Introduction to Electrochemistry

Introduction to Electrochemistry

• Electrochemistry is the branch of chemistry concerned with the interconversion of chemical and electrical energy.
• It involves oxidation-reduction (redox) reactions, where electrons are transferred between species.
• These reactions can be used to generate electricity (voltaic cells) or to drive chemical reactions using electricity (electrolytic cells).

Oxidation-Reduction Reactions

• Oxidation is the loss of electrons, while reduction is the gain of electrons.
• The species that loses electrons is the reducing agent, and the species that gains electrons is the oxidizing agent.
• Redox reactions are balanced using half-reactions.

Electrochemical Cells

• A galvanic (voltaic) cell uses a spontaneous redox reaction to generate electrical energy.
• An electrolytic cell uses electrical energy to force a non-spontaneous redox reaction to occur.
• Key components of an electrochemical cell include: electrodes, electrolytes, and an external circuit.
• The anode is where oxidation occurs, and the cathode is where reduction occurs.

Standard Electrode Potential (E°)

• Standard electrode potential (E°) is a measure of the tendency of a species to gain or lose electrons.
• It is measured in volts (V) under standard conditions (298 K, 1 atm, 1 M solutions).
• The standard hydrogen electrode (SHE) is used as a reference electrode with a standard potential of 0.00 V.

Electrochemical Series

• The electrochemical series lists the standard reduction potentials (E°) for various half-reactions in decreasing order.
• The higher the E° for a half-reaction, the greater the tendency of the species to undergo reduction.
• The electrochemical series can be used to predict the feasibility of redox reactions and the direction of electron flow in a voltaic cell.

Cell Potential (E_cell)

• The cell potential (E_cell) is the difference in standard electrode potentials between the cathode and the anode.
• E_cell = E°_cathode - E°_anode
• A positive E_cell indicates a spontaneous reaction, while a negative E_cell indicates a non-spontaneous reaction.

Nernst Equation

• The Nernst equation relates the cell potential (E) to the cell potential under standard conditions (E°) and the concentrations of the species involved.
• Its important application is in understanding how the cell potential changes with the concentration of reactants and products.
• The equation is critical for calculating non-standard cell potentials.

Types of Electrodes

• Different types of electrodes are used for electrochemical analysis, including:
  • Metal electrodes
  • Ion-selective electrodes
  • Glass electrodes
• These electrodes are chosen based on the specific chemical properties of the analyte.

Applications of Electrochemistry

• Electrochemistry has diverse applications across many fields:
  • Batteries and fuel cells
  • Corrosion
  • Electroplating
  • Analytical chemistry (e.g., voltammetry)
  • Medical devices (e.g., pacemakers)
  • Industrial processes (e.g., refining metals)

Factors Affecting Cell Potential

• Temperature impacts E_cell, as reaction rates and equilibrium positions are temperature-dependent.
• Concentration of reagents and products in the half-cell solutions influences the cell potential, affecting the potential difference between the electrodes.

Electrolysis

• Electrolysis uses electrical energy to drive a non-spontaneous redox reaction.
• The process is commonly used for the production of chemicals, refining metals, and electroplating.
• Faraday's laws of electrolysis describe the relationship between the amount of electricity used and the amount of substance produced in an electrochemical reaction.

Faraday's Laws of Electrolysis

• Faraday's first law states that the mass of a substance produced or consumed during electrolysis is directly proportional to the quantity of electricity passed through the cell.
• Faraday's second law states that the mass of different substances produced or consumed by the same quantity of electricity is directly proportional to their equivalent weights.