Electrochemistry Detailed Notes
1. Introduction:
- Electrochemistry is the branch of chemistry that deals with the relationship between electrical energy and chemical changes.
- It involves the study of redox (reduction-oxidation) reactions where transfer of electrons occurs.
- Applications include batteries, fuel cells, electroplating, and corrosion.
2. Electrochemical Cells:
- Galvanic (Voltaic) Cells: These cells convert chemical energy into electrical energy through spontaneous redox reactions.
- Example: Daniel cell.
- Consists of two half-cells connected by a salt bridge. Each half-cell contains an electrode and an electrolyte.
- Oxidation occurs at the anode, and reduction occurs at the cathode.
- Electrolytic Cells: These cells use electrical energy to drive non-spontaneous chemical reactions.
- Used in processes like electroplating, electrolysis of water, and extraction of metals.
- An external power source is required to provide the necessary energy.
3. Standard Electrode Potential (E°):
- The standard electrode potential measures the tendency of a half-cell to gain or lose electrons under standard conditions (298 K, 1 atm pressure, 1 M concentration).
- Standard Hydrogen Electrode (SHE) is used as the reference electrode with a potential of 0V.
- The potential of other electrodes is measured relative to SHE.
4. Nernst Equation:
- The Nernst equation relates the cell potential to the standard cell potential and the reaction quotient (Q).
- E = E° - (RT/nF) ln Q
- At 25°C (298 K): E = E° - (0.0591/n) log Q
- R is the gas constant (8.314 J/mol·K), T is the temperature in Kelvin, n is the number of moles of electrons transferred, F is the Faraday constant (96485 C/mol).
