Last updated: 7 October 2026
Electrochemistry may seem like several separate ideas at first. Electrolysis breaks compounds down using electricity, electroplating coats objects with metals, batteries produce electricity from chemical reactions, and fuel cells convert chemical energy directly into electrical energy. Yet all of them are built on the same underlying process: electrons are transferred from one substance to another.
Once you learn to follow those electrons, the topic becomes much more systematic. The key questions are always similar: Which particles can move? Which electrode are they attracted to? Which particle actually gains or loses electrons? And what product should we observe? We will use those questions throughout this chapter.
Try these five questions before reading the chapter. They test the particle and redox ideas that electrochemistry builds on.
- 5 quick questions
- Instant explanations
- Find misconceptions before you begin
What Is Electrolysis?
Electrolysis is the decomposition of an ionic compound, in the molten state or in aqueous solution, by the passage of electricity.
The electrical supply is not merely detecting a chemical reaction. It supplies the energy needed to force a chemical change to occur.
Electrolysis uses electrical energy to drive a chemical reaction.
Electrical energy → chemical energy
What is an electrolyte?
An electrolyte is a substance containing mobile ions that conducts electricity and is decomposed during electrolysis.
Ionic compounds conduct electricity only when their ions can move. This explains why an ionic compound can behave very differently when it is solid and when it is molten.
Solid, molten and aqueous ionic substances
Solid ionic compound: the ions are held in fixed positions within a giant ionic lattice. They cannot move through the solid, so the solid does not conduct electricity.
Molten ionic compound: the lattice has broken apart and the ions can move. The liquid conducts electricity.
Aqueous ionic solution: dissolved ions are free to move through the water, so the solution conducts electricity.
The Electrolytic Cell
An electrolytic cell must provide a path for both electrons and ions to move. It therefore contains a direct-current power supply, two electrodes and an electrolyte.
Power supply
A direct-current power supply supplies electrons to the cathode and removes electrons from the anode.
Cathode
The cathode is connected to the negative terminal of the power supply. Reduction occurs here.
Anode
The anode is connected to the positive terminal of the power supply. Oxidation occurs here.
Electrolyte
The electrolyte contains mobile cations and anions that carry charge through the liquid.
Why do ions move to different electrodes?
During electrolysis, the cathode is negatively charged and the anode is positively charged. Opposite charges attract.
| Ion | Moves towards | Possible electrode process |
|---|---|---|
| Cation | Cathode | Gains electrons and is reduced |
| Anion | Anode | Loses electrons and is oxidised |
Cathode: reduction occurs.
Anode: oxidation occurs.
The signs of the electrodes change when we later study simple chemical cells, but these definitions remain the same.
Where do electrons and ions move?
Electrons move through the external wires and conducting electrodes. Inside the electrolyte, charge is carried by moving ions.
It is therefore incorrect to imagine electrons simply travelling through the solution from one electrode to the other.
Electrolysis of Molten Ionic Compounds
Molten ionic compounds are the simplest electrolytes to analyse because water is absent. Only the ions from the ionic compound need to be considered.
In a molten binary ionic compound:
- the metal cation moves towards the cathode and gains electrons;
- the non-metal anion moves towards the anode and loses electrons.
Metal forms at the cathode.
Non-metal forms at the anode.
Example: molten sodium chloride
Molten sodium chloride contains only Na⁺ and Cl⁻ ions.
Na⁺ ions move towards the negative cathode, where they gain electrons:
Sodium ions have gained electrons, so they have been reduced.
Cl⁻ ions move towards the positive anode, where they lose electrons:
Chloride ions have lost electrons, so they have been oxidised.
The overall reaction is:
Example: molten lead(II) bromide
Molten PbBr₂ contains Pb²⁺ and Br⁻ ions.
At the cathode:
At the anode:
Notice that the number of electrons lost at the anode equals the number of electrons gained at the cathode when the half-equations are combined.
Electrolysis of Aqueous Solutions
Aqueous electrolysis requires one extra layer of reasoning because water is present. Besides the ions from the dissolved compound, H⁺ and OH⁻ associated with water must also be considered.
Several ions may therefore move towards the same electrode. However, they are not necessarily all discharged.
The process in which one competing ion is preferentially discharged is called selective discharge.
Molten versus aqueous electrolysis
Molten electrolyte: consider only the ions in the melted ionic compound.
Aqueous electrolyte: also consider H⁺ and OH⁻ from water.
Selective discharge at the cathode
Positive ions compete to gain electrons at the cathode. For the aqueous electrolytes studied at O-Level, the reactivity series provides a useful rule.
- If the metal is less reactive than hydrogen, its metal ion is usually discharged and the metal is deposited.
- If the metal is more reactive than hydrogen, hydrogen is usually produced instead.
| Competing positive ions | Product at cathode | Reason |
|---|---|---|
| Cu²⁺ and H⁺ | Copper | Copper is less reactive than hydrogen |
| Ag⁺ and H⁺ | Silver | Silver is less reactive than hydrogen |
| Na⁺ and H⁺ | Hydrogen | Sodium is more reactive than hydrogen |
| Mg²⁺ and H⁺ | Hydrogen | Magnesium is more reactive than hydrogen |
For example, copper(II) ions can be discharged as copper:
Hydrogen formation may be represented at this level as:
Selective discharge at the anode
At the anode, negative ions compete to lose electrons. For common aqueous electrolytes, two ideas are especially important.
- In a concentrated solution containing Cl⁻, Br⁻ or I⁻, the halide ion is usually discharged and the corresponding halogen forms.
- Otherwise, OH⁻ is usually discharged and oxygen is produced.
Sulfate ions, SO₄²⁻, and nitrate ions, NO₃⁻, are not normally discharged under the conditions considered at O-Level.
When OH⁻ is discharged:
When chloride ions are discharged:
All cations are attracted towards the cathode and all anions towards the anode. Selective discharge determines which particle actually reacts at the electrode.
Predicting Products of Aqueous Electrolysis
A reliable way to solve these questions is to list every ion first and then treat the two electrodes separately.
Example 1: copper(II) sulfate with inert electrodes
Aqueous CuSO₄ contains Cu²⁺ and SO₄²⁻ from copper(II) sulfate, together with H⁺ and OH⁻ associated with water.
At the cathode, Cu²⁺ competes with H⁺. Copper is less reactive than hydrogen, so Cu²⁺ is discharged:
A reddish-brown copper deposit forms on the cathode.
At the anode, sulfate ions are not normally discharged. OH⁻ is discharged instead:
Oxygen gas forms at the anode.
As electrolysis continues, Cu²⁺ ions are removed from the solution, so the blue colour of the copper(II) sulfate solution gradually becomes paler.
Example 2: concentrated aqueous sodium chloride
Concentrated sodium chloride solution is called brine. It contains Na⁺, Cl⁻, H⁺ and OH⁻.
At the cathode, sodium is more reactive than hydrogen, so hydrogen is produced instead of sodium metal.
At the anode, the high concentration of chloride ions causes Cl⁻ to be preferentially discharged:
The overall reaction may be written as:
Hydrogen is therefore produced at the cathode, chlorine at the anode, while sodium hydroxide remains in solution.
Recognising common products
| Product | Observation or test |
|---|---|
| Hydrogen | Colourless gas that gives a squeaky pop with a lighted splint |
| Oxygen | Colourless gas that relights a glowing splint |
| Chlorine | Greenish-yellow gas that bleaches damp litmus paper |
| Copper | Reddish-brown solid deposit |
| Silver | Grey or silvery metal deposit |
Why the Electrode Material Matters
So far, we have assumed that the electrodes simply conduct electricity without taking part in the reaction. Graphite and platinum are often used for this purpose because they are relatively unreactive.
These are called inert electrodes.
An electrode can also be made from a substance that participates in the electrode reaction. Such an electrode is sometimes described as an active electrode.
Inert and active electrodes
Inert electrode: provides a conducting surface but does not normally react. Graphite and platinum are common examples.
Active electrode: the electrode material itself takes part in the redox reaction. Copper electrodes in copper(II) sulfate solution are an important example.
Copper electrodes in copper(II) sulfate solution
If copper is used as the anode, copper atoms lose electrons and enter the solution as Cu²⁺ ions:
At the cathode, Cu²⁺ ions gain electrons and are deposited as copper:
Copper is therefore transferred from the anode to the cathode. If copper enters the solution from the anode at approximately the same rate as Cu²⁺ ions are removed at the cathode, the concentration of Cu²⁺ remains approximately constant.
Purification of Copper by Electrolysis
Copper used for electrical wiring must be very pure because impurities increase its electrical resistance. Electrolysis provides a way to transfer copper from an impure block onto a pure copper electrode.
Anode
The positive anode is made of impure copper. Copper atoms lose electrons and enter the electrolyte as Cu²⁺ ions.
Cathode
The negative cathode is a sheet of pure copper. Cu²⁺ ions gain electrons and are deposited here as copper.
Electrolyte
The electrolyte is aqueous copper(II) sulfate, which contains Cu²⁺ ions.
What changes?
The impure copper anode becomes smaller, while the pure copper cathode becomes thicker.
At the anode:
Copper atoms are oxidised.
At the cathode:
Copper ions are reduced.
Insoluble impurities can fall from the impure anode and collect beneath it as anode sludge.
Copper atoms leave the impure anode, become Cu²⁺ ions in the electrolyte, and are eventually deposited as copper atoms on the pure cathode.
Electroplating
Electroplating is the process of coating an object with a thin layer of another metal using electrolysis.
A metal coating may improve appearance, protect a surface from corrosion, improve wear resistance or reduce cost. For example, an object can be coated with silver rather than being made entirely from silver.
How should an electroplating cell be arranged?
Object to be plated
The object must be connected as the negative cathode.
Coating metal
The coating metal is commonly connected as the positive anode.
Electrolyte
The electrolyte must contain ions of the metal that is to be deposited.
Clean surface
The object should be cleaned so that the metal coating can adhere evenly to its surface.
Example: silver-plating a steel spoon
The spoon is connected as the cathode. A silver electrode acts as the anode, and the electrolyte contains Ag⁺ ions.
At the spoon:
Silver ions gain electrons and form a layer of silver metal on the spoon.
At a silver anode:
Silver from the anode replenishes Ag⁺ ions removed from the electrolyte.
Metal plating requires metal ions to gain electrons and become metal atoms. Since reduction occurs at the cathode, the object that receives the coating must be the cathode.
Simple Chemical Cells
Electrolysis uses electricity to force a chemical reaction. A simple chemical cell works in the opposite direction: a spontaneous chemical reaction produces an electric current.
Chemical energy → electrical energy
A simple cell can be made using two different metals in an electrolyte and joining the metals through an external circuit.
The two metals do not have the same tendency to lose electrons. The more reactive metal loses electrons more readily.
Example: magnesium and copper
Magnesium is more reactive than copper. Magnesium atoms therefore lose electrons:
Magnesium is oxidised and forms the negative anode.
The released electrons travel through the external wire from magnesium towards the copper electrode.
The copper electrode is the positive cathode, where reduction takes place. The exact species reduced depends on the electrolyte.
For example, if H⁺ ions are present, they can gain electrons at the copper electrode:
Electrons flow through the external circuit from the more reactive metal towards the less reactive metal.
What happens inside the electrolyte?
Electrons travel through the external wire, not through the electrolyte. Inside the electrolyte, ions move to maintain electrical neutrality as the electrode reactions continue.
Cations move towards the cathode, while anions move towards the anode. This rule is the same as for an electrolytic cell. What changes is the sign of the electrodes.
What determines the voltage?
For simple cells compared under similar conditions, metals that are further apart in the reactivity series generally produce a larger potential difference.
| Metal pair | Expected comparison | Reason |
|---|---|---|
| Magnesium and copper | Larger voltage | The metals are relatively far apart in reactivity |
| Iron and copper | Smaller voltage | The difference in reactivity is smaller |
| Copper and copper | No useful potential difference under identical conditions | Both electrodes have the same chemical tendency |
The exact voltage also depends on the substances and conditions used, but the reactivity-series comparison is a useful qualitative rule for the simple cells studied at O-Level.
Electrolytic Cells and Simple Chemical Cells Compared
Electrolytic cells and simple chemical cells both involve oxidation, reduction and electron transfer. Their main difference is the direction in which energy is converted.
| Feature | Electrolytic cell | Simple chemical cell |
|---|---|---|
| Main function | Uses electricity to cause chemical change | Uses chemical change to produce electricity |
| Energy conversion | Electrical → chemical | Chemical → electrical |
| Anode reaction | Oxidation | Oxidation |
| Cathode reaction | Reduction | Reduction |
| Anode sign | Positive | Negative |
| Cathode sign | Negative | Positive |
| External electrical supply | Required | Not required to drive the reaction |
Anode = oxidation.
Cathode = reduction.
These definitions work for both electrolytic and simple chemical cells.
The Hydrogen–Oxygen Fuel Cell
A fuel cell is another way of producing electricity from a redox reaction. Unlike a battery containing a fixed amount of reacting material, a fuel cell can continue operating while fuel and oxygen are supplied.
In a hydrogen–oxygen fuel cell, hydrogen reacts with oxygen to form water:
During the process, electrons travel through an external circuit, allowing electrical energy to be obtained from the chemical reaction.
Advantages of hydrogen fuel cells
- Water is the main product at the point of use.
- The fuel-cell reaction itself does not produce carbon dioxide.
- Chemical energy is converted directly into electrical energy.
- The fuel cell can continue operating while hydrogen and oxygen are supplied.
Limitations of hydrogen fuel cells
- Hydrogen has a low density and can be difficult to store compactly.
- Hydrogen is highly flammable and must be handled safely.
- Hydrogen storage and distribution require specialised equipment.
- Energy is required to manufacture hydrogen.
- If fossil fuels are used to produce the hydrogen, carbon dioxide may be released elsewhere in the overall process.
Does a hydrogen fuel cell produce no pollution at all?
The reaction inside the fuel cell produces water rather than carbon dioxide. However, the overall environmental impact also depends on how the hydrogen is produced, stored and transported.
It is therefore more accurate to say that a hydrogen fuel cell produces no carbon dioxide at the point of use, rather than claiming that the entire hydrogen system must always be pollution-free.
Common Electrochemistry Mistakes
Mistake 1: Saying that the anode is always positive
The anode is positive during electrolysis but negative in a simple chemical cell. The definition that always works is: oxidation occurs at the anode.
Mistake 2: Saying that the cathode is always negative
The cathode is negative in an electrolytic cell but positive in a simple chemical cell. Reduction always occurs at the cathode.
Mistake 3: Saying that electrons move through the electrolyte
Electrons travel through wires and conducting electrodes. Inside the electrolyte, charge is carried by moving ions.
Mistake 4: Treating molten and aqueous electrolytes in the same way
A molten electrolyte contains only the ions from the compound. In an aqueous solution, H⁺ and OH⁻ from water must also be considered.
Mistake 5: Assuming that every ion reaching an electrode is discharged
Several ions can move towards one electrode, but only the preferentially discharged particle reacts there.
Mistake 6: Depositing a very reactive metal from its aqueous solution
Metals such as sodium and magnesium are more reactive than hydrogen. Hydrogen is normally produced at the cathode instead of these metals.
Mistake 7: Discharging sulfate or nitrate ions at an inert anode
SO₄²⁻ and NO₃⁻ are not normally discharged under O-Level conditions. OH⁻ is usually discharged instead, producing oxygen.
Mistake 8: Forgetting that electrode material can react
An inert graphite electrode behaves differently from an active copper electrode. Always check what the electrodes are made of before predicting the products.
Mistake 9: Connecting the object to be plated as the anode
Metal ions must gain electrons to form the coating. Reduction occurs at the cathode, so the object to be plated must be the cathode.
Mistake 10: Saying that hydrogen fuel cells can never cause environmental impact
The fuel-cell reaction itself produces water, but manufacturing and transporting hydrogen require energy and may produce carbon dioxide depending on how the hydrogen is obtained.
Electrochemistry Checklist
For an electrolysis question, do not try to guess the products immediately. Work through the cell in a fixed order.
1. Identify the state
Is the electrolyte molten or aqueous? Water changes which ions must be considered.
2. List the ions
Write down all cations and anions present before deciding what is discharged.
3. Check the electrodes
Decide whether the electrodes are inert or whether an electrode material may take part in the reaction.
4. Predict the cathode
Reduction occurs here. For aqueous solutions, compare the metal with hydrogen.
5. Predict the anode
Oxidation occurs here. Consider halide concentration and whether OH⁻ will instead form oxygen.
6. Write half-equations
Balance atoms and charge, placing electrons on the correct side of each half-equation.
7. Predict observations
Look for gases, metal deposits, changes in electrode mass and solution colour changes.
8. For simple cells, follow electrons
The more reactive metal normally loses electrons, and electrons flow through the external circuit towards the less reactive metal.
Now apply the rules to unfamiliar electrolytes, electrode materials, electroplating, simple cells and fuel cells.
- 10 application questions
- Explanations for every option
- Test whether you can predict, not just memorise
Frequently Asked Questions
Questions Students Often Ask About Electrochemistry
Most difficulties in electrochemistry come from mixing up ion movement, electron movement and electrode reactions. These questions connect those ideas more carefully.
Why can ions move towards an electrode without being discharged?
Attraction and discharge describe two different things. All positively charged ions move towards the cathode and all negatively charged ions towards the anode. However, when several ions compete, selective discharge determines which particle actually gains or loses electrons.
Why do we consider H⁺ and OH⁻ in an aqueous electrolyte?
The word aqueous means water is present. At this level, water is treated as providing H⁺ and OH⁻ that may compete with ions from the dissolved compound. This is why an aqueous salt solution can produce different products from the molten salt.
Why can molten sodium chloride produce sodium but aqueous sodium chloride cannot?
Molten sodium chloride contains only Na⁺ and Cl⁻, so Na⁺ is the cation reduced at the cathode. In aqueous sodium chloride, hydrogen also competes for discharge. Because sodium is more reactive than hydrogen, hydrogen is produced instead.
Does “inert electrode” mean that no reaction occurs there?
No. Oxidation or reduction still occurs on the electrode surface. “Inert” means that the electrode material itself does not normally take part in the reaction.
Why does copper purification require impure copper at the anode?
Copper must first leave the impure block. At the anode, copper atoms lose electrons and become Cu²⁺ ions. These ions then gain electrons at the cathode and are deposited as pure copper.
Why must the object being electroplated be cleaned first?
Dirt, grease and oxide layers can prevent the coating metal from making good contact with the surface. Cleaning helps produce an even, firmly attached metal coating.
Why does the more reactive metal become negative in a simple chemical cell?
The more reactive metal has a greater tendency to lose electrons. Those electrons enter the external circuit from that electrode, so it is the negative electrode and the anode.
Do ions still move towards the cathode and anode in a simple chemical cell?
Yes. Cations move towards the cathode and anions towards the anode. What changes compared with electrolysis is the sign of the electrodes: in a simple chemical cell, the anode is negative and the cathode is positive.
Why are the signs of the anode and cathode reversed in electrolysis and simple cells?
During electrolysis, the external power supply removes electrons from the anode and supplies them to the cathode. In a simple chemical cell, oxidation at the anode itself releases electrons into the external circuit. The signs therefore differ, but the definitions do not: oxidation occurs at the anode and reduction occurs at the cathode.
Does a bigger difference in reactivity always give an exact predictable voltage?
Not exactly. At O-Level, the reactivity series is useful for comparing which metal pair is likely to produce a larger potential difference under similar conditions. The exact voltage also depends on the substances and experimental conditions.
Is a hydrogen fuel cell the same as a rechargeable battery?
No. A fuel cell continues working while fuel and oxygen are supplied from outside. A rechargeable battery contains its reacting materials internally and is recharged by using electricity to reverse chemical changes.
Why is hydrogen still considered useful if making it requires energy?
Hydrogen is an energy carrier. Its environmental benefit depends partly on how it is produced. If low-carbon electricity is used to manufacture the hydrogen, a fuel-cell system can have a much lower overall carbon impact than one in which the hydrogen is produced using fossil fuels.
Electrochemistry becomes much easier when every question is reduced to electron transfer. At the cathode, reduction occurs because a particle gains electrons. At the anode, oxidation occurs because a particle loses electrons. Electrolysis uses an external electrical supply to force these reactions, while simple chemical cells and fuel cells use redox reactions to drive electrons through a circuit. The apparatus may change, but the underlying Chemistry remains the same.