Welcome to Electrochemistry!
Have you ever wondered how shiny aluminium foil is made, or how rechargeable batteries power your phone? The secret lies in a fascinating branch of chemistry called electrochemistry. In this chapter, we will discover how an electric current can break apart chemical compounds to produce useful pure elements.
Don't worry if this topic feels a bit abstract at first! By breaking down the processes step-by-step, using handy memory tricks, and looking at clear visual patterns, you will master everything you need for your CCEA GCSE exam.
1. What is Electrolysis?
Electrolysis is the decomposition (breaking down) of a liquid electrolyte using a direct current (\(\text{d.c.}\)) of electricity.
Key Terms You Must Know:
• Electrolyte: A molten ionic compound or an aqueous solution containing free ions that conducts electricity and is broken down in the process.
• Electrodes: Rods (usually made of unreactive graphite/carbon or platinum) that dip into the electrolyte to carry electrical current into and out of the liquid.
• Anode: The positive electrode.
• Cathode: The negative electrode.
• Anion: A negatively charged ion (attracted to the positive anode).
• Cation: A positively charged ion (attracted to the negative cathode).
Why Must the Electrolyte Be Molten or in Solution?
In a solid ionic lattice, the ions are held tightly in fixed positions by strong ionic bonds and cannot move. Therefore, solid ionic compounds cannot conduct electricity.
When an ionic compound is melted (molten) or dissolved in water, the ionic lattice breaks down. The ions become free to move and carry electrical charge throughout the liquid.
Helpful Memory Trick: "PANIC"
To remember electrode polarities easily, remember the word PANIC:
Positive Anode, Negative Is Cathode!
And remember: Opposites attract!
• Negative ions (Anions) travel to the positive electrode (Anode).
• Positive ions (Cations) travel to the negative electrode (Cathode).
Key Takeaway: Electrolysis requires free-moving ions and a direct current. Positive ions go to the negative cathode, while negative ions go to the positive anode.
2. Electron Movement: Oxidation and Reduction
At the electrodes, ions either lose or gain electrons to become neutral atoms or molecules.
OIL RIG: The Golden Rule
To remember what happens to electrons at each electrode, use the famous mnemonic OIL RIG:
• Oxidation Is Loss of electrons
• Reduction Is Gain of electrons
What Happens at Each Electrode?
• At the Cathode (Negative Electrode): Positive ions gain electrons to become neutral atoms. Because they gain electrons, reduction always takes place at the cathode.
• At the Anode (Positive Electrode): Negative ions lose electrons to become neutral atoms or molecules. Because they lose electrons, oxidation always takes place at the anode.
Key Takeaway: Oxidation happens at the anode (loss of electrons), and reduction happens at the cathode (gain of electrons).
3. Electrolysis of Molten Compounds
When a simple binary molten compound (made of only two elements) is electrolysed, the metal always forms at the cathode and the non-metal always forms at the anode.
Example A: Molten Lead(II) Bromide (\(\text{PbBr}_2\))
When lead(II) bromide is heated until molten, it splits into lead ions (\(\text{Pb}^{2+}\)) and bromide ions (\(\text{Br}^-\)).
At the Cathode (Negative):
• Positive lead ions (\(\text{Pb}^{2+}\)) move to the cathode.
• Each lead ion gains 2 electrons to form liquid lead metal.
• Half equation: \(\text{Pb}^{2+} + 2\text{e}^- \rightarrow \text{Pb}\)
• Observation: A silvery, molten bead of lead metal forms at the cathode.
At the Anode (Positive):
• Negative bromide ions (\(\text{Br}^-\)) move to the anode.
• Two bromide ions each lose 1 electron to form a diatomic bromine molecule (\(\text{Br}_2\)).
• Half equation: \(2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^-\)
• Observation: Red-brown pungent fumes of bromine gas are observed bubbling at the anode.
Example B: Molten Lithium Chloride (\(\text{LiCl}\))
At the Cathode (Negative):
• Lithium ions gain electrons (reduction):
\(\text{Li}^+ + \text{e}^- \rightarrow \text{Li}\)
• Observation: Silvery molten metal forms.
At the Anode (Positive):
• Chloride ions lose electrons (oxidation):
\(2\text{Cl}^- \rightarrow \text{Cl}_2 + 2\text{e}^-\)
• Observation: Yellow-green gas with a sharp, pungent choking smell bubbles off.
Key Takeaway: For any molten binary salt: the metal forms at the cathode (reduction) and the non-metal forms at the anode (oxidation).
4. Industrial Extraction of Aluminium
Aluminium is more reactive than carbon, so it cannot be extracted by heating with carbon in a blast furnace. Instead, it is extracted by the electrolysis of molten aluminium oxide.
Step 1: The Raw Material
• The major ore of aluminium is bauxite.
• Bauxite is purified to form white aluminium oxide (also known as alumina, \(\text{Al}_2\text{O}_3\)).
Step 2: The Role of Cryolite
• Pure aluminium oxide has an extremely high melting point of over \(2000^\circ\text{C}\), which requires huge amounts of expensive energy to melt.
• It is dissolved in molten cryolite (an aluminium compound with a lower melting point).
• Why use cryolite?
1. It lowers the working melting point to around \(900 - 1000^\circ\text{C}\), saving massive amounts of energy and money.
2. It improves electrical conductivity.
The Electrolytic Cell Layout
• Both the anode and cathode are made of graphite (carbon).
• The cell lining acts as the cathode (negative electrode).
• Large carbon blocks dip into the molten mixture as the anodes (positive electrodes).
Reactions at the Electrodes
At the Cathode (Lining):
• Aluminium ions (\(\text{Al}^{3+}\)) are attracted to the cathode lining.
• They gain 3 electrons to form molten aluminium metal (reduction):
\(\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}\)
• Molten aluminium is denser than the electrolyte, so it sinks to the bottom of the tank and is periodically tapped off.
At the Anode (Carbon Rods):
• Oxide ions (\(\text{O}^{2-}\)) are attracted to the carbon anodes.
• They lose electrons to form oxygen gas (oxidation):
\(2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-\)
Why Must the Carbon Anodes Be Regularly Replaced?
This is a favorite exam question! At the high temperatures of the cell (\(\approx 950^\circ\text{C}\)), the oxygen gas produced at the anode reacts with the carbon/graphite anodes to produce carbon dioxide gas:
\(\text{C} + \text{O}_2 \rightarrow \text{CO}_2\)
Because of this combustion reaction, the carbon anodes gradually burn away and must be replaced regularly, adding to the operational cost.
Economic and Environmental Considerations
• Electrolysis uses massive amounts of electricity, making the extraction of new aluminium expensive.
• Recycling aluminium saves about \(95\%\) of the energy needed to extract new aluminium from bauxite, preserves natural resources, and reduces waste in landfills.
Key Takeaway: Aluminium is extracted from purified bauxite dissolved in cryolite (to lower melting point). Anodes burn away due to reaction between carbon and oxygen forming \(\text{CO}_2\).
5. Electrolysis of Dilute Sulfuric Acid
When dilute sulfuric acid (\(\text{H}_2\text{SO}_4\)) is electrolysed using inert electrodes (such as platinum or graphite), it acts essentially as the electrolysis of water.
The Ions Present
• From water: Hydrogen ions (\(\text{H}^+\)) and Hydroxide ions (\(\text{OH}^-\))
• From sulfuric acid: Hydrogen ions (\(\text{H}^+\)) and Sulfate ions (\(\text{SO}_4^{2-}\))
Electrode Reactions and Observations
At the Cathode (Negative):
• Hydrogen ions (\(\text{H}^+\)) gain electrons to form hydrogen gas (reduction):
\(2\text{H}^+ + 2\text{e}^- \rightarrow \text{H}_2\)
• Observation: Colourless bubbles of gas; gas makes a squeaky pop with a lit splint.
At the Anode (Positive):
• Hydroxide ions (\(\text{OH}^-\)) lose electrons to form oxygen gas and water (oxidation):
\(4\text{OH}^- \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^-\)
• Observation: Colourless bubbles of gas; gas relights a glowing splint.
The Volume Ratio of Gases
The chemical formula for water is \(\text{H}_2\text{O}\). During electrolysis, twice as many moles of hydrogen gas are produced as oxygen gas.
• Volume of Hydrogen : Volume of Oxygen = \(2 : 1\)
• Exactly twice as much gas collects at the cathode compared to the anode.
Key Takeaway: Electrolysis of dilute sulfuric acid produces hydrogen gas at the cathode and oxygen gas at the anode in a \(2:1\) volume ratio.
6. Common Mistakes to Avoid
• Mixing up anode and cathode: Remember PANIC (Positive Anode, Negative Is Cathode).
• Saying electrons flow through the liquid: Electrons travel only through the wires and electrodes; ions carry the current through the electrolyte.
• Forgetting diatomic molecules: Bromine (\(\text{Br}_2\)), Chlorine (\(\text{Cl}_2\)), Oxygen (\(\text{O}_2\)), and Hydrogen (\(\text{H}_2\)) exist as pairs of atoms. Make sure you balance your half equations with \(2\text{Br}^-\), \(2\text{Cl}^-\), or \(2\text{H}^+\).
• Missing the reason for replacing anodes: Anodes burn away because carbon reacts with oxygen to make carbon dioxide gas (\(\text{CO}_2\)), not just because of heat.
Quick Summary Checklist
• Electrolyte = liquid containing free ions.
• Cations (\(+\)) \(\rightarrow\) Cathode (\(-\)) [Gain electrons = Reduction].
• Anions (\(-\)) \(\rightarrow\) Anode (\(+\)) [Lose electrons = Oxidation].
• Molten lead(II) bromide \(\rightarrow\) Lead at cathode, Bromine at anode.
• Aluminium extraction: Bauxite ore, dissolved in cryolite (to lower melting point), carbon anodes burn away to form \(\text{CO}_2\).
• Dilute sulfuric acid \(\rightarrow\) Hydrogen at cathode and Oxygen at anode in a \(2:1\) ratio.