb}^{2+}\) and \(\text{Br}^-\).

Electrode	Ions Attracted	Reaction Type	Product	Half-Equation
Cathode (–)	\(\text{Pb}^{2+}\) (Cation)	Reduction (Gain \(e^-\))	Molten Lead metal	\(\text{Pb}^{2+} + 2e^- \rightarrow \text{Pb}\)
Anode (+)	\(\text{Br}^-\) (Anion)	Oxidation (Lose \(e^-\))	Bromine gas	\(2\text{Br}^- \rightarrow \text{Br}_2 + 2e^-\)

Did You Know?

Molten salts often require extremely high temperatures (over \(500^\circ\text{C}\)) to stay liquid, which is why industrial processes often look for ways to lower this temperature (like adding Cryolite in aluminium extraction, discussed below).

Case 2: Electrolysis of Aqueous Solutions (The Tricky Part!)

When a compound is dissolved in water, the electrolysis gets complicated because water itself ionizes slightly, introducing two extra ions: \(\text{H}^+\) (from the acid part of water) and \(\text{OH}^-\) (from the base part of water).

We now have four ions competing to be discharged:

The metal cation and \(\text{H}^+\) are competing at the Cathode.
The non-metal anion and \(\text{OH}^-\) are competing at the Anode.

The ion that is preferentially discharged (reacts) is the one that requires less energy.

Rule 1: At the Cathode (Cations competing: \(\text{Metal}^+\) vs. \(\text{H}^+\))

The least reactive ion is discharged (reduced).

We look at the Reactivity Series (K, Na, Ca, Mg, Al, C, Zn, Fe, Sn, Pb, H, Cu, Ag, Au).

If the metal is MORE reactive than Hydrogen (e.g., Na, K, Mg):

The metal ion stays in solution. Hydrogen gas is produced at the cathode.

\(2\text{H}^+ + 2e^- \rightarrow \text{H}_2\)
If the metal is LESS reactive than Hydrogen (e.g., Cu, Ag):

The metal ion is discharged. The pure metal is deposited on the cathode.

\(\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}\)

Analogy: Imagine the ions are running a race. The least reactive ion is the "fastest" to accept electrons, so it wins and gets discharged first. Highly reactive metals are "lazy" and stay dissolved.

Rule 2: At the Anode (Anions competing: \(\text{Halide}/\text{Sulfate}/\text{Nitrate}\) vs. \(\text{OH}^-\))

This rule depends on the type of negative ion present:

If the solution contains Halide ions (\(\text{Cl}^-\), \(\text{Br}^-\), \(\text{I}^-\)):

The Halogen is usually discharged, forming a gas (\(\text{Cl}_2\), \(\text{Br}_2\), \(\text{I}_2\)).

\(2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-\)
If the solution contains Sulfate (\(\text{SO}_4^{2-}\)) or Nitrate (\(\text{NO}_3^-\)) ions:

These complex ions are rarely discharged. Instead, Hydroxide ions (\(\text{OH}^-\)) from the water are discharged, producing Oxygen gas and water.

\(4\text{OH}^- \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4e^-\)

Common Mistake Alert!

Do not forget the state of the product. Metals (Cu, Ag) are solids. Hydrogen, Chlorine, and Oxygen are gases (bubbles!).

Industrial Applications of Electrolysis

Electrolysis is vital for manufacturing essential chemicals and extracting reactive metals.

1. Extraction of Aluminium (from Bauxite)

Aluminium metal is used in planes, cars, and foil. Since aluminium is highly reactive (above Hydrogen in the series), it must be extracted using electrolysis.

The Challenge: Aluminium oxide (\(\text{Al}_2\text{O}_3\)), which is extracted from bauxite ore, has an extremely high melting point (over \(2000^\circ\text{C}\)).

The Solution: Aluminium oxide is dissolved in molten Cryolite.

Why Cryolite? It lowers the melting point of the mixture to around \(900^\circ\text{C}\), saving huge amounts of energy and money.
Electrolyte: Molten mixture of \(\text{Al}_2\text{O}_3\) and Cryolite.
Electrodes: Graphite (Carbon) is used for both.

Products and Reactions:

At the Cathode (\(–\)): \(\text{Al}^{3+}\) ions are reduced to liquid aluminium metal, which sinks to the bottom of the cell.
\(\text{Al}^{3+} + 3e^- \rightarrow \text{Al}\)
At the Anode (\(+\)): \(\text{O}^{2-}\) ions are oxidized to oxygen gas.
\(2\text{O}^{2-} \rightarrow \text{O}_2 + 4e^-\)

Crucial Point about Aluminium Extraction

The oxygen gas produced at the carbon anodes reacts immediately with the hot carbon, forming carbon dioxide gas (\(\text{C} + \text{O}_2 \rightarrow \text{CO}_2\)). This means the carbon anodes are continuously worn away and must be replaced regularly, which adds to the cost of the process.

2. Electrolysis of Brine (Concentrated Sodium Chloride Solution)

Brine (concentrated aqueous sodium chloride, \(\text{NaCl}\)) is electrolysed to produce three extremely useful products: Chlorine, Hydrogen, and Sodium Hydroxide. This is often called the Chlor-alkali process.

Ions Present: \(\text{Na}^+\), \(\text{Cl}^-\), \(\text{H}^+\), \(\text{OH}^-\)