CHEM122 Electrochemistry

Introduction to 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.

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.

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 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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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 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.

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.

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.

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 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.