Electrochemistry

Electrochemistry is the branch of physical chemistry that studies the relationship between electrical energy and chemical changes. It primarily focuses on two phenomena: the generation of electricity from energy released during spontaneous chemical reactions, and the use of electrical energy to bring about non-spontaneous chemical transformations.

Core Concepts and Fundamental Terminologies

To understand electrochemical systems, several foundational concepts must be defined:

• Oxidation and Reduction (Redox): Oxidation involves the loss of electrons, while reduction involves the gain of electrons. These reactions always occur simultaneously.
• Electrolyte: A substance that dissociates into ions when dissolved in water or fused, allowing it to conduct electricity. Examples include aqueous NaCl, H₂SO₄, and CuSO₄.
• Electrodes: Metallic or electronic conductors through which an electric current enters or leaves an electrolyte.
  • Anode: The electrode where oxidation takes place. In a galvanic cell, it is the negative terminal; in an electrolytic cell, it is the positive terminal.
  • Cathode: The electrode where reduction takes place. In a galvanic cell, it is the positive terminal; in an electrolytic cell, it is the negative terminal.

Classification of Electrochemical Cells

Electrochemical cells are broadly categorized into two primary types based on their operational mechanism.

Feature	Electrochemical / Galvanic / Voltaic Cell	Electrolytic Cell
Energy Conversion	Converts Chemical Energy into Electrical Energy.	Converts Electrical Energy into Chemical Energy.
Spontaneity	Based on spontaneous redox reactions (Δ G < 0).	Based on non-spontaneous redox reactions (Δ G > 0).
Anode Sign	Negative (-)	Positive (+)
Cathode Sign	Positive (+)	Negative (-)
Salt Bridge	Required to maintain electrical neutrality.	Not required; both electrodes are usually in the same container.
Key Example	Daniell Cell, Lithium-ion battery (during discharge).	Electroplating units, Aluminium extraction (Hall-Héroult process).

Electrochemical Series and Electrode Potential

When a metal electrode is dipped into its own ionic solution, it develops a potential difference called the electrode potential. When measured under standard conditions (298 K, 1 atm pressure, 1 M concentration) relative to a Standard Hydrogen Electrode (SHE), it is called the Standard Electrode Potential (E°).

The Electrochemical Series

The electrochemical series is an arrangement of various electrodes in the decreasing or increasing order of their standard reduction potentials.

• Strongest Reducing Agents: Elements with highly negative reduction potentials (e.g., Lithium, Potassium) sit at the top/bottom depending on the convention, meaning they lose electrons very easily.
• Strongest Oxidizing Agents: Elements with highly positive reduction potentials (e.g., Fluorine) accept electrons very easily.

Applications of the Series

• Predicting Feasibility: A redox reaction is spontaneous only if the total electromotive force (EMF) of the cell (E°_cell = E°_cathode – E°_anode) is positive.
• Displacement Reactions: A metal with a lower reduction potential (more reactive) will displace a metal with a higher reduction potential from its salt solution. For example, Zinc displaces Copper from CuSO₄ solution.

Laws Governing Electrolysis

Electrolysis is the process of decomposition of an electrolyte by passing continuous electric current through it. This process is quantitatively governed by Michael Faraday’s two laws.

Faraday’s First Law of Electrolysis

The mass (m) of any substance deposited or liberated at any electrode during electrolysis is directly proportional to the quantity of electricity (Q) passed through the electrolyte.

Where I is current in amperes, t is time in seconds, and z is the electrochemical equivalent of the substance.

Faraday’s Second Law of Electrolysis

When the same quantity of electricity is passed through several electrolytic solutions connected in series, the masses of the substances liberated at the electrodes are directly proportional to their chemical equivalent weights (E).

Commercial Cells and Batteries

Batteries act as practical applications of galvanic cells. They are classified into three major groups based on usability.

Primary Cells

These cells cannot be recharged because the chemical reaction occurs only once, and the cell becomes dead after a period of use.

• Dry Cell (Leclanché Cell): Uses a zinc container as the anode, a carbon (graphite) rod surrounded by MnO₂ as the cathode, and a moist paste of NH₄Cl and ZnCl₂ as the electrolyte. Commonly used in transistors and clocks.
• Mercury Cell: Consists of zinc-mercury amalgam as anode and a paste of HgO and carbon as cathode. It provides a constant voltage of 1.35 V throughout its life, making it suitable for low-current devices like hearing aids and watches.

Secondary Cells

These cells can be recharged by passing an electric current through them in the opposite direction, reversing the chemical reactions.

• Lead-Acid Storage Battery: Used commonly in automobiles and invertors. It consists of a lead anode and a grid of lead packed with lead dioxide (PbO₂) as the cathode. The electrolyte is a 38% solution of sulphuric acid (H₂SO₄).
• Nickel-Cadmium (Ni-Cd) Cell: Has a longer life than lead-acid cells but is more expensive. It utilizes a cadmium anode and a nickel dioxide cathode.
• Lithium-Ion Batteries: The mainstay of modern portable electronics and Electric Vehicles (EVs). They utilize lithium intercalation compounds as electrodes rather than metallic lithium, ensuring higher energy density, lower weight, and no memory effect.

Fuel Cells

Galvanic cells designed to convert the energy of combustion of fuels like hydrogen, methane, or methanol directly into electrical energy.

• Hydrogen-Oxygen Fuel Cell: Used extensively in space programmes (such as the Apollo space missions). The water vapour produced during its operation is condensed and added to the drinking water supply for astronauts. It operates with an efficiency of nearly 70%, significantly higher than thermal power plants.

Corrosion as an Electrochemical Phenomenon

Corrosion is the slow conversion of metal surfaces into undesirable compounds (like oxides, sulphides, or carbonates) due to interaction with atmospheric gases and moisture.

Mechanism of Rusting of Iron

Rusting is an electrochemical process that occurs on uneven iron surfaces.

• At Anode Spot: Iron oxidizes to ferrous ions: Fe → Fe²⁺ + 2e^-
• At Cathode Spot: Oxygen in the presence of H^+ ions (derived from dissolved CO₂ or SO₂ in water) is reduced: O₂ + 4H^+ + 4e^- → 2H₂O
• The ferrous ions (Fe²⁺) are further oxidized by atmospheric oxygen to ferric ions (Fe³⁺), forming hydrated ferric oxide (Fe₂O₃ · xH₂O), which is rust.

Methods of Prevention

• Barrier Protection: Painting, oiling, or greasing the metal surface to block contact with air and moisture.
• Galvanization: Coating iron with a thin layer of zinc. Zinc has a lower reduction potential than iron, meaning it oxidizes preferentially, protecting the underlying iron even if the zinc layer is scratched or broken (Sacrificial Protection).
• Cathodic Protection: Connecting the iron structure to a more reactive metal block (like Magnesium or Zinc) via a wire. The more reactive metal acts as the sacrificial anode, protecting the iron which acts as the cathode.

Key Historical Facts and Scientific Trivia

• Luigi Galvani & Alessandro Volta: In the late 18th century, Galvani observed contraction of frog leg muscles when touched with different metals, calling it “animal electricity.” Volta disproved the biological necessity, demonstrating that electricity could be generated solely by inorganic materials, leading to the invention of the Voltaic Pile (the first electrical battery) in 1800.
• The Hall-Héroult Process: Aluminium extraction was prohibitively expensive until 1886 when Charles Martin Hall and Paul Héroult independently discovered that electrolyzing a molten mixture of alumina (Al₂O₃) and cryolite (Na₃AlF₆) drastically lowered the melting point and allowed viable industrial production.
• The Demarcation of Corrosion: Gold and Platinum do not undergo corrosion under standard atmospheric conditions because they have highly positive standard reduction potentials, making them chemically inert to atmospheric oxygen and moisture.

Last Modified: May 26, 2026