Electrochemistry: Redox Reactions, Electrochemical Cells, and Applications

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 $\text{MnO}_4^- (aq) + \text{H}_2\text{C}_2\text{O}_4 (aq) \rightarrow \text{Mn}^{2+} (aq) + \text{CO}_2 (g)$, 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 $\text{H}_2\text{O}$.

  4. Balance H by adding $\text{H}^+$.

  5. Balance charge by adding electrons ($e^-$).

  6. Combine the half-reactions and cancel electrons.

• Steps in Basic Solution:

  1. Balance as in acidic solution.

  2. Add $\text{OH}^-$ to both sides to neutralize $\text{H}^+$, forming $\text{H}_2\text{O}$.

  3. Simplify and balance.

• Example: Balancing $\text{Cr}_2\text{O}_7^{2-} + \text{H}_2\text{O}_2 \rightarrow \text{Cr}^{3+} + \text{O}_2$ 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, $\text{Zn}(s) | \text{Zn}^{2+}(aq) || \text{Cu}^{2+}(aq) | \text{Cu}(s)$.

• 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 ($E^\circ_{cell}$)

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

• Calculation:

  • $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

• Example: For $\text{Fe}^{3+}/\text{Fe}^{2+}$ and $\text{Cl}_2/\text{Cl}^-$ half-cells, use tabulated $E^\circ$ values to calculate $E^\circ_{cell}$.

Gibbs Free Energy and Cell Potential

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

• $\Delta G^\circ = -n F E^\circ_{cell}$

• Where:

  • $n$ = number of moles of electrons transferred

  • $F$ = Faraday's constant ($96,485$ C/mol $e^-$)

  • $E^\circ_{cell}$ = standard cell potential (V)

• Example: If $E^\circ_{cell} = 1.10$ V and $n = 2$, $\Delta G^\circ = -2 \times 96,485 \times 1.10$ J

Equilibrium Constant and Cell Potential

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

• $E^\circ_{cell} = \frac{0.0592}{n} \log K$ (at 25°C)

• Example: Calculate $K$ for a reaction with $E^\circ_{cell} = 0.54$ V and $n = 2$.

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:

  • $\text{mass} = \frac{ItM}{nF}$

  • Where:

    • $I$ = current (A)

    • $t$ = time (s)

    • $M$ = molar mass (g/mol)

    • $n$ = number of electrons

    • $F$ = 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

Component	Function	Location
Anode	Oxidation occurs; electrons released	Left side in cell notation
Cathode	Reduction occurs; electrons accepted	Right side in cell notation
Salt Bridge	Maintains charge balance	Between half-cells
Electrolyte	Conducts ions	In each half-cell

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.