Electrolysis

By Ali Raza
Roots International School & Colleges
Fundamentals of Electrolysis

Electrolysis
Electrolysis refers to the chemical decomposition of ionic compounds when subjected to an electric
current. This decomposition occurs in a specially designed setup called an electrolytic cell.
Electrolytic Cells
An electrolytic cell is the apparatus used for electrolysis, consisting of:
Electrodes: These are conductors through which electricity enters or leaves the cell, typically made
of inert materials like platinum or graphite.
Electrolyte: A substance containing free ions that make the substance electrically conductive. It can
be either a molten ionic compound or an aqueous solution of salts, acids, or bases.
Electric Current in Electrolysis
The electric current in electrolysis serves as the driving force for chemical reactions. It facilitates the
movement of electrons and ions, thus enabling the decomposition of compounds.
Detailed Roles of Electrolytic Cell Components
Anode: The Site of Oxidation
Function: The anode is the positively charged electrode where oxidation, the loss of electrons,
occurs.
Process: Negative ions (anions) move towards the anode, where they give up their excess electrons
and are discharged as atoms or molecules.

Cathode: The Site of Reduction

Function: The cathode is the negatively charged electrode where reduction, the gain of electrons,
takes place.
Process: Positive ions (cations) are attracted to the cathode, where they acquire electrons and are
discharged.
Electrolyte: The Conductive Medium
Function: The electrolyte contains ions that move towards the respective electrodes, allowing the
flow of electric current through the cell.
Types: Electrolytes can be solid or liquid, such as molten salts or aqueous solutions.
Charge Transfer in Electrolysis
Electron Movement
Direction: Electrons flow from the cathode to the anode through the external circuit.
Role: This movement facilitates the reduction and oxidation processes at the respective electrodes.

Ion Movement in Electrolyte

Understanding Electrolysis Through Examples
Example: Electrolysis of Molten Sodium Chloride

At the Cathode: Sodium ions (Na+) receive electrons and form sodium metal.

At the Anode: Chloride ions (Cl-) release electrons to produce chlorine gas.
Electrolysis of Water
At the Cathode: Hydrogen ions (H+) from water gain electrons to form hydrogen gas.

At the Anode: Hydroxide ions (OH-) lose electrons, resulting in oxygen gas and water.
Factors Influencing Electrolysis
Nature of the Electrolyte
Ionic Composition: Determines the types of ions present and their reactivity.

Conductivity: Influences the efficiency of ion movement and overall reaction rate.

Type of Electrodes

Inert Electrodes: Do not participate in the chemical reactions, e.g., platinum or graphite.

Active Electrodes: Can react with the electrolyte or the products of electrolysis.

Concentration of the Electrolyte

Ion Availability: Higher concentrations increase the number of ions available for reaction.

Reaction Rate: Concentration affects the rate at which ions migrate to the electrodes.
Electrolysis Products and Observations
Molten Lead(II) Bromide
Introduction
Molten lead(II) bromide undergoes electrolysis to produce lead metal and bromine gas. This process
vividly demonstrates the fundamental principles of electrolysis.

Observations
Colour Change: Initially, lead(II) bromide appears as a solid with an orange hue. Upon heating, it
transitions into a clear, molten state.
Electrode Reactions: At the anode, a noticeable emission of brown bromine gas occurs. At the
cathode, the reduction of lead ions results in the formation of silvery droplets of lead.

Products
At the Anode (Positive Electrode): Bromine gas (Br2) is produced, evident from its distinct colour
and odour.
At the Cathode (Negative Electrode): Metallic lead (Pb) is deposited, identifiable by its
characteristic silvery appearance.

Chemical Equations
Anode: 2Br⁻ → Br2 + 2e⁻ (Oxidation)
Cathode: Pb²⁺ + 2e⁻ → Pb (Reduction)
Electrolysis of Solutions
When solutions are electrolyzed, gases are usually produced.

The gases produced can be collected in test tubes to be identified by later simple tests.

The electrolysis of solutions is more complicated than electrolysis of molten compounds, because the
products at the electrode come from the electrolyte as well from water.
Selective Discharge of Ions

At Cathode
• Positive ions from the electrolyte are discharged if they are H+ ions or ions of less
reactive metals such as Cu+2, Pb+2, or Ag+
• Positive ions of reactive metals such as Na+, K+ or Ca+2 are not discharged in the
presence of water. Instead H+ ions from the water are discharged and H 2 gas is
produced.

At Anode
• Negative ions from the electrolyte are discharged if they are halide ions such as I-, Br-,
and Cl-.
• SO42- and NO3- are not discharged. Instead OH- ions from the water are discharged and O2
gas is produced.
Dilute Sulfuric Acid with Inert Electrodes
Introduction
Electrolysing dilute sulfuric acid demonstrates the production of oxygen and hydrogen gases,
embodying the principle of water electrolysis.

Observations
Gas Release: The generation of gas at both electrodes is a key observation.
Solution Change: The solution remains clear, but its conductivity gradually decreases, indicating the
consumption of ions.

Products
At the Anode: Oxygen gas (O2), which is crucial in various industrial processes, is produced.
At the Cathode: Hydrogen gas (H2), a potential clean fuel, is liberated.

Chemical Equations
Anode: 4OH⁻ → O2 + 2H2O + 4e⁻ (Oxidation)
Cathode: 2H⁺ + 2e⁻ → H2 (Reduction)
Concentrated Aqueous Sodium Chloride
Introduction
The electrolysis of concentrated sodium chloride solution, commonly known as brine, results in the
production of chlorine gas, hydrogen gas, and sodium hydroxide.

Observations
Gas Evolution: Evident bubbling at both electrodes indicates gas production.
Colour Change: The solution remains largely clear, with a possible yellow tinge near the anode due
to chlorine gas.

Products
At the Anode: Chlorine gas (Cl2), with its distinctive greenish-yellow colour and pungent odour, is
liberated.
At the Cathode: Hydrogen gas (H2) is formed, alongside hydroxide ions (OH⁻), contributing to the
formation of sodium hydroxide (NaOH) in the solution.

Chemical Equations
Anode: 2Cl⁻ → Cl2 + 2e⁻ (Oxidation)
Cathode: 2H+ + 2e⁻ → H2 (Reduction)
Aqueous Copper(II) Sulfate with Inert and Copper
Electrodes
With Inert Electrodes
Observations
Colour Change: The characteristic blue colour of the solution fades, indicating the depletion of
copper ions.
Electrode Changes: Copper metal deposits on the cathode, while oxygen gas evolves at the anode.
Products
At the Anode: Oxygen gas (O2), a byproduct of water electrolysis in this context.
At the Cathode: Copper metal (Cu), which appears as a reddish deposit.
Chemical Equations
Anode: 4OH⁻ → O2 + 2H2O + 4e⁻ (Oxidation)
Cathode: Cu²⁺ + 2e⁻ → Cu (Reduction)

With Copper Electrodes

Observations
Cathode Change: The cathode experiences a gain in mass as copper is deposited.
Anode Change: The anode gradually loses mass due to the dissolution of copper into the solution.
Products
At the Anode: Copper ions (Cu²⁺) are released into the solution, maintaining the concentration of
the electrolyte.
At the Cathode: Copper metal (Cu) is deposited, enhancing the cathode's mass.
Chemical Equations