CORE Chemistry (9222) | Electrolysis

💡 Study Notes: Chemical Changes – Electrolysis 💡

Hello future chemist! Get ready to explore one of the most exciting ways to use energy in Chemistry: Electrolysis. This is where electricity is used to force a chemical reaction to happen, literally splitting compounds apart. Don't worry if this sounds complicated—we will break down the process step by step, using simple analogies to help you master this topic!

1. Understanding Electrolysis: Definitions and Setup

What is Electrolysis?

Electrolysis means "splitting with electricity." It is the process where an electric current passes through a substance, causing a non-spontaneous chemical reaction (a reaction that wouldn't happen on its own) to occur, resulting in the decomposition (breaking down) of the substance.

The Essential Components (The Electrolytic Cell)

To perform electrolysis, you need four main things:

• Power Source: Provides the direct current (DC) to push the reaction.
• Electrodes: Two solid conductors (usually made of carbon/graphite or metal) that carry the current into and out of the substance.
• Electrolyte: The substance being broken down. This must be a liquid (either molten or dissolved in water) and contain free-moving ions.
• External Circuit: The wires connecting the power source and electrodes.

Key Term Alert!

• Electrolyte: A liquid or solution containing ions that can conduct electricity (e.g., molten salt, sulfuric acid solution).
• Non-electrolyte: A substance that does not conduct electricity because it contains no free ions (e.g., pure water, sugar solution).

Analogy: Think of the electrolyte as a busy swimming pool full of people (ions). The electrodes are the diving boards where the people line up before they jump (react). The battery is the coach shouting instructions!

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2. The Key Players: Electrodes and Ions

In electrolysis, the electrodes are given special names based on their charge and what happens there. Remember: Opposites attract!

Electrodes: Anode and Cathode

The battery dictates the charge of the electrodes:

1. The electrode connected to the positive terminal of the battery is the Anode.
2. The electrode connected to the negative terminal of the battery is the Cathode.

🔥 Mnemonic 1 (P.A.N.I.C):
Positive Anode, Negative Is Cathode.

Ions: Cations and Anions

• Cations: Positively charged ions (metals or \(H^+\)). They are attracted to the negative electrode (the Cathode).
• Anions: Negatively charged ions (non-metals or groups like \(OH^-\)). They are attracted to the positive electrode (the Anode).

What Happens at the Electrodes?

When ions reach the electrodes, they either gain or lose electrons to become neutral atoms or molecules (the final product).

• At the Cathode (Negative): Cations (positive ions) arrive and gain electrons. Gaining electrons is called Reduction. (Products are usually metals or hydrogen gas.)
• At the Anode (Positive): Anions (negative ions) arrive and lose electrons. Losing electrons is called Oxidation. (Products are usually non-metals or oxygen gas.)

🔥 Mnemonic 2 (O.I.L R.I.G):
Oxidation Is Loss (of electrons).
Reduction Is Gain (of electrons).

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3. Electrolysis of Molten (Fused) Ionic Compounds

The simplest type of electrolysis is when the compound is melted (fused) but not dissolved in water.

Since no water is present, there is no competition! The products are simply the elements that make up the ionic compound.

Step-by-Step Example: Molten Lead Bromide (\(PbBr_2\))

1. Ions Present: \(Pb^{2+}\) (Cation) and \(Br^-\) (Anion).
2. At the Cathode (-): The positive lead ions move here. They gain electrons to form neutral lead metal.

   \(Pb^{2+} + 2e^- \rightarrow Pb\) (Liquid metal)

3. At the Anode (+): The negative bromide ions move here. They lose electrons to form bromine molecules (a brownish gas).

   \(2Br^- \rightarrow Br_2 + 2e^-\) (Gas)

Key Takeaway for Molten Electrolysis: You always get the metal at the cathode and the non-metal at the anode.

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4. Electrolysis of Aqueous Solutions (The Competition)

When an ionic compound is dissolved in water (an aqueous solution), the process gets slightly trickier because water itself breaks down into ions: \(H^+\) and \(OH^-\).

This means that at both electrodes, there are two types of ions competing to react:

• At the Cathode (-): Metal ions (from the salt) vs. \(H^+\) ions (from water).
• At the Anode (+): Non-metal ions (from the salt) vs. \(OH^-\) ions (from water).

A) Predicting the Product at the Cathode (Reduction)

The ion that reacts is the one that is less reactive. We use the reactivity series to decide:

• If the metal in the salt is MORE reactive than Hydrogen (e.g., Sodium, Potassium, Calcium): Hydrogen ions (\(H^+\)) react instead.

  Product: Hydrogen gas (\(H_2\)) is produced.

• If the metal in the salt is LESS reactive than Hydrogen (e.g., Copper, Silver): The metal ion reacts instead.

  Product: Pure metal is deposited on the electrode.

Don't worry if this seems tricky at first. Just check your reactivity series: if the metal is high up, you get Hydrogen instead!

B) Predicting the Product at the Anode (Oxidation)

Generally, the \(OH^-\) ions (from water) are easier to oxidize than most other anions.

• If the salt contains Sulfate (\(SO_4^{2-}\)) or Nitrate (\(NO_3^-\)) ions: The \(OH^-\) ions react instead.

  Product: Oxygen gas (\(O_2\)) is produced.

• If the salt contains Halide ions (Chloride \(Cl^-\), Bromide \(Br^-\), Iodide \(I^-\)): The halide ions are often chosen, especially if the solution is concentrated.

  Product: The Halogen element (e.g., Chlorine gas, \(Cl_2\)) is produced.

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5. Industrial Uses of Electrolysis

Electrolysis is not just a lab experiment; it is vital for many industrial processes, especially involving highly reactive elements.

A) Extraction of Reactive Metals

Metals high up in the reactivity series (like Aluminium, Sodium, Potassium) form very stable ionic bonds. We cannot use carbon (a common reducing agent) to extract them, because carbon is not reactive enough to steal the oxygen away from them.

Therefore, we must use electrolysis of their molten compounds (usually oxides or chlorides).

Did you know? Aluminium is the most abundant metal in the Earth's crust, but because it is so reactive, its extraction requires huge amounts of electrical energy, making it an expensive process.

The Hall–Héroult Process (Extraction of Aluminium):

1. Aluminium oxide (Alumina, \(Al_2O_3\)) has a very high melting point (\(2000^\circ C\)).
2. It is dissolved in molten Cryolite. Cryolite acts as a solvent and lowers the operating temperature to about \(950^\circ C\), saving energy.
3. The mixture is electrolysed using large carbon electrodes.
4. At the Cathode: Aluminium ions gain electrons to form molten Aluminium metal.

   \(Al^{3+} + 3e^- \rightarrow Al\)

5. At the Anode: Oxygen gas is produced. This oxygen reacts continuously with the hot carbon anodes, causing them to burn away (forming \(CO_2\)). The anodes must be regularly replaced.

B) Purification and Electroplating (Briefly)

• Purification: Electrolysis can be used to produce extremely pure metals (like copper), which is essential for electrical wiring. Impure metal is used as the anode, and pure metal is deposited on the cathode.
• Electroplating: This is using electrolysis to coat one metal object with a thin layer of another metal (e.g., plating cheap metal cutlery with silver or chrome for protection/appearance). The plating metal is used as the anode, and the object to be plated is the cathode.

Key Takeaway for Uses: Electrolysis is necessary for reactive metal extraction because those metals cannot be reduced using cheaper methods like heating with carbon.