What is electrolysis and how does it work in electrochemistry?

Electrolysis is a chemical process used in electrochemistry where electrical energy is used to drive a non-spontaneous chemical reaction. It involves passing an electric current through an electrolyte, causing the movement of ions towards the electrodes, where oxidation and reduction reactions occur. This process is fundamental in the IB MYP Science curriculum for understanding how chemical changes can be induced by electricity.

What are the main components required for an electrolysis setup?

An electrolysis setup typically requires an electrolyte, which is a substance containing free ions, and two electrodes: an anode and a cathode. The anode is the positive electrode where oxidation occurs, and the cathode is the negative electrode where reduction takes place. A power source is also necessary to provide the electrical energy needed to drive the reaction.

How is electrolysis used in real-world applications?

Electrolysis has several practical applications, including the extraction of metals from ores, such as aluminum and copper, and the electroplating of metals to improve their appearance and resistance to corrosion. It is also used in the production of chemical compounds like chlorine and hydrogen gas. Understanding these applications is important for students studying electrochemistry in the IB MYP Science curriculum.

What is the role of electrodes in the process of electrolysis?

In electrolysis, electrodes serve as the sites where oxidation and reduction reactions occur. The anode attracts anions (negatively charged ions) and facilitates their oxidation, while the cathode attracts cations (positively charged ions) and facilitates their reduction. The choice of electrode material can affect the efficiency and outcome of the electrolysis process, making it a key concept in electrochemistry studies.

Why is electrolysis considered a non-spontaneous reaction?

Electrolysis is considered a non-spontaneous reaction because it requires an external source of energy, in the form of electricity, to proceed. Unlike spontaneous reactions, which occur naturally and release energy, electrolysis needs energy input to drive the chemical changes. This distinction is crucial for understanding the principles of electrochemistry in the IB MYP Science curriculum.

Why is "Mixing up cathode and anode polarity." a common mistake in Electrolysis?

Why it happens: Confusing with batteries/galvanic cells where polarity is reversed. How to avoid it: In ELECTROLYSIS, the supply DRIVES current → cathode is connected to the NEGATIVE terminal of the supply. Use: cAThode = cATtacks cATions (positive ions).

Why is "Saying Na metal forms at the cathode in aqueous NaCl." a common mistake in Electrolysis?

Why it happens: Na⁺ is the only cation from the salt. How to avoid it: Sodium is TOO REACTIVE — H⁺ (from water) is discharged instead, producing H₂. Sodium can only be obtained by electrolysing MOLTEN NaCl.

Why is "Writing 'Cu²⁺ + 2e⁻ → Cu' at the anode." a common mistake in Electrolysis?

Why it happens: Slip on direction. How to avoid it: At ANODE, electrons FLOW OUT into the circuit — so the equation has electrons on the RIGHT (atom → ion + e⁻). At CATHODE, electrons IN — ion + e⁻ → atom.

Electrolysis cell schematic

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Electrolysis — IB MYP Sciences: using electricity to split ionic compounds and electroplate metals

Electrolysis is the use of electricity to break compounds apart or coat objects with metal. This MYP Sciences note covers what happens at each electrode, the products from molten and aqueous electrolytes, and the rule that lets you predict the products in any case.

At a glance

Electrolysis: electric current passed through a molten or dissolved ionic compound to break it down.
The substance being electrolysed is the ELECTROLYTE.
Two electrodes: ANODE (+) and CATHODE (−).
Cations (+ ions) move to the cathode; anions (− ions) move to the anode.
At the cathode: REDUCTION (cation + electrons → atom).
At the anode: OXIDATION (anion → atom + electrons).
Molten ionic compounds: simple — metal at cathode, non-metal at anode.
Aqueous solutions: H+ and OH− also present; less reactive species discharged at each electrode.
Electroplating: coating an object with a thin metal layer (e.g. silver-plating a spoon).

What you’ll learn

Mapped to the IB MYP Sciences subject guide (2026 onwards).

MYP Sciences A — Define electrolysis, electrolyte, cathode, anode.
MYP Sciences A — Predict the products of molten and aqueous electrolysis.
MYP Sciences A — Write half-equations for cathode and anode reactions.
MYP Sciences A — Describe electroplating and a real-world use.
MYP Sciences D — Discuss the energy cost and environmental impact of industrial electrolysis.

Electrolysis setup

Battery + two electrodes dipped in a molten or dissolved ionic compound.

An electrolysis cell has three parts:

Electrolyte: the ionic compound, MOLTEN or DISSOLVED in water (ions must be free to move).
Electrodes: two conducting rods (usually carbon or platinum) dipped in the electrolyte.
DC power supply: a battery or DC source, making one electrode positive (ANODE) and the other negative (CATHODE).

Electrolysis of molten ionic compounds

Easy: metal at cathode, non-metal at anode.

Molten electrolytes contain ONLY the ions from the compound — no water. So the products are simple:

Molten lead bromide (PbBr₂):

At cathode (−): Pb²⁺ + 2 e⁻ → Pb. Lead metal forms.
At anode (+): 2 Br⁻ → Br₂ + 2 e⁻. Bromine gas (orange-brown vapour).

Molten sodium chloride (NaCl):

At cathode: Na⁺ + e⁻ → Na. Sodium metal.
At anode: 2 Cl⁻ → Cl₂ + 2 e⁻. Chlorine gas.

The metal is always discharged at the cathode and the non-metal at the anode. This is how very reactive metals (K, Na, Ca, Mg, Al) are extracted from their compounds.

Molten: only the compound's ions present.
Metal forms at cathode; non-metal at anode.
Used to extract reactive metals (Na from NaCl, Al from Al₂O₃).

Electrolysis of aqueous solutions

Water adds H⁺ and OH⁻ to the mix. The LESS reactive species is discharged at each electrode.

Water is slightly ionised: H₂O ⇌ H⁺ + OH⁻. So in an aqueous electrolyte, both the salt's ions AND H⁺/OH⁻ are present at the electrodes.

Rule at the cathode: the LESS REACTIVE species is discharged.

If the cation is LESS reactive than H⁺ (e.g. Cu²⁺, Ag⁺) → the metal forms.
If the cation is MORE reactive than H⁺ (e.g. Na⁺, K⁺, Ca²⁺, Mg²⁺) → H₂ gas forms instead.

Rule at the anode: again, less reactive discharged.

Halide ions (Cl⁻, Br⁻, I⁻) are usually discharged → halogen gas.
Otherwise OH⁻ (from water) is discharged → O₂ gas.

Examples:

Solution	Cathode	Anode
NaCl (aq)	H₂ (Na too reactive)	Cl₂
CuSO₄ (aq)	Cu (less reactive than H⁺)	O₂ (no halide)
H₂SO₄ (aq)	H₂	O₂
Dilute Na₂SO₄ (aq)	H₂	O₂ (essentially electrolysing water)

CuSO₄ is a great teaching example: blue colour (Cu²⁺) fades as the copper deposits at the cathode, and a coating of bright orange copper builds up.

Aqueous: salt ions + H⁺/OH⁻ from water.
Cathode rule: less reactive species discharged (Cu²⁺ → Cu, but Na⁺ stays and H₂ comes off).
Anode rule: halides discharged first; otherwise O₂ from OH⁻.

Electroplating

Deposit a thin layer of a metal on an object — for decoration or protection.

Electroplating is electrolysis used to put a thin layer of one metal onto an object. The object to be plated is the CATHODE; the plating metal is the ANODE; the electrolyte contains ions of the plating metal.

For silver-plating a spoon:

Cathode (−): the spoon.
Anode (+): a piece of pure silver.
Electrolyte: silver nitrate solution (AgNO₃).

What happens:

At the cathode (spoon): Ag⁺ + e⁻ → Ag. Silver coats the spoon.
At the anode (silver): Ag → Ag⁺ + e⁻. Silver dissolves into solution.
Net effect: silver moves from the anode to the spoon. The AgNO₃ solution stays roughly the same concentration.

Uses include:

Decoration: jewellery (gold/silver-plated steel).
Protection: chrome-plating bumpers and tools.
Improving conductivity: electroplating copper onto electrical contacts.

MYP criterion D extension: why is electroplating preferred over making things from solid silver or gold? Cost — a thin layer needs only a tiny amount of the expensive metal, while the bulk is cheap steel underneath.

Object to plate = cathode. Plating metal = anode. Electrolyte = metal-ion solution.
Metal moves from anode to cathode.
Used for decoration, corrosion protection and contact conductivity.

Quick recap

Electrolysis = breaking down an ionic compound using electricity.
Cathode (−): reduction. Anode (+): oxidation.
Molten: metal at cathode, non-metal at anode.
Aqueous: less reactive species discharged at each electrode.
Electroplating: object = cathode; plating metal = anode; metal-ion solution.

Memorise this

Verbatim phrases and definitions MYP criterion-A markschemes credit.

Cathode (−): cations move there; REDUCTION happens
Anode (+): anions move there; OXIDATION happens
Molten ionic: metal at cathode, non-metal at anode
Aqueous: cathode → less reactive species (Cu wins over Na)
Aqueous anode: halide if present, else O₂ from OH⁻

How it’s examined

Criterion A and C are the main tests. Predict products of a molten or aqueous electrolyte, write balanced half-equations, and describe electroplating setups.

Sources: IB MYP Sciences guide (IBO, official subject guide). Last reviewed 2026-05-25.

Step-by-step worked examples — Electrolysis

Step-by-step solutions to past-paper-style questions on electrolysis, written exactly the way a tutor would explain them at the board.

1Molten lead bromide
Getting started• molten, half-equations
Question
Molten PbBr₂ is electrolysed using inert electrodes. Write half-equations for each electrode and name the products.
Step-by-step solution
1. Step 1
  Pb²⁺ goes to cathode (−). Half-equation: Pb²⁺ + 2e⁻ → Pb.
2. Step 2
  Br⁻ goes to anode (+). Half-equation: 2Br⁻ → Br₂ + 2e⁻.
3. Step 3
  Products: lead metal at cathode, bromine vapour at anode.
Answer
Cathode: Pb²⁺ + 2e⁻ → Pb (lead). Anode: 2Br⁻ → Br₂ + 2e⁻ (bromine).
2Aqueous CuSO₄
Building confidence• aqueous
Question
Predict the products of electrolysing aqueous CuSO₄ with inert electrodes.
Step-by-step solution
1. Step 1
  At cathode: Cu²⁺ vs H⁺. Cu is LESS reactive than H, so Cu wins → Cu metal deposited.
2. Step 2
  At anode: SO₄²⁻ vs OH⁻. SO₄²⁻ contains S already at max oxidation — OH⁻ discharged → 4OH⁻ → 2H₂O + O₂ + 4e⁻.
3. Step 3
  Products: copper at cathode, oxygen at anode.
Answer
Cathode: copper metal. Anode: oxygen gas. (Blue Cu²⁺ solution fades.)
3Aqueous NaCl
Building confidence• aqueous
Question
Predict the products of electrolysing concentrated brine (aqueous NaCl).
Step-by-step solution
1. Step 1
  At cathode: Na⁺ vs H⁺. Na is FAR more reactive than H — H⁺ is discharged → H₂ gas.
2. Step 2
  At anode: Cl⁻ vs OH⁻. Cl⁻ is high concentration; halides usually discharge first → Cl₂ gas.
3. Step 3
  NaOH remains in solution as the third useful product.
Answer
Cathode: H₂ (Na too reactive). Anode: Cl₂. Solution: NaOH.
Examiner tip
This is the chlor-alkali industrial process — produces three valuable products from cheap brine.
4Setting up an electroplating cell
Stretch• electroplating
Question
A student wants to silver-plate a steel spoon. Describe the setup: electrode positions, electrolyte and what happens.
Step-by-step solution
1. Step 1
  Spoon: cathode (−), connected to the negative terminal.
2. Step 2
  Pure silver: anode (+), connected to the positive terminal.
3. Step 3
  Electrolyte: silver nitrate solution (AgNO₃), so Ag⁺ ions are available in solution.
4. Step 4
  Action: Ag → Ag⁺ + e⁻ at anode (silver dissolves); Ag⁺ + e⁻ → Ag at cathode (silver coats spoon). The AgNO₃ solution stays roughly constant.
Answer

Electrolysis – Study Notes & Past Paper Style Questions | IB MYP Sciences (Years 1-5) | Tutopiya

Electrolysis

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Electrolysis — frequently asked questions

Electrolysis — IB MYP Sciences: using electricity to split ionic compounds and electroplate metals

At a glance

What you’ll learn

Electrolysis setup

Electrolysis of molten ionic compounds

Electrolysis of aqueous solutions

Electroplating

Quick recap

Memorise this

How it’s examined

Step-by-step worked examples — Electrolysis

1Molten lead bromide

2Aqueous CuSO₄

3Aqueous NaCl

4Setting up an electroplating cell

Key Definitions and Keywords — Electrolysis

Electrolysis

Electrolyte

Cathode

Anode

Electroplating

Common Mistakes and Misconceptions — Electrolysis

✕Mixing up cathode and anode polarity.

✕Saying Na metal forms at the cathode in aqueous NaCl.

✕Writing 'Cu²⁺ + 2e⁻ → Cu' at the anode.