Electrolysis

Purpose of the electrolysis required practical

To observe and identify products formed at electrodes

Type of electrolysis in required practical

Aqueous electrolysis

Electrolyte used in required practical

Copper sulfate solution

Formula of copper sulfate

CuSO₄

Electrodes used

Graphite (inert) electrodes

Why electrodes must be inert

So they do not react with the electrolyte

Power supply used

DC power supply

Set up of electrolysis practical

Two graphite electrodes placed in copper sulfate solution and connected to power supply

What happens when power is switched on

Ions move towards electrodes

Cathode charge

Negative

Anode charge

Positive

Ion moving to cathode

Cu²⁺

Ion moving to anode

OH⁻

What happens at the cathode

Copper ions gain electrons

Cathode half equation

Cu²⁺ + 2e⁻ → Cu

Observation at cathode

Pink/brown copper solid forms

What happens at the anode

Hydroxide ions lose electrons

Anode half equation

4OH⁻ → O₂ + 2H₂O + 4e⁻

Observation at anode

Bubbles of gas

Gas produced at anode

Oxygen

Test for oxygen

Relights a glowing splint

Why copper forms instead of hydrogen

Copper is less reactive than hydrogen

Why sulfate ions do not react

Sulfate ions are very stable

Main hazard in electrolysis practical

Production of gases

Safety precaution

Wear eye protectionMolten electrolysis

Why compounds must be molten

So ions are free to move

Ions present in molten electrolysis

Only ions from the compound

Difference between molten and aqueous electrolysis

Molten has no water ions

Molten sodium chloride ions

Na⁺ and Cl⁻

Product at cathode in molten NaCl

Sodium metal

Cathode half equation for molten NaCl

Na⁺ + e⁻ → Na

Product at anode in molten NaCl

Chlorine gas

Anode half equation for molten NaCl

2Cl⁻ → Cl₂ + 2e⁻

Why sodium forms in molten electrolysis

No hydrogen present

Why molten electrolysis is dangerous

Produces very reactive metalsReactivity series

Top of reactivity series

Potassium

Order of common metals

Potassium, Sodium, Calcium, Magnesium, Aluminium, Zinc, Iron, Copper, Silver, Gold

Metals above carbon in reactivity series

Extracted by electrolysis

Metals below carbon in reactivity series

Extracted by reduction with carbon

Why aluminium uses electrolysis

More reactive than carbon

Why iron does not use electrolysis

Less reactive than carbon

Hydrogen position in reactivity series

Between aluminium and zinc

Metals above hydrogen

React with acids

Metals below hydrogen

Do not react with acidsIons present in aqueous solutions

Rule at cathode in aqueous electrolysis

Less reactive metal or hydrogen produced

Rule at anode in aqueous electrolysis

Halide gives halogen, otherwise oxygen

Why hydrogen often forms

Hydrogen is less reactive than many metals

Why oxygen forms at anode

OH⁻ ions are oxidisedElectrolysis

Ionic compound

A compound made of positive and negative ions

Electrolyte

A molten or aqueous ionic compound that conducts electricity

Why ionic compounds conduct when molten

Ions are free to move

Why ionic compounds do not conduct when solid

Ions are fixed in a lattice

Cation

A positively charged ion

Anion

A negatively charged ion

Cathode

The negative electrode where reduction occurs

Anode

The positive electrode where oxidation occurs

Reduction

Gain of electrons

Oxidation

Loss of electrons

Mnemonic for oxidation and reduction

OIL RIG (Oxidation Is Loss, Reduction Is Gain)

Mnemonic for electrodes

RED CAT, OX AN

Movement of ions in electrolysis

Cations move to the cathode, anions move to the anode

What happens at the cathode

Reduction occurs

What happens at the anode

Oxidation occurs

Metal ions at the cathode

Gain electrons to form metal atoms

Example cathode half equation

Cu²⁺ + 2e⁻ → Cu

Non-metal ions at the anode

Lose electrons to form non-metal molecules

Example anode half equation

2Cl⁻ → Cl₂ + 2e⁻

Why products form at electrodes

Ions gain or lose electrons

State of products at electrodes

Solid metals or gases

Why electrodes must conduct electricity

To allow electrons to flow

Why aluminium is extracted by electrolysis

It is too reactive to be reduced by carbon

Ore of aluminium

Bauxite

Main aluminium compound

Aluminium oxide (Al₂O₃)

Why aluminium oxide has a high melting point

Strong ionic bonds

Purpose of cryolite in aluminium extraction

Lowers the melting point

Benefit of cryolite

Reduces energy required

Cathode reaction for aluminium

Al³⁺ + 3e⁻ → Al

Anode reaction for aluminium

O²⁻ → O₂ + 4e⁻

Why carbon anodes wear away

Oxygen reacts with carbon to form CO₂

Main disadvantage of aluminium extraction

Uses a lot of electricity

Environmental issue with aluminium extraction

CO₂ emissions and high energy use

Aqueous solution

Ionic compound dissolved in water

Ions present in aqueous electrolysis

Ions from compound plus H⁺ and OH⁻

Rule for cathode in aqueous electrolysis

Less reactive metal or hydrogen is produced

Hydrogen half equation

2H⁺ + 2e⁻ → H₂

Rule for anode in aqueous electrolysis

Halides form halogen, otherwise oxygen

Chloride ions at anode

2Cl⁻ → Cl₂ + 2e⁻

Hydroxide ions at anode

4OH⁻ → O₂ + 2H₂O + 4e⁻

Why oxygen forms instead of sulfate

Sulfate ions are very stable

Electrolysis of sodium chloride solution

Hydrogen at cathode, chlorine at anode

Why sodium is not formed

Hydrogen is less reactive than sodium

Purpose of electrolysis required practical

To identify products at electrodes

Electrolyte used in practical

Copper sulfate solution

Electrodes used in practical

Inert graphite electrodes

Observation at cathode

Copper metal deposited

Observation at anode

Oxygen gas produced

Test for oxygen

Relights a glowing splint

Test for hydrogen

Squeaky pop with a lit splint

Why graphite electrodes are used

They are inert

Main safety risk

Gas production