Voltaic Cells and Corrosion Notes

Voltaic Cells

• Voltaic cells (or galvanic cells) convert chemical energy into electrical energy using spontaneous oxidation-reduction reactions.

• Batteries and fuel cells are types of voltaic cells.

• Electrons move from the anode (where oxidation occurs) to the cathode (where reduction occurs) through an external wire.

• Ion movement in the solution balances the electron flow.

• Common dry cell types include zinc-carbon, alkaline, and mercury batteries.

Types of Dry Cells

• Zinc-Carbon Dry Cells:

  • Zinc container acts as the anode: $Zn(s) → Zn^{2+}(aq) + 2e^-$

  • Cathode involves reduction of $MnO<em>2$: $2MnO</em>2(s) + H<em>2O(l) + 2e^- → Mn</em>2O_3(s) + 2OH^-(aq)$

• Alkaline Batteries:

  • Use a paste of $Zn$ metal and potassium hydroxide instead of a solid metal anode.

  • Anode half-reaction: $Zn(s) + 2OH^-(aq) → Zn(OH)_2(s) + 2e^-$

  • Cathode reaction is the same as in zinc-carbon cells.

• Mercury Batteries:

  • Anode reaction is the same as in alkaline batteries.

  • Cathode half-reaction: $HgO(s) + H_2O(l) + 2e^- → Hg(l) + 2OH^-(aq)$

Fuel Cells

• Fuel cells continuously supply reactants and remove products.

• Reactions used in space program fuel cells:

  • Cathode: $O<em>2(g) + 2H</em>2O(l) + 4e^- → 4OH^-(aq)$

  • Anode: $2H<em>2(g) + 4OH^-(aq) → 4e^- + 4H</em>2O(l)$

  • Net: $2H<em>2 + O</em>2 → 2H_2O$

• Fuel cells are efficient and have low emissions.

Corrosion

• Corrosion is an electrochemical process with significant economic impact.

• Rust (hydrated iron(III) oxide) forms via: $4Fe(s) + 3O<em>2(g) + xH</em>2O(l) → 2Fe<em>2O</em>3•xH_2O(s)$

• Mechanism:

  • Anode: $Fe(s) → Fe^{2+}(aq) + 2e^-$

  • Cathode: $O<em>2(g) + 2H</em>2O(l) + 4e^- → 4OH^-(aq)$

• Water and oxygen must be present for corrosion.

Preventing Corrosion

• Methods include painting metal or galvanizing (coating with zinc).

• Cathodic protection uses a sacrificial anode (e.g., zinc) that is more easily oxidized than iron.

Electric Potential

• Electric potential (voltage) is the driving force on electrons, measured in volts (V).

• Current is the movement of electrons, measured in amperes (A).

• Electrode potential is the potential difference between an electrode and its solution.

Electrode Potentials

• Standard hydrogen electrode (SHE) is a reference with an arbitrary potential of $0.00 V$.

  • Anodic reaction: $H_2(g) → 2H^+(aq) + 2e^-$

• Standard electrode potential ($E^0$) is measured relative to SHE, expressed as reduction potentials.

• Effective oxidizing agents have positive $E^0$ values; effective reducing agents have negative $E^0$ values.

• When a half-reaction is written as oxidation, the sign of $E^0$ is reversed.

Calculating Cell Potential

• A spontaneous reaction will have a positive value for $E^0<em>{cell}$, calculated as: $E^0</em>{cell} = E^0<em>{cathode} - E^0</em>{anode}$

• The half-reaction with the more negative standard reduction potential will be the anode.