Galvanic (Voltaic) Cells and Electrochemistry Study Notes

Galvanic (Voltaic) Cells

Introduction to Galvanic Cells

Galvanic cells, also known as voltaic cells, are electrochemical cells that convert chemical energy into electrical energy through spontaneous redox reactions. They are fundamental in understanding how batteries work and are a key topic in introductory chemistry.

• Galvanic Cell: An electrochemical cell that generates electrical current from a spontaneous chemical reaction.

• Redox Reaction: A chemical reaction involving the transfer of electrons between two species.

• Electrodes: The two solid conductors (anode and cathode) where oxidation and reduction occur.

• Salt Bridge: A device that maintains electrical neutrality by allowing ions to flow between the two half-cells.

Components of a Galvanic Cell

• Anode: The electrode where oxidation occurs (loss of electrons).

• Cathode: The electrode where reduction occurs (gain of electrons).

• Half-Cell: Each compartment containing an electrode and an electrolyte solution.

• External Circuit: The wire connecting the two electrodes, allowing electron flow.

Cell Notation and Line Notation

Cell notation is a shorthand way to represent the reactions occurring in a galvanic cell.

• Format: Anode | Anode solution || Cathode solution | Cathode

• Example: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Standard Cell Potential ($E^ ext{cell}$)

The standard cell potential is the voltage produced by a galvanic cell under standard conditions (1 M concentration, 1 atm pressure, 25°C).

• Formula:

$E^ ext{cell} = E^ ext{cathode} - E^ ext{anode}$

• Positive $E^ ext{cell}$: Indicates a spontaneous reaction.

• Standard Reduction Potentials: Used to calculate $E^ ext{cell}$ from tables.

Electrochemical Reactions in Galvanic Cells

Oxidation and Reduction Half-Reactions

Each half-cell contains a half-reaction. The anode undergoes oxidation, and the cathode undergoes reduction.

• Oxidation (at Anode): $ext{M} ightarrow ext{M}^{n+} + n e^-$

• Reduction (at Cathode): $ext{M}^{n+} + n e^- ightarrow ext{M}$

• Example: In a cell with Al and Pb:

Oxidation: $ext{Al} ightarrow ext{Al}^{3+} + 3e^-$ Reduction: $ext{Pb}^{2+} + 2e^- ightarrow ext{Pb}$

Electron Flow and Ion Movement

• Electron Flow: Electrons move from the anode to the cathode through the external wire.

• Ion Flow: Cations move toward the cathode, anions move toward the anode through the salt bridge to maintain charge balance.

• Direction: Electrons always flow from the site of oxidation (anode) to the site of reduction (cathode).

Changes in Concentration

• Anode Solution: Concentration of metal ions increases as the metal is oxidized.

• Cathode Solution: Concentration of metal ions decreases as ions are reduced and deposited as solid metal.

Worked Examples and Applications

Example 1: Al/Pb Cell

• Cell Diagram: Al electrode in 1.0 M Al(NO3)3, Pb electrode in 1.0 M Pb(NO3)2, connected by a salt bridge.

• Standard Cell Potential: $1.53 ext{ V}$

• Half-Reactions:

Oxidation (Anode): $ext{Al} ightarrow ext{Al}^{3+} + 3e^-$ Reduction (Cathode): $ext{Pb}^{2+} + 2e^- ightarrow ext{Pb}$

• Electron Flow: From Al to Pb (right in the diagram).

• Concentration Changes: [Al3+] increases, [Pb2+] decreases.

• Line Notation: Al | Al3+ || Pb2+ | Pb

Example 2: Co/Ni Cell

• Cell Diagram: Co electrode in Co2+ solution, Ni electrode in Ni2+ solution.

• Standard Cell Potential: $0.03 ext{ V}$

• Half-Reactions:

Oxidation (Anode): $ext{Co} ightarrow ext{Co}^{2+} + 2e^-$ Reduction (Cathode): $ext{Ni}^{2+} + 2e^- ightarrow ext{Ni}$

• Electron Flow: From Co to Ni.

• Concentration Changes: [Co2+] increases, [Ni2+] decreases.

• Line Notation: Co | Co2+ || Ni2+ | Ni

Example 3: Fe/Mg Cell

• Cell Diagram: Fe electrode in Fe2+ solution, Mg electrode in Mg2+ solution.

• Standard Cell Potential: $1.93 ext{ V}$

• Half-Reactions:

Oxidation (Anode): $ext{Mg} ightarrow ext{Mg}^{2+} + 2e^-$ Reduction (Cathode): $ext{Fe}^{2+} + 2e^- ightarrow ext{Fe}$

• Electron Flow: From Mg to Fe.

• Concentration Changes: [Mg2+] increases, [Fe2+] decreases.

• Line Notation: Mg | Mg2+ || Fe2+ | Fe

Example 4: Zn/Ag Cell

• Cell Diagram: Zn electrode in Zn2+ solution, Ag electrode in Ag+ solution.

• Standard Cell Potential: $1.56 ext{ V}$

• Half-Reactions:

Oxidation (Anode): $ext{Zn} ightarrow ext{Zn}^{2+} + 2e^-$ Reduction (Cathode): $ext{Ag}^{+} + e^- ightarrow ext{Ag}$

• Electron Flow: From Zn to Ag.

• Concentration Changes: [Zn2+] increases, [Ag+] decreases.

• Line Notation: Zn | Zn2+ || Ag+ | Ag

Summary Table: Galvanic Cell Properties

Cell	Anode (Oxidation)	Cathode (Reduction)	Electron Flow	Standard Cell Potential (V)	Concentration Change	Line Notation
Al/Pb	Al	Pb	Al → Pb	1.53	[Al3+] ↑, [Pb2+] ↓	Al | Al3+ || Pb2+ | Pb
Co/Ni	Co	Ni	Co → Ni	0.03	[Co2+] ↑, [Ni2+] ↓	Co | Co2+ || Ni2+ | Ni
Fe/Mg	Mg	Fe	Mg → Fe	1.93	[Mg2+] ↑, [Fe2+] ↓	Mg | Mg2+ || Fe2+ | Fe
Zn/Ag	Zn	Ag	Zn → Ag	1.56	[Zn2+] ↑, [Ag+] ↓	Zn | Zn2+ || Ag+ | Ag

Key Terms and Concepts

• Anode: Site of oxidation, loses electrons.

• Cathode: Site of reduction, gains electrons.

• Salt Bridge: Maintains charge balance by allowing ion flow.

• Standard Cell Potential: Calculated using standard reduction potentials.

• Electron Flow: Always from anode to cathode.

Additional info:

• Standard reduction potentials are typically found in tables and are essential for calculating cell voltages.

• Galvanic cells are used in batteries, corrosion studies, and electrochemical sensors.