CH 112: Electrochemical Cells and Potential

Unit 7: Topics 11 - 16

Topic 12: Electrochemical Cells and Potential

Electrochemistry Overview

• Definition: Electrochemistry is the branch of chemistry that examines the transformations between chemical and electrical energy.

• Redox Reactions: These reactions are the sum of two half-reactions:
    - Reduction Half-Reaction: A reactant gains electrons.
    - Oxidation Half-Reaction: A reactant loses electrons.

• Key Point: Reduction and oxidation half-reactions occur simultaneously, and the number of electrons gained during reduction must exactly match the number lost during oxidation.

Oxidation-Reduction Reactions

• Examples of Oxidation-Reduction Reactions:
    - $ext{Zn}(s) <br>ightarrow ext{Zn}^{2+}(aq) + 2 e^-$ (Oxidation)
    - $ext{Cu}^{2+}(aq) + 2 e^- <br>ightarrow ext{Cu}(s)$ (Reduction)
    - Combined Reaction: $ext{Zn}(s) + ext{Cu}^{2+}(aq) <br>ightarrow ext{Zn}^{2+}(aq) + ext{Cu}(s)$

Electrochemical Cell

• Definition: An electrochemical cell is a device that converts chemical energy into electrical work or electrical work into chemical energy.

• Voltaic Cell: A type of electrochemical cell where chemical energy is transformed into electrical energy through a spontaneous redox reaction (commonly known as a battery).

• Cell Diagram: A visual representation (symbols) that shows how the components of an electrochemical cell are connected.

Cell Components

• Electrodes:
    - Anode: The electrode where oxidation (loss of electrons) occurs.
    - Cathode: The electrode where reduction (gain of electrons) occurs.

• Salt Bridge: A component that connects the two solutions of the cell, balances the flow of electrons, and eliminates the accumulation of charge in either compartment.

Voltaic vs. Electrolytic Cell

Voltaic Cell

• Spontaneous cell reaction converts chemical energy into electrical energy.

Electrolytic Cell

• Electrical energy is used to drive a nonspontaneous cell reaction.

• Anode: Negative terminal (oxidation).

• Cathode: Positive terminal (reduction).

• Power Supply: Supplies energy required to drive the reaction.

• Background Electrolyte: Maintains ionic balance in the solution.

Mass Change in Electrodes

• At the Anode:
    - The zinc anode loses mass due to oxidation: $ext{Zn}(s) <br>ightarrow ext{Zn}^{2+}(aq) + 2 e^-$

• At the Cathode:
    - The copper cathode gains mass due to reduction: $ext{Cu}^{2+}(aq) + 2 e^- <br>ightarrow ext{Cu}(s)$

Writing Cell Diagrams

1. Positioning: Write the chemical symbol of the anode on the far left and the cathode on the far right, with a double vertical line indicating the salt bridge between them.
      - For Example: $ext{Zn}(s) ext{ . . . . . } ext{ || } ext{ . . . . . } ext{Cu}(s)$

2. Representing Phase Changes: Work from the electrodes toward the bridge, using vertical lines to indicate phase changes and noting the ions or compounds that represent the electrolytes surrounding the electrode.
      - Example: $ext{Zn}(s) | ext{Zn}^{2+}(aq) || ext{Cu}^{2+}(aq) | ext{Cu}(s)$

3. Indicate Concentrations: Note concentrations of dissolved species and partial pressures of gases if applicable.
      - Example: $ext{Zn}(s) | ext{Zn}^{2+}(1.00 M) || ext{Cu}^{2+}(1.00 M) | ext{Cu}(s)$

Standard Potentials

• Standard Reduction Potential (E°): The potential of a half-reaction under standard conditions (all reactants and products in their standard states at 25°C).

• Standard Cell Potential (E° cell): Measures how forcefully an electrochemical cell can pump electrons through an external circuit; calculated as:
    - $E°<em>{cell} = E°</em>{cathode} - E°_{anode}$

• Standard Conditions: Includes concentrations of 1 M and a partial pressure of gases at 1 bar.

Calculating Standard Potentials

Example Calculation for E° cell:

  - $E°<em>{cathode} [Cu^{2+} ightarrow Cu(s)] = 0.342 V$   - $E°</em>{anode} [Zn^{2+} <br>ightarrow Zn(s)] = -0.762 V$
  - Calculation:
  $E°_{cell} = (0.342 V) - (-0.762 V) = 1.104 V$

Zinc-Air Battery

• Chemical Reaction: $2 ext{Zn}(s) + ext{O}_2(g) <br>ightarrow 2 ext{ZnO}(s)$

• Standard Cell Potentials:
    - $E°<em>{cathode} = 0.401 V$   - $E°</em>{anode} = -1.25 V$

• Calculation of Standard Cell Potential:
    - $E°_{cell} = (0.401 V) - (-1.25 V) = 1.65 V$

Faraday's Laws and Gibbs Free Energy

• Equations:
    - $ext{ΔG}<em>{cell} = -nF E</em>{cell}$
    - Where:
      - F = Faraday constant = $9.65 imes 10^4 ext{ C/(mol e^-)}$
      - n = number of moles of electrons transferred in the reaction.

Nernst Equation

• Used for calculating cell potential under non-standard conditions:
    - $ext{ΔG} = ext{ΔG}° + RT ext{ln} Q$
    - $E_{cell} = E_{cell}° - rac{RT}{nF} ext{ln} Q$

• At 298 K, this can be simplified to:
    - $E_{cell} = E_{cell}° - 0.0592 rac{V}{n} ext{log} Q$

Corrosion and Inhibition

• Corrosion: A process where a metal is oxidized by substances in its environment:
    - Example Reaction: $4 ext{Fe}(s) + 3 ext{O}_2(g) <br>ightarrow 2 ext{Fe}_2 ext{O}_3(s)$

• Factors Promoting Corrosion:
    - Presence of water
    - Presence of electrolytes
    - Contact between dissimilar metals

• Corrosion Inhibition Techniques:
    - Cathodic Protection with sacrificial anodes.

Solutions and Applications

• Lead-Acid Batteries: Reactions in lead-acid batteries during discharge and recharge.

• Electrolysis and Electroplating: Using electricity to drive a nonspontaneous redox reaction, with applications in metal plating and refining.

• Fuel Cells: Allows for direct conversion of chemical energy from fuel into electricity with byproducts of water and heat.