Ch. 19 Electrochemistry

Redox Reactions

Electrochemistry is the study of chemical processes that cause electrons to move, which is the basis for redox (reduction-oxidation) reactions. These reactions are fundamental to understanding how chemical energy is converted to electrical energy and vice versa.

• Oxidation: The loss of electrons by a species. The species that loses electrons is said to be oxidized.

• Reduction: The gain of electrons by a species. The species that gains electrons is said to be reduced.

• Oxidizing Agent: The substance that causes oxidation by accepting electrons (itself is reduced).

• Reducing Agent: The substance that causes reduction by donating electrons (itself is oxidized).

• Example: In the reaction , Mn is reduced and C is oxidized.

Balancing Redox Reactions

Redox reactions must be balanced for both mass and charge. The process differs depending on whether the reaction occurs in acidic or basic solution.

• Steps in Acidic Solution:

  1. Write the oxidation and reduction half-reactions.

  2. Balance all elements except H and O.

  3. Balance O by adding .

  4. Balance H by adding .

  5. Balance charge by adding electrons ().

  6. Combine the half-reactions and cancel electrons.

• Steps in Basic Solution:

  1. Balance as in acidic solution.

  2. Add to both sides to neutralize , forming .

  3. Simplify and balance.

• Example: Balancing in acidic solution.

Electrochemical Cells

Electrochemical cells convert chemical energy into electrical energy (or vice versa) through redox reactions. There are two main types:

• Voltaic (Galvanic) Cell: Produces electrical current from a spontaneous chemical reaction.

• Electrolytic Cell: Uses electrical current to drive a nonspontaneous chemical reaction.

• Cell Notation: Represents the components and reactions in an electrochemical cell. For example, .

• Line Notation: Anode (oxidation) is written on the left, cathode (reduction) on the right. Double vertical lines () indicate a salt bridge.

Key Components

• Anode: Electrode where oxidation occurs; electrons are released.

• Cathode: Electrode where reduction occurs; electrons are accepted.

• Salt Bridge: Maintains electrical neutrality by allowing ion flow.

Standard Cell Potential ()

The standard cell potential is the voltage produced by an electrochemical cell under standard conditions (1 M, 1 atm, 25°C).

• Calculation:

• Example: For and half-cells, use tabulated values to calculate .

Gibbs Free Energy and Cell Potential

The relationship between cell potential and Gibbs free energy is given by:

• Where:

  • = number of moles of electrons transferred

  • = Faraday's constant ( C/mol )

  • = standard cell potential (V)

• Example: If V and , J

Equilibrium Constant and Cell Potential

The equilibrium constant () for a redox reaction is related to the standard cell potential:

• (at 25°C)

• Example: Calculate for a reaction with V and .

Electrolysis and Faraday's Laws

Electrolysis is the process of driving a nonspontaneous reaction using electrical energy. Faraday's laws relate the amount of substance produced at an electrode to the amount of electric charge passed.

• Key Formula:

  • Where:

    • = current (A)

    • = time (s)

    • = molar mass (g/mol)

    • = number of electrons

    • = Faraday's constant

• Example: Calculate the mass of Ag plated in 67 minutes at 8.70 A.

Electrodes and Cell Construction

Electrodes are materials that conduct electricity and participate in redox reactions. A sacrificial electrode is used to prevent corrosion of another metal by preferentially oxidizing.

• Sacrificial Electrode: A metal that oxidizes more easily than the protected metal, often used in cathodic protection.

• Example: Iron is protected from corrosion by attaching a more easily oxidized metal such as zinc.

Summary Table: Electrochemical Cell Components

Additional info:

• These notes expand upon the practice questions by providing definitions, formulas, and context for key electrochemistry concepts.

• Students should be familiar with identifying oxidation and reduction, balancing redox reactions, calculating cell potentials, and applying Faraday's laws for quantitative electrolysis problems.