Unit 3: Electrochemistry - Overview
This unit explores the interplay between chemical reactions and electrical energy. It explains
how chemical reactions generate electricity and, conversely, how electricity can drive chemical
changes. The unit covers:
Electrical conductivity: Differentiating between metallic and electrolytic conduction.
Electrolysis: Examining the process of driving chemical reactions using electrical energy.
Galvanic (Voltaic) cells: Exploring how chemical reactions produce electricity.
Industrial applications of electrochemistry: Highlighting real-world uses of these
principles.
Electrical Conductivity
The sources distinguish between two types of electrical conductivity:
Metallic Conductivity: This refers to the flow of electricity through metals. The charge
carriers are free electrons, also known as delocalized electrons, which are loosely bound
to metal atoms and can move freely throughout the metallic structure.
Electrolytic Conductivity: This involves the flow of electricity through electrolytes,
which are substances that conduct electricity when dissolved in water or molten. The
charge carriers in this case are ions – positively charged cations and negatively charged
anions.
Electrolysis
Electrolysis is a process that utilizes electrical energy to bring about a non-spontaneous chemical
reaction. This reaction is typically the breakdown of a compound into its constituent elements.
Key components and concepts within electrolysis include:
Electrolytic cell: The apparatus where electrolysis takes place.
Electrolyte: The substance being electrolysed, either molten or dissolved in a solvent.
Electrodes: Electrically conductive materials immersed in the electrolyte, providing a
path for electrons to enter and exit the cell.
Anode: The positive electrode where oxidation occurs.
Cathode: The negative electrode where reduction occurs.
Half-cell reactions: Individual oxidation and reduction reactions occurring at each
electrode.
Cell reaction: The overall balanced chemical equation representing the net change
during electrolysis.
Galvanic (Voltaic) Cells
These cells, unlike electrolytic cells, convert chemical energy into electrical energy through
spontaneous redox reactions. Different types of voltaic cells exist, with the Daniell cell being a
classic example.
The Daniell Cell
The Daniell cell demonstrates the workings of a simple voltaic cell. It comprises:
Two half-cells: One with a zinc electrode in zinc sulphate solution, and the other with a
copper electrode in copper sulphate solution.
Salt bridge: A tube containing an electrolyte solution, connecting the two half-cells. This
bridge allows ion flow to maintain electrical neutrality and complete the circuit.
In the Daniell cell:
Zinc acts as the anode (negative electrode) where oxidation occurs: Zn(s) → Zn2+(aq) +
2e-
Copper acts as the cathode (positive electrode) where reduction occurs: Cu2+(aq) + 2e-
→ Cu(s)
The overall cell reaction is: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Types of Voltaic Cells
The sources distinguish between two main types:
Primary cells: Non-rechargeable and discarded after use, such as the Leclanché cell
(common dry cell).
Secondary cells: Rechargeable by passing an electric current to reverse the cell reactions,
like the lead-acid battery used in vehicles.
Industrial Applications of Electrolysis
Electrolysis plays a crucial role in various industrial processes, including:
Extraction of metals: Used to produce reactive metals like aluminium from their ores.
Purification of metals: Refining metals like copper to achieve high purity.
Electroplating: Coating objects with a thin layer of another metal for protection or
aesthetics.
