What Is Electrolysis?

Electrolysis is the use of electrical energy to drive a chemical reaction that would not normally occur spontaneously.

The word can be thought of as:

electro = electricity

lysis = splitting or breaking apart

During electrolysis, an external source of electricity forces chemical changes to occur.

The overall energy transformation is:

electrical energy → chemical energy

Some energy is also transferred to the surroundings as:

thermal energy.

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Why Is Electrolysis Needed?

Some chemical reactions occur spontaneously.

Others do not.

A non-spontaneous reaction requires an external energy supply before it can proceed.

Electrolysis provides this energy using:

electricity.

For example, electrical energy can be used to:

• decompose compounds
• produce elements
• extract reactive metals
• purify metals
• electroplate objects
• produce industrial chemicals
• generate hydrogen from water

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The Electrolytic Cell

Electrolysis takes place in an:

electrolytic cell.

A simple electrolytic cell contains:

• an external power supply
• two electrodes
• an electrolyte
• connecting wires

The power supply provides the electrical energy needed to drive the:

non-spontaneous chemical reaction.

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The Electrolyte

The electrolyte is a substance containing mobile ions.

These ions must be able to move.

Therefore electrolysis generally requires an ionic substance that is:

molten

or:

dissolved in water.

In these states, the ions are free to move toward the electrodes.

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Why Solid Ionic Compounds Do Not Conduct

In a solid ionic compound, positive and negative ions are held in a:

fixed crystal lattice.

The ions cannot move freely.

Therefore the solid cannot normally conduct electricity through ionic movement.

When the substance is melted or dissolved:

the ions become mobile.

The electrolyte can then conduct electricity.

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The Electrodes

The two conductors placed in the electrolyte are called:

electrodes.

They are:

anode

and

cathode.

During electrolysis:

Anode = positive electrode

Cathode = negative electrode

This is important because the electrode signs are different from those in a spontaneous galvanic cell.

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The Cathode

In an electrolytic cell, the cathode is connected to the:

negative terminal of the power supply.

Electrons are supplied to the cathode.

Positive ions are attracted toward it.

Positive ions are called:

cations.

At the cathode, cations can:

gain electrons.

Gaining electrons is:

reduction.

Therefore:

Cathode = Reduction

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The Anode

In an electrolytic cell, the anode is connected to the:

positive terminal of the power supply.

Negative ions are attracted toward it.

Negative ions are called:

anions.

At the anode, particles can:

lose electrons.

Losing electrons is:

oxidation.

Therefore:

Anode = Oxidation.

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Remember: AnOx and RedCat

A useful memory aid is:

AnOx

Anode = Oxidation

and:

RedCat

Reduction = Cathode

These rules apply to both:

electrolytic cells and galvanic cells.

What changes between the two types of cells is the sign of each electrode.

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Ion Movement During Electrolysis

Opposite charges attract.

Therefore:

Cations (+) → Cathode (−)

and:

Anions (−) → Anode (+)

This movement of ions through the electrolyte allows charge to be transported through the:

solution or molten substance.

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Electrons and Ions Follow Different Paths

It is important to distinguish between:

electron movement

and:

ion movement.

Electrons move through:

• wires
• electrodes
• the external electrical circuit

Ions move through:

• molten electrolytes
• electrolyte solutions

Therefore:

electrons move through conductors

while:

ions move through the electrolyte.

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Oxidation and Reduction

Electrolysis always involves:

redox reactions.

Remember:

Oxidation = loss of electrons

Reduction = gain of electrons

Another useful memory aid is:

OIL RIG

Oxidation Is Loss

Reduction Is Gain

The "loss" and "gain" refer to:

electrons.

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Reduction at the Cathode

Suppose copper ions are present:

Cu²⁺

At the cathode, a copper ion can gain two electrons:

Cu²⁺ + 2e⁻ → Cu

The copper ion has gained electrons.

Therefore it has been:

reduced.

Solid copper is deposited at the cathode.

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Oxidation at the Anode

Suppose chloride ions are present:

Cl⁻

At the anode:

2Cl⁻ → Cl₂ + 2e⁻

The chloride ions lose electrons.

Therefore they have been:

oxidized.

Chlorine gas is produced.

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Electron Transfer

Electrolysis therefore involves electron transfer.

At the cathode:

particles gain electrons

At the anode:

particles lose electrons

The external power supply drives this movement of charge and forces the chemical reaction to:

continue.

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Example: Molten Sodium Chloride

Consider molten sodium chloride:

NaCl

When molten, it contains mobile:

Na⁺ ions

and:

Cl⁻ ions.

The ions move toward oppositely charged electrodes.

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At the Cathode: Sodium

Sodium ions move toward the:

negative cathode.

Each sodium ion gains an electron:

Na⁺ + e⁻ → Na

This is:

reduction.

Sodium metal is produced.

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At the Anode: Chlorine

Chloride ions move toward the:

positive anode.

Chloride ions lose electrons:

2Cl⁻ → Cl₂ + 2e⁻

This is:

oxidation.

Chlorine gas is produced.

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Overall Electrolysis of Molten Sodium Chloride

The two half-equations are:

Cathode:

2Na⁺ + 2e⁻ → 2Na

Anode:

2Cl⁻ → Cl₂ + 2e⁻

Combining them gives:

2NaCl → 2Na + Cl₂

Electrical energy is required to drive this:

non-spontaneous decomposition reaction.

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Why Must Sodium Chloride Be Molten?

Solid sodium chloride contains ions, but they are locked into a:

crystal lattice.

They cannot move to the electrodes.

When sodium chloride melts:

Na⁺ and Cl⁻ ions become mobile.

The molten compound can therefore conduct electricity and undergo:

electrolysis.

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Example: Electrolysis of Water

Electricity can also be used to decompose water.

The overall reaction can be represented as:

2H₂O → 2H₂ + O₂

Electrical energy is supplied to produce:

hydrogen gas and oxygen gas.

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Products of Water Electrolysis

During water electrolysis:

hydrogen gas forms at the cathode.

oxygen gas forms at the anode.

The volume of hydrogen produced is approximately:

twice the volume of oxygen.

This matches the overall equation:

2H₂O → 2H₂ + O₂

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Electrolysis and Hydrogen Production

Electrolysis can be used to produce hydrogen as an:

energy carrier.

The energy transformation is:

electrical energy → chemical energy stored in hydrogen

Later, hydrogen can be used in a fuel cell:

chemical energy → electrical energy.

This creates a possible energy-storage pathway.

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Electrolysis and Galvanic Cells

Electrolytic cells and galvanic cells both involve:

• electrodes
• electrolytes
• electron transfer
• oxidation
• reduction
• redox reactions

However, they differ in a fundamental way.

A galvanic cell uses a spontaneous chemical reaction to:

produce electrical energy.

An electrolytic cell uses electrical energy to:

force a non-spontaneous chemical reaction.

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Comparing Electrolytic and Galvanic Cells

The most important rule to remember is:

oxidation always occurs at the anode

and:

reduction always occurs at the cathode.

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Why Do the Electrode Signs Change?

In a galvanic cell, the chemical reaction itself releases electrons at the anode.

Therefore the anode acts as the:

negative electrode.

In an electrolytic cell, an external power supply pulls electrons away from the anode and pushes electrons toward the cathode.

Therefore:

electrolytic anode = positive

electrolytic cathode = negative.

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Energy Transformations

Energy transformation is one of the most important differences between the two systems.

Galvanic Cell

Chemical energy → electrical energy

The spontaneous reaction releases energy.

Electrolytic Cell

Electrical energy → chemical energy

Electrical energy forces the non-spontaneous reaction.

Some energy is also transferred into:

thermal energy.

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Electrolysis Does Not Create Energy

Electrolysis does not create energy.

The electrical energy supplied by the power source is transformed.

Some becomes:

chemical potential energy

stored in the products.

Some is transferred as:

heat.

This follows the:

law of conservation of energy.

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Electrolysis as Energy Storage

Consider water electrolysis.

Electricity is supplied:

electrical energy

↓

Water is decomposed

↓

Hydrogen is produced

↓

Energy is stored chemically in hydrogen.

The hydrogen can later be used in a fuel cell:

chemical energy → electrical energy.

This means electrolysis can be part of an:

energy-storage system.

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Electroplating

One important application of electrolysis is:

electroplating.

Electroplating uses electricity to deposit a thin layer of one metal onto another object.

Examples include:

• silver plating
• gold plating
• nickel plating
• chromium plating
• copper plating

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How Electroplating Works

The object being coated is usually connected as the:

cathode.

Metal ions in the electrolyte move toward the cathode.

At the cathode they gain electrons:

Mⁿ⁺ + ne⁻ → M

Metal atoms are deposited onto the surface.

This forms a thin:

metal coating.

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Why Electroplate Objects?

Electroplating can be used to:

• improve appearance
• reduce corrosion
• improve wear resistance
• provide useful surface properties
• reduce the amount of expensive metal required

For example, an object can receive a thin layer of gold without being made entirely from:

gold.

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Metal Purification

Electrolysis can also be used to purify metals.

Copper is an important example.

Impure copper can be used as the:

anode.

Pure copper is deposited onto the:

cathode.

This process is called:

electrorefining.

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Why Is Pure Copper Important?

Copper is widely used for:

electrical wiring.

Impurities can affect its electrical conductivity.

Electrolytic purification can produce very pure copper suitable for:

• electrical cables
• electronics
• electrical equipment

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Extracting Reactive Metals

Some metals are too reactive to be extracted easily by chemical reduction with carbon.

These include metals such as:

• aluminum
• sodium
• magnesium

Electrolysis can be used to extract these metals from:

molten ionic compounds.

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Aluminum Production

Aluminum is produced industrially using electrolysis of aluminum-containing compounds.

Electrical energy is required to reduce aluminum ions:

Al³⁺ + 3e⁻ → Al

Because aluminum production requires large amounts of electricity, the cost and environmental impact of the electricity supply are:

important considerations.

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Electrolysis in Industry

Electrolysis is used on a very large scale.

Important applications include:

• aluminum production
• chlorine production
• hydrogen production
• sodium hydroxide production
• copper purification
• electroplating

Therefore electrolysis is an important connection between:

chemistry, electricity, and industry.

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Predicting Ion Movement

To predict where an ion will move, first identify its:

charge.

Positive ion:

cation → cathode

Negative ion:

anion → anode

For example:

Cu²⁺ → cathode

Na⁺ → cathode

Cl⁻ → anode

Br⁻ → anode

O²⁻ → anode

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Predicting the Reaction

Once the electrode is identified, determine whether the particle must gain or lose electrons.

At the cathode:

gain electrons → reduction

At the anode:

lose electrons → oxidation

This gives a useful sequence:

Identify charge → identify electrode → identify electron transfer → identify oxidation or reduction.

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Worked Example 1

Where will Mg²⁺ ions move?

Mg²⁺ is:

positively charged.

Therefore it moves toward the:

negative cathode.

At the cathode:

Mg²⁺ + 2e⁻ → Mg

This is:

reduction.

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Worked Example 2

Where will Br⁻ ions move?

Br⁻ is:

negatively charged.

Therefore it moves toward the:

positive anode.

At the anode:

2Br⁻ → Br₂ + 2e⁻

This is:

oxidation.

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Worked Example 3

Consider molten magnesium chloride:

MgCl₂

The ions present are:

Mg²⁺ and Cl⁻.

At the cathode:

Mg²⁺ + 2e⁻ → Mg

At the anode:

2Cl⁻ → Cl₂ + 2e⁻

Overall:

MgCl₂ → Mg + Cl₂

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Worked Example 4

A student says:

"Reduction occurs at the negative electrode."

Is this always correct?

No.

Reduction always occurs at the:

cathode.

In an electrolytic cell the cathode is negative.

In a galvanic cell the cathode is positive.

Therefore it is safer to remember:

RedCat: Reduction at Cathode.

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Worked Example 5

A student says:

"The anode is always positive."

Is this correct?

No.

The anode is positive in an:

electrolytic cell.

The anode is negative in a:

galvanic cell.

However:

oxidation always occurs at the anode.

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Worked Example 6

Why does molten sodium chloride conduct electricity while solid sodium chloride does not?

Both contain ions.

However, in solid sodium chloride the ions are:

fixed in position.

When molten, the ions can:

move freely.

Moving ions transport charge through the electrolyte.

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Worked Example 7

What energy transformation occurs when hydrogen is produced from water by electrolysis?

Electrical energy → chemical energy

The chemical energy is stored in the products, including:

hydrogen.

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Worked Example 8

A renewable-energy system uses excess solar electricity to electrolyze water.

During the day:

solar energy → electrical energy → chemical energy in hydrogen

Later, a fuel cell can convert:

chemical energy → electrical energy.

Electrolysis therefore provides a way of converting electrical energy into a:

storable chemical form.

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Common Mistake: Electrolysis Produces Electricity

Electrolysis does not primarily produce electricity.

It:

uses electricity.

The external power supply drives a chemical reaction.

A galvanic cell does the opposite:

it produces electrical energy from a spontaneous chemical reaction.

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Common Mistake: The Anode Is Always Negative

The anode is not defined by its charge.

The anode is defined as the electrode where:

oxidation occurs.

In electrolysis:

anode = positive.

In a galvanic cell:

anode = negative.

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Common Mistake: The Cathode Is Always Positive

The cathode is defined as the electrode where:

reduction occurs.

In electrolysis:

cathode = negative.

In a galvanic cell:

cathode = positive.

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Common Mistake: Electrons Move Through the Electrolyte

Electrons mainly travel through the:

external electrical circuit.

Ions transport charge through the:

electrolyte.

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Common Mistake: Solid Ionic Compounds Can Always Be Electrolyzed

The ions in a solid ionic compound cannot move freely.

The compound usually needs to be:

molten or dissolved

before its ions can move and participate effectively in electrolysis.

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Common Mistake: Oxidation Means Adding Oxygen

Some oxidation reactions involve oxygen, but the more general definition is:

oxidation = loss of electrons.

Similarly:

reduction = gain of electrons.

These definitions are especially useful in:

electrochemistry.

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Common Mistake: Electrolysis Destroys Energy

Energy is not destroyed.

Electrical energy is transformed into:

chemical energy and thermal energy.

The total amount of energy is conserved.

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Check Your Understanding

1. Define electrolysis.
2. What is the main purpose of electrolysis?
3. Why does electrolysis require an external power supply?
4. What type of chemical reaction is driven by electrolysis?
5. What are the main components of an electrolytic cell?
6. What is an electrolyte?
7. Why must ions in an electrolyte be able to move?
8. Why does solid sodium chloride not conduct electricity through ionic movement?
9. Why does molten sodium chloride conduct electricity?
10. What is the cathode?
11. What is the anode?
12. What is the charge of the cathode in an electrolytic cell?
13. What is the charge of the anode in an electrolytic cell?
14. Where do cations move?
15. Where do anions move?
16. Where does reduction occur?
17. Where does oxidation occur?
18. What does OIL RIG mean?
19. What does AnOx mean?
20. What does RedCat mean?
21. Write the half-equation for the reduction of Na⁺.
22. Write the half-equation for the oxidation of Cl⁻.
23. Describe the electrolysis of molten sodium chloride.
24. What products are formed during the electrolysis of molten NaCl?
25. What products can be formed during the electrolysis of water?
26. Why is approximately twice as much hydrogen as oxygen produced from water?
27. What is the overall energy transformation during electrolysis?
28. Where is some of the supplied electrical energy stored?
29. Compare electrolysis with a galvanic cell.
30. Which type of cell involves a spontaneous reaction?
31. Which type of cell requires an external electrical supply?
32. What happens to energy in a galvanic cell?
33. What happens to energy in an electrolytic cell?
34. Why do the signs of the electrodes differ between electrolytic and galvanic cells?
35. Explain why oxidation still occurs at the anode in both types of cells.
36. Explain why reduction still occurs at the cathode in both types of cells.
37. What is electroplating?
38. Why is the object being electroplated usually the cathode?
39. Give three reasons why objects are electroplated.
40. How can electrolysis be used to purify copper?
41. Why is electrolysis used to extract very reactive metals?
42. Write the reduction half-equation for Al³⁺.
43. Predict which electrode Mg²⁺ moves toward.
44. Predict which electrode Br⁻ moves toward.
45. Describe the difference between electron movement and ion movement in an electrolytic cell.
46. Explain why electrolysis is an example of energy conversion.
47. Explain how electrolysis can be used to store renewable energy as hydrogen.
48. Compare an electrolytic cell and a galvanic cell in terms of energy, spontaneity, electrode signs, oxidation, and reduction.
49. Explain why the statement "the anode is always positive" is incorrect.
50. Describe the complete process of electrolysis from the electrical power supply to the formation of products at the electrodes.

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Key Terms

Electrolysis: Use of electrical energy to drive a non-spontaneous chemical reaction.

Electrolytic cell: Electrochemical cell in which electrical energy drives chemical change.

Electrolyte: Substance containing mobile ions that can transport charge.

Electrode: Conductor through which electrons enter or leave an electrochemical system.

Anode: Electrode where oxidation occurs.

Cathode: Electrode where reduction occurs.

Cation: Positively charged ion attracted toward the cathode during electrolysis.

Anion: Negatively charged ion attracted toward the anode during electrolysis.

Oxidation: Loss of electrons.

Reduction: Gain of electrons.

Redox reaction: Reaction involving both oxidation and reduction.

Non-spontaneous reaction: Reaction that requires an external energy input to proceed.

Half-equation: Equation showing electron transfer during either oxidation or reduction.

Electroplating: Use of electrolysis to deposit a layer of material, usually a metal, onto an object.

Electrorefining: Use of electrolysis to purify a metal.

Electrolysis of water: Use of electrical energy to decompose water, producing hydrogen and oxygen.

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Key Takeaways

• Electrolysis uses electrical energy to drive a non-spontaneous chemical reaction.
• The overall energy transformation is mainly electrical energy → chemical energy.
• An electrolytic cell requires an external power supply.
• The electrolyte must contain mobile ions.
• In an electrolytic cell, the anode is positive and the cathode is negative.
• Cations move toward the cathode.
• Anions move toward the anode.
• Oxidation always occurs at the anode.
• Reduction always occurs at the cathode.
• The rules AnOx and RedCat work for both electrolytic and galvanic cells.
• Electrons travel through the external circuit, while ions transport charge through the electrolyte.
• Electrolysis and galvanic cells involve opposite overall energy transformations.
• A galvanic cell converts chemical energy → electrical energy.
• An electrolytic cell converts electrical energy → chemical energy.
• Electrolysis can be used for metal extraction, electroplating, metal purification, chemical production, and hydrogen production.
• Electrolysis of water can convert electrical energy into chemical energy stored in hydrogen.
• Electrolysis is therefore important in chemistry, industry, materials processing, and emerging energy-storage systems.