Electrolytic Cell – Definition, How It Works, and Examples

An electrolytic cell is an electrochemical device that uses electrical energy to drive a nonspontaneous chemical reaction. Electrolytic cells are essential to modern chemistry, metallurgy, manufacturing, and energy technology because they enable reactions that would not occur on their own. These cells perform processes such as electroplating, metal purification, water electrolysis, chlorine production, and rechargeable battery charging. Unlike galvanic cells, which convert chemical energy into electrical energy, electrolytic cells consume electrical energy to force chemical change.

Electrolytic cells operate according to the principles of oxidation-reduction (redox) chemistry. Oxidation occurs at the anode, while reduction occurs at the cathode. However, because an external power source drives the reaction, the electrode charges are opposite those in a galvanic cell. The anode is positive and the cathode is negative in an electrolytic cell.

The study of electrolysis helped establish the foundations of electrochemistry during the nineteenth century. Scientists such as Luigi Galvani, Alessandro Volta, Michael Faraday, and Humphry Davy contributed to the understanding of how electricity and chemical reactions relate to one another. Today, electrolytic cells are central to industrial chemistry and emerging clean-energy technologies such as hydrogen fuel production.

Key Takeaways: Electrolytic Cell

• An electrolytic cell uses electrical energy to drive a nonspontaneous redox reaction.
• Oxidation occurs at the anode and reduction occurs at the cathode.
• In an electrolytic cell, the anode is positive and the cathode is negative.
• Electrons flow from the power source to the cathode and from the anode back to the power source.
• Conventional current flows opposite electron flow.
• Electrolytic cells perform electroplating, electrorefining, water splitting, and metal extraction.
• Faraday’s laws relate the amount of substance produced to the electrical charge passed through the cell.
• Electrolytic cells differ from galvanic cells because they require an external power source.

Purpose of an Electrolytic Cell

An electrolytic cell forces chemical reactions that are thermodynamically unfavorable under normal conditions. In other words, the cell uses electricity to make a nonspontaneous reaction occur.

The main purposes of electrolytic cells include:

• Producing chemical substances
• Purifying metals
• Coating surfaces with metals
• Decomposing compounds
• Storing energy in rechargeable batteries
• Generating gases such as hydrogen and oxygen

For example, aluminum extraction from bauxite requires electrolysis because aluminum ions do not spontaneously reduce to metallic aluminum. Similarly, electroplating deposits thin layers of metals such as chromium, silver, or gold onto objects.

History of the Electrolytic Cell

The development of the electrolytic cell closely followed the discovery of electricity and electrochemistry.

Early Discoveries

In the late eighteenth century, Luigi Galvani observed that frog muscles contracted in the presence of electricity. Alessandro Volta later developed the voltaic pile, the first practical battery, in 1800. This invention provided a continuous source of electric current for chemical experiments.

Humphry Davy and Electrolysis

Soon after Volta’s discovery, English chemist Humphry Davy used electrolysis to isolate highly reactive metals, including sodium, potassium, calcium, and magnesium. His experiments demonstrated that electricity could decompose compounds into elements.

Faraday’s Contributions

Michael Faraday established the quantitative laws of electrolysis during the 1830s. He showed that the amount of material produced during electrolysis depends on the electric charge passed through the system. Faraday also introduced important electrochemical terminology, including electrode, anode, cathode, electrolyte, and ion.

Industrial Applications

During the nineteenth and twentieth centuries, electrolysis became essential for industrial chemistry. Important developments included:

• Electroplating
• Aluminum production by the Hall-Héroult process
• Chlor-alkali production of chlorine and sodium hydroxide
• Copper electrorefining

Modern electrolytic technology now supports renewable energy storage and hydrogen fuel production.

What Is an Electrolytic Cell?

An electrolytic cell is a type of electrochemical cell that converts electrical energy into chemical energy by driving a nonspontaneous redox reaction.

The cell contains:

• Two electrodes
• An electrolyte containing ions
• An external power source

The power source pushes electrons through the external circuit, forcing oxidation and reduction reactions to occur at the electrodes.

Electrolytic Cell vs Electrochemical Cell

An electrochemical cell is any device that converts chemical energy and electrical energy between one another through redox reactions.

There are two main categories:

1. Galvanic (voltaic) cells
2. Electrolytic cells

A galvanic cell generates electricity from a spontaneous reaction, while an electrolytic cell consumes electricity to drive a nonspontaneous reaction.

Thus, every electrolytic cell is an electrochemical cell, but not every electrochemical cell is electrolytic.

Electrolytic Cell vs Galvanic Cell

An electrolytic cell and a galvanic cell both involve redox reactions and electron transfer, but they operate in opposite ways.

Electrolytic Cell vs Electrolytic Capacitor

Students sometimes confuse these terms because of the similar names.

An electrolytic capacitor stores electrical charge using an electrolyte and oxide layer, but it does not perform electrolysis during normal operation.

An electrolytic cell drives chemical reactions through electrolysis.

How an Electrolytic Cell Works

An electrolytic cell uses an external voltage source to force electrons to move in a direction opposite that favored by the spontaneous reaction.

The power source pulls electrons away from the anode and pushes electrons toward the cathode.

Oxidation at the Anode

Oxidation always occurs at the anode:$$\text{Anode: } \text{species loses electrons}$$

Example:$$2Cl^- \rightarrow Cl_2(g) + 2e^-$$

Reduction at the Cathode

Reduction always occurs at the cathode:$$\text{Cathode: } \text{species gains electrons}$$

Example:$$Cu^{2+} + 2e^- \rightarrow Cu(s)$$

Electron Flow

Electrons move:

• From the power supply to the cathode
• From the anode back to the power supply

Current Flow

Conventional current flows opposite electron flow:

• From the positive terminal toward the cathode through the external circuit

Components and Construction of an Electrolytic Cell

Electrodes

The electrodes conduct electricity and provide surfaces for redox reactions.

• Anode: site of oxidation
• Cathode: site of reduction

Electrodes may be:

• Inert (graphite, platinum)
• Reactive (copper, nickel, silver)

Electrolyte

The electrolyte contains mobile ions that carry charge through the solution or molten material.

Examples include:

• Molten NaCl
• CuSO₄(aq)
• H₂SO₄(aq)

External Power Supply

A battery or DC power source supplies electrical energy needed to force the reaction.

Container

The container holds the electrolyte and electrodes. Some cells use porous barriers or membranes to separate products.

Electrolytic Cell Diagram and Labeling Conventions

Students often confuse electrode signs and electron flow directions.

Key points to remember:

• Oxidation always occurs at the anode.
• Reduction always occurs at the cathode.
• In electrolytic cells:
  • Anode = positive
  • Cathode = negative
• Electrons move through wires.
• Ions move through the electrolyte.

A simple mnemonic also helps:

• AN OX
  • Anode = Oxidation
• RED CAT
  • Reduction = Cathode

Electrolysis of Molten Compounds

Molten ionic compounds contain only ions from the compound itself. Because water is absent, the ions directly undergo oxidation and reduction.

Example: molten NaCl

Cathode:$$Na^+ + e^- \rightarrow Na(l)$$

Anode:$$2Cl^- \rightarrow Cl_2(g) + 2e^-$$

Products:

• Sodium metal
• Chlorine gas

Electrolysis of Aqueous Solutions

Aqueous solutions contain dissolved ions and water molecules, so water may react instead of the dissolved ions.

For example, aqueous NaCl produces hydrogen gas rather than sodium metal because water reduces more easily than sodium ions:

Cathode:$$2H_2O + 2e^- \rightarrow H_2(g) + 2OH^-(aq)$$

Anode:$$2Cl^- \rightarrow Cl_2(g) + 2e^-$$

Products:

• Hydrogen gas
• Chlorine gas
• Sodium hydroxide in solution

General Rules for Aqueous Electrolysis

At the cathode:

• Less reactive metal ions (Cu²⁺, Ag⁺) usually reduce to metals.
• Highly reactive metal ions (Na⁺, K⁺, Ca²⁺) usually do not reduce; water produces hydrogen gas instead.

At the anode:

• Halide ions often oxidize to halogen gases.
• Otherwise, water commonly oxidizes to oxygen gas.

Predicting Electrode Reactions

Predicting reactions in electrolysis requires identifying which species undergo oxidation and reduction.

Step 1: Identify All Species Present

List ions, solvent molecules, and electrode materials.

For aqueous NaCl:

• Na⁺
• Cl⁻
• H₂O

Step 2: Determine Possible Reductions

Reduction occurs at the cathode.

Possible reductions:$$Na^+ + e^- \rightarrow Na$$$$2H_2O + 2e^- \rightarrow H_2 + 2OH^-$$

Water reduces more easily than sodium ions in aqueous solution.

Step 3: Determine Possible Oxidations

Oxidation occurs at the anode.

Possible oxidations:$$2Cl^- \rightarrow Cl_2 + 2e^-$$$$2H_2O \rightarrow O_2 + 4H^+ + 4e^-$$

Under many conditions, chloride oxidizes preferentially.

Step 4: Combine Half-Reactions

Balance electrons and add the half-reactions.

Worked Example Problems

Example 1: Electrolysis of Molten Sodium Chloride

Predict the products of electrolysis for molten NaCl.

Step 1: Identify Species

Molten NaCl contains:

• Na⁺
• Cl⁻

Step 2: Cathode Reaction

Reduction:$$Na^+ + e^- \rightarrow Na(l)$$

Step 3: Anode Reaction

Oxidation:$$2Cl^- \rightarrow Cl_2(g) + 2e^-$$

Step 4: Balance and Combine

Multiply the sodium reaction by 2:$$2Na^+ + 2e^- \rightarrow 2Na$$

Overall reaction:$$2NaCl(l) \rightarrow 2Na(l) + Cl_2(g)$$

Example 2: Electroplating Copper

A silver spoon is electroplated using CuSO₄ solution and copper electrodes.

Cathode

Copper ions reduce onto the spoon:$$Cu^{2+} + 2e^- \rightarrow Cu(s)$$

Anode

Copper metal dissolves:$$Cu(s) \rightarrow Cu^{2+} + 2e^-$$

Result

Copper transfers from the anode to the spoon.

Faraday’s Law of Electrolysis

Faraday’s law relates the amount of substance produced to the quantity of electric charge passed through the cell.

The basic equation is:

$m=\frac{ItM}{nF}$​

Where:

• $m$ = mass produced
• $I$ = current
• $t$ = time
• $M$ = molar mass
• $n$ = electrons transferred
• $F$ = Faraday constant = $96485 \, C/mol$

Worked Example Using Faraday’s Law

How much copper deposits when a 2.50 A current runs through CuSO₄ solution for 1.50 hours?

Step 1: Write the Reduction Reaction

$$Cu^{2+} + 2e^- \rightarrow Cu$$

Thus:

• $n = 2$
• $M = 63.55 \, g/mol$

Step 2: Convert Time

$$1.50 \, h \times 3600 \, s/h = 5400 \, s$$

Step 3: Calculate Charge

$$Q = It$$$$Q = (2.50)(5400) = 13500 \, C$$

Step 4: Apply Faraday’s Law

$$m = \frac{(13500)(63.55)}{(2)(96485)}$$$$m \approx 4.45 \, g$$

Therefore, 4.45 g of copper deposits.

Standard Reduction Potentials and Electrolysis

Electrode potentials help predict whether a reaction occurs spontaneously or requires external energy.

A spontaneous galvanic reaction has:$$E^\circ_{cell} > 0$$

A nonspontaneous electrolytic reaction has:$$E^\circ_{cell} < 0$$

The external power supply in an electrolytic cell provides enough voltage to overcome the negative cell potential.

For example, decomposition of water requires external voltage because the reaction is nonspontaneous under standard conditions.

Overpotential and Real-World Efficiency

Real electrolytic cells often require more voltage than theoretical calculations predict.

This extra required voltage is called overpotential or overvoltage.

Causes include:

• Resistance in the electrolyte
• Slow electrode kinetics
• Gas bubble formation
• Electrode surface effects

For example, water electrolysis theoretically requires about 1.23 V under standard conditions, but practical systems often require 1.8–2.2 V.

Overpotential affects:

• Industrial energy consumption
• Hydrogen production efficiency
• Battery charging
• Corrosion prevention systems

Applications of Electrolytic Cells

Electrolytic cells have numerous industrial and technological uses.

Electroplating

Electroplating deposits thin metal coatings for:

• Corrosion resistance
• Decoration
• Improved conductivity

Examples include chrome plating and gold plating.

Electrorefining

Electrorefining purifies metals such as copper.

Impure copper serves as the anode, while pure copper deposits at the cathode.

Metal Extraction

Electrolysis extracts reactive metals including:

• Aluminum
• Sodium
• Magnesium

Water Electrolysis

Electrolysis splits water into hydrogen and oxygen gases:
$$2H_2O(l) \rightarrow 2H_2(g) + O_2(g)$$

Hydrogen production is important for clean-energy technologies.

Chlor-Alkali Process

Electrolysis of brine produces:

• Chlorine gas
• Hydrogen gas
• Sodium hydroxide

These chemicals are important industrial feedstocks.

Rechargeable Batteries

Charging rechargeable batteries involves electrolytic processes that reverse spontaneous discharge reactions.

Hall-Héroult Process

The Hall-Héroult process extracts aluminum from molten aluminum oxide dissolved in cryolite.

Overall reaction:$$2Al_2O_3(l) \rightarrow 4Al(l) + 3O_2(g)$$

Electrolytic Cells and Renewable Energy

Electrolysis is increasingly important in clean-energy technology.

Green Hydrogen

Renewable electricity can electrolyze water to produce hydrogen fuel without fossil fuels.

Energy Storage

Electrolysis allows conversion of excess renewable electricity into chemical energy that is easier to store and transport.

Fuel Cell Connection

Fuel cells reverse electrolysis by converting hydrogen and oxygen back into electricity and water.

Safety Considerations in Electrolysis

Potential hazards include:

• Explosive hydrogen gas
• Toxic chlorine gas
• Hot molten salts
• Electrical shock
• Corrosive electrolytes

Safe laboratory practices include:

• Good ventilation
• Eye protection
• Low-voltage power supplies
• Avoiding confined gas buildup

Common Misconceptions

The Cathode Is Always Positive

This statement is false.

• In galvanic cells, the cathode is positive.
• In electrolytic cells, the cathode is negative.

The cathode is defined by reduction, not charge.

Electrons Flow Through the Solution

Electrons move through the external circuit, not through the electrolyte solution.

Ions carry charge through the electrolyte.

Oxidation Means Oxygen Is Present

Oxidation simply means loss of electrons. Oxygen does not need to participate.

Current and Electron Flow Are the Same

They move in opposite directions.

• Electrons flow from negative to positive.
• Conventional current flows from positive to negative.

Electrolysis Always Produces Pure Metals

Side reactions and impurities often occur, especially in aqueous solutions.

FAQs

What Is the Difference Between Electrolysis and an Electrolytic Cell?

Electrolysis is the process of using electricity to drive a chemical reaction. An electrolytic cell is the device that performs electrolysis.

Why Is the Anode Positive in an Electrolytic Cell?

The external power supply pulls electrons away from the anode, leaving it positively charged.

Why Is the Cathode Negative?

The power source pushes electrons toward the cathode, giving it a negative charge.

Which Direction Do Electrons Flow?

Electrons flow:

• From the negative terminal of the power supply to the cathode
• From the anode back to the positive terminal

Which Direction Does Current Flow?

Conventional current flows opposite electron flow.

Can Water Be Electrolyzed?

Yes. Water electrolysis produces hydrogen and oxygen gas.

Why Are Inert Electrodes Sometimes Necessary?

Inert electrodes prevent the electrode itself from reacting. Platinum and graphite are common choices.

Why Does Electrolysis of Aqueous Solutions Often Produce Hydrogen Gas?

Water is sometimes easier to reduce than dissolved metal ions, especially for highly reactive metals.

Is Electrolysis Endothermic?

Electrolysis requires energy input, but the exact thermodynamics depend on the specific reaction.

Are Electrolytic Cells Used in Everyday Life?

Yes. Applications include:

• Rechargeable batteries
• Metal plating
• Aluminum production
• Water purification
• Hydrogen generation

References and Further Reading

• Harris, Daniel C. (2010). Quantitative Chemical Analysis (8th ed.). New York: W. H. Freeman and Company. ISBN 978-1-4292-1815-3.
• Mortimer, Robert G. (2008). Physical Chemistry (3rd ed.). Amsterdam: Academic Press/Elsevier. ISBN 978-0-12-370617-1.
• Murphy, B.; Murphy, C.; Hathaway, B., eds. (1997). “Electrochemistry I: Galvanic Cells”. A Working Method Approach for Introductory Physical Chemistry Calculations. Cambridge, United Kingdom: Royal Society of Chemistry. ISBN 0-85404-553-8.
• Skoog, Douglas A.; West, Donald M.; Holler, F. James; Crouch, Stanley R. (2014). Fundamentals of Analytical Chemistry. Belmont, CA: Brooks/Cole, Cengage Learning. ISBN 978-0-495-55828-6.
• Wendt, Hartmut; Kolb, Dieter M.; Engelmann, Gerald E.; Ziegler, Jörg C. (2011). “Electrochemistry, 1. Fundamentals”, in Wiley-VCH Verlag GmbH & Co. KGaA (ed.). Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim, Germany: Wiley-VCH Verlag GmbH & Co. KGaA. doi:10.1002/14356007.a09_183.pub4. ISBN 978-3-527-30673-2.