CHEM122 Electrochemistry

Introduction to Electrochemistry

Overview of Electrochemistry

• Electrochemistry is the study of the interchange of chemical and electrical energy, primarily through oxidation-reduction (redox) reactions.

• These reactions can generate electric current from chemical reactions or use electric current to induce chemical changes.

• Key applications include batteries, fuel cells, and electrolysis.

Key Concepts in Electrochemistry

• Cell Potential: The voltage produced by a galvanic cell, indicating the driving force behind the electrochemical reaction.

• Gibbs Free Energy (ΔG): A thermodynamic potential that measures the maximum reversible work obtainable from a thermodynamic system at constant temperature and pressure.

Importance of Redox Reactions

• Redox reactions involve the transfer of electrons; oxidation is the loss of electrons, while reduction is the gain of electrons.

• Assigning oxidation states is crucial for identifying oxidizing and reducing agents in reactions.

Galvanic Cells

Principles of Galvanic Cells

• Galvanic cells convert chemical energy into electrical energy through spontaneous redox reactions.

• A typical galvanic cell consists of two half-cells: an anode (oxidation) and a cathode (reduction).

• Example reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) illustrates the flow of electrons from zinc to copper.

Components of a Galvanic Cell

• Anode: The electrode where oxidation occurs (e.g., Zn(s) → Zn2+(aq) + 2e–).

• Cathode: The electrode where reduction occurs (e.g., Cu2+(aq) + 2e– → Cu(s)).

• Salt Bridge: Maintains electrical neutrality by allowing the flow of ions between the half-cells.

Redox Couples and Standard Reduction Potentials

• Standard reduction potentials (E°) measure the tendency of a species to be reduced, with the standard hydrogen electrode (SHE) set at 0.0 V.

• Example: Cu2+(aq) + 2e– → Cu(s) has E° = +0.34 V, indicating a strong oxidizing agent.

Gibbs Free Energy and Cell Potential

Relationship Between ΔG and E°cell

• The relationship is given by the equation ΔG° = -nFE°cell, where n is the number of moles of electrons transferred, F is Faraday's constant, and E°cell is the cell potential.

• A negative ΔG° indicates a spontaneous reaction, which corresponds to a positive E°cell.

Calculating Cell Potentials

• For the reaction Zn + Cu2+ → Zn2+ + Cu, the cell potential can be calculated as E°cell = E°(Cu2+/Cu) - E°(Zn2+/Zn).

• Example calculation: E°cell = 0.337 V - (-0.763 V) = +1.10 V, confirming the spontaneity of the reaction.

Applications of Gibbs Free Energy in Electrochemistry

• Gibbs free energy calculations help predict the feasibility of electrochemical reactions under standard and non-standard conditions.

• Understanding ΔG° allows for the design of more efficient batteries and fuel cells.

Concentration Effects on Cell Potential

Nernst Equation

• The Nernst equation relates cell potential to the concentrations of reactants and products: E = E° - (RT/nF)ln(Q), where Q is the reaction quotient.

• This equation allows for the calculation of cell potential under non-standard conditions, highlighting the effect of concentration on cell performance.

Impact of Concentration on Galvanic Cells

• Changes in concentration can significantly affect the cell potential; for instance, increasing the concentration of reactants generally increases the cell potential.

• Understanding these effects is crucial for optimizing the performance of electrochemical devices like batteries.

Practical Applications

• Real-world applications include lead storage batteries, dry cell batteries, and fuel cells, which all rely on the principles of electrochemistry and cell potential.

• The design and efficiency of these devices can be improved by manipulating concentration and understanding the underlying electrochemical principles.

Section 1: Fundamental Electrochemical Reactions

Detailed Key Concepts of Electrochemical Reactions

• Electrochemical reactions involve the transfer of electrons between species, often represented in half-reaction format.

• Example reaction: 2Al(s) + 3Mn2+(aq) → 2Al3+(aq) + 3Mn(s) illustrates the oxidation of aluminum and reduction of manganese ions.

• Le Chatelier's principle applies to these reactions, indicating that changes in concentration can shift the equilibrium position.

• The cell potential (E) is influenced by the concentrations of reactants and products, which can be calculated using the Nernst equation.

Gibbs Free Energy and Nernst Equation

• The relationship between Gibbs free energy (ΔG) and cell potential (E) is given by ΔG° = -nFE° and ΔG = ΔG° + RTlnQ.

• The Nernst equation, E = E° - (RT/nF)lnQ, shows how cell potential varies with concentration.

• At standard conditions (25°C), the Nernst equation simplifies to E = E° - 0.05916/n log Q, allowing for easier calculations.

• The equilibrium constant (K) can be derived from electrochemical measurements, linking thermodynamics and electrochemistry.

Concentration Cells and Their Potentials

• Concentration cells consist of two half-cells with different concentrations of the same ion, driving the reaction due to concentration gradients.

• Example: Ag+ + e– → Ag with E° = 0.80 V demonstrates how concentration differences create a potential difference.

• The Nernst equation is used to calculate the cell potential based on the concentration ratio, E = E° - (0.0591/n)log(Q).

• A concentration cell with 1 M Ag+ and 0.1 M Ag+ results in a small driving force, illustrating the concept of concentration gradients.

Section 2: Batteries and Energy Storage

Types of Batteries

• Galvanic cells can be connected in series to form batteries, with the total potential being the sum of individual cell potentials.

• Lead storage batteries consist of lead anodes, lead dioxide cathodes, and sulfuric acid electrolytes, providing ~2 V per cell.

• Dry cell batteries, such as zinc-carbon and alkaline batteries, are lightweight and efficient, with different half-reactions depending on the electrolyte used.

• Rechargeable batteries, like nickel-cadmium (Ni-Cad), allow for reversible reactions, although they pose toxicity issues due to cadmium.

Electrochemical Reactions in Batteries

• The overall reaction in a lead storage battery is Pb + PbO2 + 2H+ + 2HSO4– → 2PbSO4 + 2H2O, demonstrating the conversion of chemical energy to electrical energy.

• Dry cell batteries utilize reactions between zinc and manganese dioxide, producing a cell potential of approximately 1.5 volts.

• Alkaline batteries improve longevity by using KOH or NaOH as an electrolyte, reducing corrosion of the zinc anode.

• The efficiency of rechargeable batteries is highlighted by the reversibility of reactions, making them suitable for repeated use.

Section 3: Electrolysis and Corrosion

Electrolysis and Its Applications

• Electrolysis involves using electrical energy to drive non-spontaneous chemical reactions, such as the decomposition of water into hydrogen and oxygen.

• The half-reactions for water electrolysis are 2H2O → O2 + 4H+ + 4e– and 4H2O + 4e– → 2H2 + 4OH–, with a total cell potential of -2.06 V.

• Electroplating is a practical application of electrolysis, where metals are deposited onto surfaces, such as silver or copper.

• The time required for electroplating can be calculated using the formula: time = total charge (C) / current (A).

Corrosion and Its Prevention

• Corrosion is the oxidation of metals, often leading to structural failure, and is driven by electrochemical processes.

• The corrosion of iron involves anodic and cathodic regions, where iron oxidizes and electrons travel to react with oxygen, forming rust (Fe2O3·nH2O).

• Cathodic protection is a method to prevent corrosion by attaching a more reactive metal (e.g., magnesium) to the metal needing protection, sacrificing itself instead of the iron.

• Understanding the electrochemical potential helps in designing materials and coatings to resist corrosion effectively.