Electrolysis
1. Principles of Electrolysis
Learning outcomes
- I can explain the purpose of electrolysis.
- I can describe how electrical energy drives chemical change.
- I can distinguish between electrolysis and galvanic cells.
- I can identify oxidation and reduction during electrolysis.
- I can explain the energy transformations involved.
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.
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
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.
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.
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.
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.
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
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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
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.
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.
Comparing Electrolytic and Galvanic Cells
| Feature | Electrolytic Cell | Galvanic Cell |
|---|---|---|
| Reaction | Non-spontaneous | Spontaneous |
| Main purpose | Use electricity to cause chemical change | Produce electricity from chemical change |
| Energy conversion | Electrical → chemical | Chemical → electrical |
| External power source | Required | Not required to drive the reaction |
| Anode reaction | Oxidation | Oxidation |
| Cathode reaction | Reduction | Reduction |
| Anode sign | Positive | Negative |
| Cathode sign | Negative | Positive |
The most important rule to remember is:
oxidation always occurs at the anode
and:
reduction always occurs at the cathode.
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.
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.
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.
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.
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
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.
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.
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.
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
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.
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.
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.
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
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.
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.
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.
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
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.
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.
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.
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.
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.
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.
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.
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.
Common Mistake: Electrons Move Through the Electrolyte
Electrons mainly travel through the:
external electrical circuit.
Ions transport charge through the:
electrolyte.
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.
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.
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.
Check Your Understanding
- Define electrolysis.
- What is the main purpose of electrolysis?
- Why does electrolysis require an external power supply?
- What type of chemical reaction is driven by electrolysis?
- What are the main components of an electrolytic cell?
- What is an electrolyte?
- Why must ions in an electrolyte be able to move?
- Why does solid sodium chloride not conduct electricity through ionic movement?
- Why does molten sodium chloride conduct electricity?
- What is the cathode?
- What is the anode?
- What is the charge of the cathode in an electrolytic cell?
- What is the charge of the anode in an electrolytic cell?
- Where do cations move?
- Where do anions move?
- Where does reduction occur?
- Where does oxidation occur?
- What does OIL RIG mean?
- What does AnOx mean?
- What does RedCat mean?
- Write the half-equation for the reduction of Na⁺.
- Write the half-equation for the oxidation of Cl⁻.
- Describe the electrolysis of molten sodium chloride.
- What products are formed during the electrolysis of molten NaCl?
- What products can be formed during the electrolysis of water?
- Why is approximately twice as much hydrogen as oxygen produced from water?
- What is the overall energy transformation during electrolysis?
- Where is some of the supplied electrical energy stored?
- Compare electrolysis with a galvanic cell.
- Which type of cell involves a spontaneous reaction?
- Which type of cell requires an external electrical supply?
- What happens to energy in a galvanic cell?
- What happens to energy in an electrolytic cell?
- Why do the signs of the electrodes differ between electrolytic and galvanic cells?
- Explain why oxidation still occurs at the anode in both types of cells.
- Explain why reduction still occurs at the cathode in both types of cells.
- What is electroplating?
- Why is the object being electroplated usually the cathode?
- Give three reasons why objects are electroplated.
- How can electrolysis be used to purify copper?
- Why is electrolysis used to extract very reactive metals?
- Write the reduction half-equation for Al³⁺.
- Predict which electrode Mg²⁺ moves toward.
- Predict which electrode Br⁻ moves toward.
- Describe the difference between electron movement and ion movement in an electrolytic cell.
- Explain why electrolysis is an example of energy conversion.
- Explain how electrolysis can be used to store renewable energy as hydrogen.
- Compare an electrolytic cell and a galvanic cell in terms of energy, spontaneity, electrode signs, oxidation, and reduction.
- Explain why the statement "the anode is always positive" is incorrect.
- Describe the complete process of electrolysis from the electrical power supply to the formation of products at the electrodes.
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.
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.