Galvanic cells, electrolytic cells, cell potential

1. Introduction to Electrochemical Cells

1. Electrochemical cells convert chemical energy into electrical energy or vice versa.
2. They are classified into two types: Galvanic cells (voltaic cells) and electrolytic cells.
3. In both types, redox reactions occur, involving electron transfer.

2. Galvanic Cells

1. A Galvanic cell generates electrical energy from spontaneous redox reactions.
2. Composed of two half-cells, each containing an electrode and an electrolyte.
3. The anode is the site of oxidation, and the cathode is the site of reduction.
4. Electrons flow from the anode to the cathode through an external circuit.
5. A salt bridge or porous membrane maintains electrical neutrality by allowing ion flow.
6. Example: The Daniel cell, where Zn is oxidized to Zn²⁺ and Cu²⁺ is reduced to Cu.

3. Electrolytic Cells

1. An electrolytic cell uses electrical energy to drive a non-spontaneous redox reaction.
2. The anode is positively charged (oxidation occurs), and the cathode is negatively charged (reduction occurs).
3. Example: Electrolysis of water to produce hydrogen and oxygen gases.
4. Used in industrial processes like electroplating, metal extraction, and electrorefining.

4. Cell Potential

1. The cell potential (E_cell) is the difference in potential between the two electrodes.
2. Measured in volts and determined by the nature of the electrodes and electrolytes.
3. A positive E_cell indicates a spontaneous reaction (Galvanic cell).
4. A negative E_cell requires external energy (Electrolytic cell).
5. The standard electrode potential (E^°) is measured under standard conditions (1 M concentration, 1 atm pressure, 25°C).
6. The Nernst equation relates E_cell to concentration and temperature: 
   E = E^° - (RT/nF)lnQ
7. Where:
   • E = cell potential under non-standard conditions.
   • E^° = standard cell potential.
   • R = universal gas constant (8.314 J/mol·K).
   • T = temperature in Kelvin.
   • n = number of electrons transferred.
   • F = Faraday’s constant (96,485 C/mol).
   • Q = reaction quotient.

5. Applications of Electrochemical Cells

1. Galvanic cells are used in batteries, such as dry cells, lithium-ion batteries, and lead-acid batteries.
2. Electrolytic cells are used in industrial processes like chlor-alkali process and aluminum extraction.
3. Fuel cells generate energy with high efficiency and minimal pollution.

6. Key Points

1. Galvanic cells convert chemical energy to electrical energy, while electrolytic cells do the opposite.
2. The anode is the site of oxidation, and the cathode is the site of reduction in both cell types.
3. The salt bridge maintains charge balance in a Galvanic cell.
4. Standard electrode potentials (E^°) determine cell potential.
5. The Nernst equation adjusts cell potential for non-standard conditions.
6. Applications include batteries, electroplating, and metal refining.