Electrochemical cells transform chemical energy into electrical energy through redox processes. These cells are essential for many uses, such as industrial operations, batteries, and electrolysis.
Definition and Working Principle
The anode (anode is generally positive side of cell) and cathode (cathode is generally negative side of cell) of an electrochemical cell are both submerged in an electrolyte. At these electrodes, a redox reaction (Oxidation–reduction reactions, commonly known as redox reactions, are reactions that involve the transfer of electrons from one species to another.
The species that loses electrons is said to be oxidized, while the species that gains electrons is said to be reduced.) takes place, producing an electric current. While the reduction reaction happens at the cathode, where electrons are acquired, the oxidation reaction happens at the anode, where electrons are lost. Electrical energy is produced when electrons move from the anode to the cathode via an external circuit.
Types of Electrochemical Cells
Electrochemical cells are broadly classified into two categories:
1.Galvanic (Voltaic) Cells
Voltaic or galvanic cells use spontaneous redox reactions to produce electrical energy. Each of the two half-cells that make up this system has an electrode and an electrolyte solution. The half-cells are connected by a salt bridge, which permits ion passage to preserve electrical neutrality. One kind of galvanic cell that is frequently examined is the Daniell cell, which has the following configuration:
Anode (Oxidation): Zinc electrode (Zn) in ZnSO₄ solution
Cathode (Reduction): Copper electrode (Cu) in CuSO₄ solution
Electrode Reactions:
At Anode: Zn → Zn²⁺ + 2e⁻
At Cathode: Cu²⁺ + 2e⁻ → Cu
An electric current is created when electrons go from the zinc electrode to the copper electrode via an external wire. The Daniell cell finds extensive application in power sources and batteries.
2. Electrolytic Cells
Electrolytic cells generate a non-spontaneous redox reaction with electrical energy. They are employed in metal refining, electroplating, and electrolysis of water. An external power source is necessary for an electrolytic cell to operate, in contrast to a galvanic cell.
The electrolysis of water, which uses electricity to break it down into hydrogen and oxygen gases, is a typical example:
At Anode (Oxidation): 2H₂O → O₂ + 4H⁺ + 4e⁻
At Cathode (Reduction): 4H⁺ + 4e⁻ → 2H₂
Here, the anode is positive, and the cathode is negative, opposite to the setup in a Galvanic cell.
Difference Between Galvanic and Electrolytic Cells
Feature | Galvanic Cell | Electrolytic Cell |
Energy Conversion | Chemical → Electrical | Electrical → Chemical |
Reaction Type | Spontaneous | Non-spontaneous |
Anode Polarity | Negative | Positive |
Cathode Polarity | Positive | Negative |
Example | Daniell Cell | Electrolysis of Water |
Applications of Electrochemical Cells
Electrochemical cells are widely used in various industries and daily life:
1.Batteries: Used in mobile phones, cars, and laptops, including lithium-ion, lead-acid, and nickel-cadmium batteries.
2. Electrolysis: Production of chlorine, sodium hydroxide, and hydrogen gas.
3. Electroplating: Coating of metals like gold, silver, and chromium on objects.
4. Corrosion Prevention: Sacrificial anodes protect metals from rusting.
5. Fuel Cells: Generate electricity using hydrogen and oxygen with high efficiency.
Electrochemical Series:
When various electrodes are arranged in order of increasing or decreasing value of their standard reduction potentials we get a series called electrochemical series.
Reduction half reaction | Standard reduction potential E0 (in volt) |
Li++ e– – Li(s) | -3.05 |
K+ + e– – K(s) | -2.93 |
Ba2+ + 2e– – Ba(s) | -2.90 |
Ca2+ + 2e– – Ca(s) | -2.87 |
Na+ + e– – Na(s) | -2.72 |
Mg2+ + 2e– – Mg(s) | -2.36 |
Al3+ + 3e– – Al(s) | -1.66 |
Mn2+ + 2e– – Mn(s) | -1.18 |
2H2O(l) + 2e– – H2(g) 2OH–(aq) | -0.83 |
Zn2+ + 2e– – Zn(s) | -0.76 |
Cr3+ + 3e– – Cr(s) | -0.74 |
Fe2+ + 2e– – Fe(s) | 0.44 |
Cd2+ + 2e– – Cd(s) | 0.40 |
Co2+ + 2e– – Co(s) | 0.28 |
Ni2+ + 2e– – Ni(s) | 0.25 |
Sn2+ + 2e– – Sn(s) | 0.14 |
Pb2+ + 2e – Pb(s) | 0.13 |
2H+ + 2e– – H2(g) | 0.00 |
Sn4+ + 2e– – Sn2+ | +0.15 |
Cu2+ + 2e– – Cu(s) | +0.34 |
Cu+ + e– – Cu(s) | +0.52 |
I2 + 2e– – 2I– | +0.54 |
Fe3+ + e– – Fe2+ | +0.77 |
Ag+ + e– – Ag(s) | +0.80 |
Br2 + 2e– – 2Br– | +1.09 |
O2 (g)+ 4H+(aq) + 4e– – 2H2O | +1.23 |
Cl2 + 2e– – 2Cl– | +1.36 |
Au3+ + 3e– – Au(s) | +1.40 |
Co3+ + e– – Co2+ | +1.81 |
F2(g) + 2e– – 2F– | +2.87 |
Summary
Modern technology relies heavily on electrochemical cells, which give options for energy conversion and storage. Their functions, distinctions, and uses helps to know the important electrochemistry concepts. These cells transformed science and engineering, from basic batteries to industrial electrolysis.
An electrochemical cell transforms chemical energy into electrical energy through redox processes, (Galvanic cell) or electrical energy into chemical energy (Electrolytic cell).
An electrolytic cell needs external electricity to power a non-spontaneous reaction, whereas a galvanic cell produces electrical energy from spontaneous redox reactions.
A Galvanic cell consists of:
Anode (where oxidation occurs)
Cathode (where reduction occurs)
Electrolyte solutions
Salt bridge (maintains charge balance)
External circuit (for electron flow)
In a Daniell cell:
Zinc (Zn) oxidizes at the anode: Zn → Zn²⁺ + 2e⁻
Copper ions (Cu²⁺) reduce at the cathode: Cu²⁺ + 2e⁻ → Cu
Electrons flow from zinc to copper, generating electricity.
By permitting ion movement between the two half-cells, the salt bridge keeps electrical neutrality intact, avoiding charge accumulation and guaranteeing constant electron flow.
Electrochemical cells are used in:
Batteries (Lithium-ion, lead-acid, dry cells)
Electrolysis processes (metal refining, water splitting)
Electroplating (gold, silver, chrome coating)
Fuel cells (powering vehicles and industries)
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