Electrolysis

1. Overview

Electrolysis is a fundamental process in chemistry used to decompose ionic compounds into their constituent elements using electricity. This process is essential for the industrial extraction of reactive metals (like aluminum), the production of useful chemicals like chlorine, and for protecting or decorating metals through electroplating.

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Key Definitions

• Electrolysis: The decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current.
• Electrolyte: The molten or aqueous substance that undergoes electrolysis; it contains moving ions which carry the charge.
• Anode: The positive electrode.
• Cathode: The negative electrode.
• Anion: A negatively charged ion (attracted to the anode).
• Cation: A positively charged ion (attracted to the cathode).
• Inert Electrode: Electrodes that do not react with the electrolyte or products (usually graphite/carbon or platinum).

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Core Content

The Electrolytic Cell

General Rules for Electrode Products

1. Metals or Hydrogen are always formed at the cathode (negative electrode).
2. Non-metals (other than hydrogen) are always formed at the anode (positive electrode).

Electrolysis of Molten Lead(II) Bromide ($PbBr_2$)

• Electrolyte: Molten lead(II) bromide, $PbBr_2(l)$
• Observations at Anode (+): Brown fumes of bromine gas are evolved.
• Observations at Cathode (-): A silver bead of molten lead is formed.
• Word Equation: lead(II) bromide → lead + bromine
• Symbol Equation: $PbBr_2(l) \rightarrow Pb(l) + Br_2(g)$

Electrolysis of Concentrated Aqueous Sodium Chloride ($NaCl$)

• Electrolyte: Concentrated $NaCl(aq)$ (Brine).
• Observations at Anode (+): Bubbles of pale green gas (chlorine).
• Observations at Cathode (-): Bubbles of colorless gas (hydrogen).
• Note: Sodium remains in the solution as part of sodium hydroxide ($NaOH$).

Electrolysis of Dilute Sulfuric Acid ($H_2SO_4$)

• Electrolyte: $H_2SO_4(aq)$
• Observations at Anode (+): Bubbles of colorless gas (oxygen).
• Observations at Cathode (-): Bubbles of colorless gas (hydrogen).
• Volume Ratio: The volume of hydrogen produced is twice the volume of oxygen ($2:1$ ratio).

Electroplating

• Purpose: To improve appearance (e.g., silver-plated jewelry) and to increase resistance to corrosion (e.g., chromium-plated car parts).
• Process:
  1. The cathode is the object to be plated (e.g., a spoon).
  2. The anode is the pure metal you want to plate with (e.g., silver).
  3. The electrolyte is an aqueous solution of a soluble salt of the plating metal (e.g., silver nitrate).

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Extended Content (Extended Curriculum Only)

Transfer of Charge

• External Circuit: Electrons flow from the positive terminal of the battery to the anode, and from the cathode to the negative terminal.
• Electrolyte: Charge is carried by the movement of ions. Cations ($+$) move to the cathode; Anions ($-$) move to the anode.
• At Electrodes:
  • Anode: Oxidation occurs (loss of electrons). $2Cl^-(aq) \rightarrow Cl_2(g) + 2e^-$
  • Cathode: Reduction occurs (gain of electrons). $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$

Electrolysis of Aqueous Copper(II) Sulfate ($CuSO_4$)

1. Using Inert Electrodes (Carbon/Graphite):
   • Cathode: Pink-brown solid copper coats the electrode.
   • Anode: Bubbles of oxygen gas.
   • Observation: The blue color of the solution fades as $Cu^{2+}$ ions are removed.
2. Using Copper Electrodes:
   • Cathode: Increases in mass as copper is deposited.
   • Anode: Decreases in mass as copper atoms dissolve to become ions.
   • Observation: The blue color of the solution remains constant.

Predicting Products in Aqueous Solutions

When multiple ions are present, the "Ease of Discharge" rules apply:

• At the Cathode: The less reactive element is discharged. (Usually $H^+$ vs. a metal; if the metal is above Hydrogen in the reactivity series, $H_2$ gas forms).
• At the Anode:
  • In concentrated halides (e.g., $NaCl$), the halide ion ($Cl^-$, $Br^-$, $I^-$) is discharged.
  • In dilute solutions or solutions with sulfates/nitrates, oxygen is discharged from $OH^-$ ions.

How to write a half-equation Half-equations show what happens at one electrode only. Follow these steps:

1. Write the ion that arrives at the electrode.
2. Add or remove electrons to balance the charge.
3. At the cathode, ions gain electrons (reduction): $Cu^{2+} + 2e^- \rightarrow Cu$
4. At the anode, ions lose electrons (oxidation): $2Br^- \rightarrow Br_2 + 2e^-$
5. Check: atoms balanced? Charges balanced? If both sides add up, you are correct.

The trickiest half-equation is oxygen formation from hydroxide ions. Four OH⁻ ions combine, losing 4 electrons, and producing water as a by-product: $4OH^- \rightarrow O_2 + 2H_2O + 4e^-$. This appears at the anode whenever the solution is dilute or contains sulfate/nitrate ions.

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Key Equations

Process	Location	Half-Equation
Reduction	Cathode	$M^{n+} + ne^- \rightarrow M$
Oxidation	Anode	$2X^- \rightarrow X_2 + 2e^-$
Hydrogen formation	Cathode	$2H^+(aq) + 2e^- \rightarrow H_2(g)$
Oxygen formation	Anode	$4OH^-(aq) \rightarrow O_2(g) + 2H_2O(l) + 4e^-$
Lead formation	Cathode	$Pb^{2+}(l) + 2e^- \rightarrow Pb(l)$
Chlorine formation	Anode	$2Cl^-(aq) \rightarrow Cl_2(g) + 2e^-$

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Common Mistakes to Avoid

• ❌ Wrong: Saying electrons flow through the electrolyte.
  • ✓ Right: Electrons only flow through the wires; ions carry the charge through the electrolyte.
• ❌ Wrong: Confusing the charges of the electrodes.
  • ✓ Right: Use the mnemonic PANIC: Positive Anode, Negative Is Cathode.
• ❌ Wrong: Forgetting that hydrogen is diatomic ($H_2$) and oxygen is diatomic ($O_2$).
  • ✓ Right: Always write them as $H_2(g)$ and $O_2(g)$ in equations.

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Exam Tips

• Command Words:
  • If asked to "State" an observation, write what you see (e.g., "bubbles," "colorless gas," "pink solid"). Do not just name the gas.
  • If asked to "Describe" the electrolysis, mention the movement of ions, the observations at both electrodes, and the names of the products.
• The 2:1 Ratio: In the electrolysis of water or dilute sulfuric acid, there is always twice as much Hydrogen gas (at the cathode) as Oxygen gas (at the anode) because the formula is $H_2O$.
• Common Contexts: Expect questions about the extraction of Aluminum (using cryolite) or the purification of copper.
• State Symbols: Ensure molten substances are $(l)$ and dissolved substances are $(aq)$. Forgetting these is a common way to lose easy marks.

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Exam-Style Questions

Practice these original exam-style questions to test your understanding. Each question mirrors the style, structure, and mark allocation of real Cambridge 0620 Theory papers.

Exam-Style Question 1 — Short Answer [5 marks]

Question:

Electrolysis is used in the extraction of aluminium from its ore, bauxite.

(a) Define electrolysis. [2]

(b) Aluminium oxide is dissolved in molten cryolite before electrolysis. State two reasons why this is necessary. [2]

(c) State at which electrode aluminium is produced during the electrolysis of aluminium oxide. [1]

Worked Solution:

(a)

1. Electrolysis is the decomposition of a compound... Electrolysis is the decomposition of an ionic compound. [This specifies the process]
2. ...when molten or in solution... ...when molten or in aqueous solution, by the passage of an electric current. [Completes the definition]

How to earn full marks:

• Mention breaking down a compound (1 mark)
• Mention using electricity AND being molten or in solution (1 mark)

(b)

1. Cryolite lowers the melting point of aluminium oxide... Cryolite lowers the melting point of aluminium oxide. [This reduces the energy needed for electrolysis]
2. ...and increases its conductivity. Cryolite increases the conductivity of the electrolyte. [This allows the current to flow more easily]

How to earn full marks:

• Lowering the melting point (1 mark)
• Increasing conductivity (1 mark)

(c)

1. Aluminium is a metal, so it is produced at the cathode. Cathode

How to earn full marks:

• Correct electrode (1 mark)

Common Pitfall: Remember that electrolysis only works if the compound is molten or dissolved. Solid ionic compounds don't conduct electricity because the ions are fixed in place. Also, cryolite isn't just a solvent; it plays a crucial role in reducing the energy needed for the process.

Exam-Style Question 2 — Short Answer [6 marks]

Question:

A student investigates the electrolysis of copper(II) sulfate solution using inert carbon electrodes.

(a) State the colour change observed in the electrolyte during the electrolysis. [1]

(b) Write the ionic half-equation for the reaction that occurs at the cathode. [2]

(c) Name the gas evolved at the anode. [1]

(d) Explain why the mass of the anode remains unchanged during the electrolysis. [2]

Worked Solution:

(a)

1. Copper(II) sulfate is blue and is being used up. The blue colour fades.

How to earn full marks:

• Mention the blue colour fading or becoming lighter (1 mark)

(b)

1. Copper ions are reduced at the cathode... $Cu^{2+}(aq) + 2e^- \longrightarrow Cu(s)$ [This shows the correct ions and electrons]

How to earn full marks:

• Correct reactants and products (1 mark)
• Correct balancing and state symbols (1 mark)

(c)

1. Oxygen is produced at the anode. Oxygen

How to earn full marks:

• Correct gas (1 mark)

(d)

1. The electrodes are inert... The electrodes are made of carbon (graphite) which is inert. [State the electrodes don't react]
2. ...so they are not involved in the reaction. Therefore, the carbon atoms do not react or dissolve during electrolysis. [Explain the consequence]

How to earn full marks:

• State that the electrodes are inert (1 mark)
• State that the carbon does not react (1 mark)

Common Pitfall: Many students forget that in the electrolysis of aqueous solutions, water can also be electrolyzed. In this case, with copper(II) sulfate, the copper ions are preferentially discharged at the cathode. At the anode, oxygen is produced from the water, not from the sulfate ions.

Exam-Style Question 3 — Extended Response [9 marks]

Question:

An electroplating company plates iron spoons with silver to improve their appearance.

(a) Describe the process of electroplating an iron spoon with silver. In your answer, state the electrolyte used, the electrode materials, and explain what happens at each electrode. [5]

(b) Suggest two reasons why a company might choose to electroplate an object. [2]

(c) The same company also plates iron objects with chromium. The same current is used for both silver and chromium electroplating. Given that the relative atomic mass of silver is 108 and chromium is 52, and that silver ions ($Ag^+$) require one electron for reduction while chromium ions ($Cr^{3+}$) require three, predict which metal will deposit a greater mass in the same time. Show your reasoning. [2]

Worked Solution:

(a)

1. The iron spoon is made the cathode (negative electrode)... The iron spoon is made the cathode (negative electrode). [The object to be plated is always the cathode]
2. ...and a silver bar is made the anode (positive electrode). A silver bar is made the anode (positive electrode). [The plating metal is always the anode]
3. The electrolyte is a solution containing silver ions... The electrolyte is a solution containing silver ions, such as silver nitrate ($AgNO_3$). [The electrolyte must contain ions of the plating metal]
4. At the cathode, silver ions are reduced to silver metal... At the cathode, silver ions ($Ag^+$) are reduced to silver metal ($Ag$) and deposited on the spoon: $Ag^+ + e^- \longrightarrow Ag$. [Reduction occurs at the cathode]
5. ...and at the anode, silver metal is oxidised to silver ions. At the anode, silver metal is oxidised to silver ions ($Ag^+$), replenishing the electrolyte: $Ag \longrightarrow Ag^+ + e^-$. [Oxidation occurs at the anode, replacing the ions lost at the cathode]

How to earn full marks:

• Iron spoon as cathode (1 mark)
• Silver bar as anode (1 mark)
• Silver nitrate (or other suitable silver salt) as electrolyte (1 mark)
• Reduction of silver ions at cathode (1 mark)
• Oxidation of silver metal at anode (1 mark)

(b)

1. Electroplating improves appearance... To improve appearance (e.g., to make it look more attractive). [Aesthetic reasons]
2. ...and provides resistance to corrosion. To provide resistance to corrosion (e.g., to prevent rusting). [Functional reasons]

How to earn full marks:

• Improve appearance (1 mark)
• Increase resistance to corrosion (1 mark)

(c)

1. For silver, 1 mole requires 1 mole of electrons, so 108g is deposited per mole of electrons. For silver, 1 mole requires 1 mole of electrons, so $\boxed{108 \text{ g}}$ is deposited per mole of electrons. [Consider the mass of silver per mole of electrons]
2. For chromium, 1 mole requires 3 moles of electrons, so 52g is deposited per 3 moles of electrons, or 52/3 = 17.3g per mole of electrons. For chromium, 1 mole requires 3 moles of electrons, so $\boxed{52 \text{ g}}$ is deposited per 3 moles of electrons, or $\boxed{17.3 \text{ g}}$ is deposited per mole of electrons. [Consider the mass of chromium per mole of electrons]
3. Therefore, silver will deposit a greater mass. Therefore, silver will deposit a greater mass in the same time.

How to earn full marks:

• Correctly relate moles of electrons to mass of silver deposited (1 mark)
• Correctly relate moles of electrons to mass of chromium deposited AND conclude silver deposits more (1 mark)

Common Pitfall: In electroplating, remember the object you want to plate is always the cathode. Also, when comparing the mass of different metals deposited during electrolysis, make sure you account for the number of electrons each ion requires for discharge. It's not just about the relative atomic mass!

Exam-Style Question 4 — Extended Response [10 marks]

Question:

A student investigates the electrolysis of aqueous copper(II) chloride, $CuCl_2(aq)$, using carbon electrodes.

(a) State the products formed at the cathode and the anode. [2]

(b) Write the ionic half-equations for the reactions occurring at the cathode and the anode. [2]

(c) Describe the observations the student would make at each electrode during the electrolysis. [2]

(d) If the student replaced the carbon electrodes with copper electrodes, predict how the observations at each electrode would change. Explain your answer. [4]

Worked Solution:

(a)

1. Copper(II) chloride contains $Cu^{2+}$ and $Cl^-$ ions. Copper is produced at the cathode... Cathode: Copper [Metals are produced at the cathode]
2. ...and chlorine is produced at the anode. Anode: Chlorine [Non-metals (other than hydrogen) are produced at the anode]

How to earn full marks:

• Copper at the cathode (1 mark)
• Chlorine at the anode (1 mark)

(b)

1. At the cathode, copper ions gain electrons... Cathode: $Cu^{2+}(aq) + 2e^- \longrightarrow Cu(s)$ [This shows the correct reduction reaction]
2. ...and at the anode, chloride ions lose electrons. Anode: $2Cl^-(aq) \longrightarrow Cl_2(g) + 2e^-$ [This shows the correct oxidation reaction]

How to earn full marks:

• Correct cathode half-equation (1 mark)
• Correct anode half-equation (1 mark)

(c)

1. At the cathode, a brown solid is deposited... Cathode: A brown solid (copper) is deposited on the electrode. [Describe the copper formed]
2. ...and at the anode, a green/yellow gas is evolved. Anode: A green/yellow gas (chlorine) is evolved. [Describe the chlorine formed]

How to earn full marks:

• Brown solid deposited at the cathode (1 mark)
• Green/yellow gas evolved at the anode (1 mark)

(d)

1. At the cathode, copper is still formed. Cathode: The observation would remain the same, as copper ions are still reduced to copper metal. [Explain that reduction still takes place at the cathode]
2. At the anode, the copper electrode dissolves. Anode: The copper electrode would dissolve and the electrolyte would turn blue. [Explain that oxidation takes place at the anode]
3. ...as copper atoms are oxidised to copper ions. ...as copper atoms are oxidised to copper ions ($Cu(s) \longrightarrow Cu^{2+}(aq) + 2e^-$) replenishing the $Cu^{2+}$ ions in the solution. [Explaining the oxidation of copper into copper ions]
4. ...so no chlorine is produced. Therefore, no chlorine gas is produced. [Explaining that chlorine is no longer produced]

How to earn full marks:

• Cathode observation remains the same (1 mark)
• Anode dissolves (1 mark)
• Electrolyte turns blue (1 mark)
• No chlorine produced (1 mark)

Common Pitfall: When the electrodes are made of the same metal as the electrolyte's cation (like copper in copper(II) chloride), the anode dissolves instead of producing a gas. This is because it's energetically easier to oxidize the copper electrode than to oxidize the chloride ions. Always consider the electrode material when predicting electrolysis products.