🔌Electrochemistry Unit 1 – Introduction to Electrochemistry

Key Concepts and Definitions

• Electrochemistry studies the interrelation of electrical and chemical changes in a system
• Involves the transfer of electrons between chemical species through oxidation-reduction (redox) reactions
• Oxidation occurs when a species loses electrons and its oxidation number increases
• Reduction occurs when a species gains electrons and its oxidation number decreases
• Electrochemical cells convert chemical energy into electrical energy (galvanic cells) or vice versa (electrolytic cells)
• Electrodes are conductors that allow the flow of electrons in an electrochemical cell
  • Anode is the electrode where oxidation occurs
  • Cathode is the electrode where reduction occurs
• Electrolytes are substances that conduct electricity when dissolved in water due to the presence of mobile ions

Electrochemical Cells and Reactions

• Galvanic cells spontaneously convert chemical energy into electrical energy
  • Also known as voltaic cells or batteries
  • Consist of two half-cells connected by a salt bridge or porous membrane
• Electrolytic cells use an external power source to drive a non-spontaneous redox reaction
  • Used in electroplating, electrolysis, and electrochemical synthesis
• Redox reactions involve the transfer of electrons between species
  • Oxidation and reduction always occur simultaneously
  • The species that loses electrons is oxidized (reducing agent)
  • The species that gains electrons is reduced (oxidizing agent)
• Salt bridge maintains electrical neutrality in the cell by allowing the flow of ions between half-cells
• Standard hydrogen electrode (SHE) is used as a reference to measure the potential of other half-cells

Redox Reactions and Half-Cells

• Half-reactions represent the oxidation or reduction process occurring at each electrode
  • Oxidation half-reaction occurs at the anode
  • Reduction half-reaction occurs at the cathode
• Standard reduction potentials ($E^0$) measure the tendency of a species to be reduced under standard conditions
  • More positive $E^0$ indicates a greater tendency to be reduced
  • Tabulated in a standard reduction potential table
• Cell notation describes the components of an electrochemical cell
  • Anode | Anode Electrolyte || Cathode Electrolyte | Cathode
• Balancing redox reactions using the half-reaction method
  • Separate the oxidation and reduction half-reactions
  • Balance atoms and charges in each half-reaction
  • Multiply the half-reactions to equalize the number of electrons transferred
  • Add the half-reactions and cancel out common terms

Electrodes and Electrolytes

• Electrodes are conductors that allow the flow of electrons in an electrochemical cell
  • Can be made of metals, graphite, or other conductive materials
  • Inert electrodes (platinum, gold) do not participate in the redox reaction
  • Active electrodes (zinc, copper) participate in the redox reaction
• Electrolytes are substances that conduct electricity when dissolved in water due to the presence of mobile ions
  • Strong electrolytes completely dissociate into ions (strong acids, strong bases, soluble salts)
  • Weak electrolytes partially dissociate into ions (weak acids, weak bases)
• Concentration cells have the same species in both half-cells but at different concentrations
  • Potential is determined by the concentration gradient
• Ion-selective electrodes respond selectively to a specific ion in solution
  • Used in pH meters and other analytical applications

Nernst Equation and Cell Potential

• Nernst equation relates the cell potential to the standard reduction potentials and concentrations of reactants and products
  • $E_\text{cell} = E_\text{cell}^0 - \frac{RT}{nF} \ln Q$
  • $E_\text{cell}^0$ is the standard cell potential
  • $R$ is the gas constant, $T$ is the temperature, $n$ is the number of electrons transferred, and $F$ is Faraday's constant
  • $Q$ is the reaction quotient, which depends on the concentrations of reactants and products
• Standard cell potential ($E_\text{cell}^0$) is the cell potential under standard conditions (1 M concentrations, 1 atm pressure, 25°C)
  • Calculated as the difference between the standard reduction potentials of the half-reactions
  • $E_\text{cell}^0 = E_\text{cathode}^0 - E_\text{anode}^0$
• Gibbs free energy change ($\Delta G$) is related to the cell potential
  • $\Delta G = -nFE_\text{cell}$
  • Negative $\Delta G$ indicates a spontaneous reaction
• Concentration cells have a non-zero potential due to the concentration gradient
  • Potential is given by a simplified Nernst equation: $E_\text{cell} = \frac{RT}{nF} \ln \frac{[\text{ox}]}{[\text{red}]}$

Applications in Daily Life

• Batteries are galvanic cells that convert chemical energy into electrical energy
  • Primary batteries are single-use and cannot be recharged (alkaline batteries)
  • Secondary batteries are rechargeable (lithium-ion batteries)
• Fuel cells generate electricity from the oxidation of a fuel (hydrogen, methanol)
  • Used in space exploration, transportation, and stationary power generation
• Corrosion is an electrochemical process that involves the oxidation of a metal
  • Rusting of iron is a common example of corrosion
  • Cathodic protection prevents corrosion by making the metal the cathode in an electrochemical cell
• Electroplating is an electrolytic process that deposits a thin layer of metal onto a surface
  • Used in jewelry making, automotive industry, and electronics manufacturing
• Electrolysis is used to produce pure substances or drive non-spontaneous reactions
  • Electrolysis of water produces hydrogen and oxygen gases
  • Hall-Héroult process produces aluminum via electrolysis of molten alumina

Lab Techniques and Safety

• Potentiometry measures the potential difference between two electrodes
  • Used in pH meters, ion-selective electrodes, and redox titrations
• Coulometry measures the amount of charge passed during an electrolytic process
  • Used to determine the concentration of an analyte or the stoichiometry of a reaction
• Cyclic voltammetry studies the redox behavior of a system by varying the potential and measuring the current
  • Provides information about the reversibility and kinetics of a redox reaction
• Proper handling and disposal of chemicals are essential for lab safety
  • Wear personal protective equipment (gloves, goggles, lab coat)
  • Work in a well-ventilated area and use a fume hood when necessary
  • Dispose of waste according to regulations and guidelines
• Electrical safety is crucial when working with electrochemical equipment
  • Avoid contact with exposed wires or connections
  • Use grounded outlets and ground fault circuit interrupters (GFCIs)
  • Follow the manufacturer's instructions and safety guidelines

Challenges and Future Directions

• Developing high-performance, sustainable, and cost-effective energy storage systems
  • Improving the energy density, cycle life, and safety of batteries
  • Exploring new materials and chemistries for advanced batteries (solid-state, lithium-sulfur, lithium-air)
• Enhancing the efficiency and durability of fuel cells
  • Developing low-cost, high-performance catalysts
  • Optimizing the design and manufacturing processes
• Addressing the environmental impact of electrochemical technologies
  • Recycling and proper disposal of batteries and other electrochemical devices
  • Minimizing the use of toxic or scarce materials
• Advancing the understanding of fundamental electrochemical processes
  • Investigating the structure and dynamics of the electrode-electrolyte interface
  • Modeling and simulating complex electrochemical systems
• Expanding the applications of electrochemistry in various fields
  • Biomedical devices (sensors, drug delivery systems)
  • Environmental monitoring and remediation
  • Chemical synthesis and catalysis