Galvanic (Voltaic) Cells and Electrochemistry

Galvanic (Voltaic) Cells

Introduction to Galvanic Cells

A Galvanic cell (also called a voltaic cell) is a type of electrochemical cell that produces electricity through a spontaneous redox reaction. These cells are commonly used in batteries to convert chemical energy into electrical energy.

• Purpose: To generate electricity from spontaneous chemical reactions.

• Key Function: Converts chemical energy into electrical energy.

Example: The main purpose of a galvanic cell is to generate electricity (not to consume electricity or allow only oxidation).

Galvanic Cell Components

The main components of a galvanic cell are:

• Electrodes:

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

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

• Salt Bridge: A tube containing an electrolyte solution that allows ions to flow between the two half-cells, maintaining electrical neutrality.

• Voltmeter: Measures the potential difference (voltage) generated by the cell.

Diagram: A typical galvanic cell consists of two half-cells connected by a salt bridge and external wire, with a voltmeter measuring the cell potential.

Example: Galvanic Cell Reactions

• Given the reaction: $\text{Mg (s)} + \text{Cu}^{2+} (aq) \rightarrow \text{Mg}^{2+} (aq) + \text{Cu (s)}$

• Key Points:

  • Electrons flow from the anode (Mg) to the cathode (Cu).

  • Mg is oxidized (loses electrons), Cu2+ is reduced (gains electrons).

  • The salt bridge allows ions to move and maintain charge balance.

Galvanic Cell Electrodes and Electron Flow

Half-Cell Reactions

Each galvanic cell has two half-cells: one for oxidation and one for reduction.

• Anode: Site of oxidation; electrons are released and flow through the external circuit to the cathode.

• Cathode: Site of reduction; electrons are accepted from the external circuit.

Electron Flow: Always from anode to cathode.

Example: Calculating Electron Transfer

• Given: $\text{Ga}^{3+} + 3e^- \rightarrow \text{Ga}$

• Number of electrons transferred per Ga atom reduced: 3

Cell Potential and Spontaneity

Cell Potential ($E_{cell}$)

The cell potential is the voltage produced by a galvanic cell. It is calculated as:

• $E_{cell} = E_{cathode} - E_{anode}$

A positive $E_{cell}$ indicates a spontaneous reaction.

Spontaneity of Redox Reactions

• Spontaneous redox reactions have a positive $E_{cell}$ and a negative $\Delta G$ (Gibbs free energy change).

• Nonspontaneous reactions require external energy input (electrolytic cells).

Type of Cell	Spontaneity	$E_{cell}$	$\Delta G$
Galvanic (Voltaic)	Spontaneous	Positive	Negative
Electrolytic	Nonspontaneous	Negative	Positive

Example: Equilibrium Constant and Cell Potential

• A reduction reaction with an equilibrium constant $K = 4.4 \times 10^8$ is spontaneous and has a positive $E_{cell}$.

Practice Problems and Applications

Identifying Anode and Cathode

• The anode is where oxidation occurs; the cathode is where reduction occurs.

• Electrons flow from anode to cathode through the external circuit.

• Cations migrate toward the cathode; anions migrate toward the anode via the salt bridge.

Example: Nickel and Silver Cell

• Given: $\text{Ni (s)} + 2\text{Ag}^+ (aq) \rightarrow \text{Ni}^{2+} (aq) + 2\text{Ag (s)}$

• At the anode: Ni is oxidized to Ni2+ (loses electrons).

• At the cathode: Ag+ is reduced to Ag (gains electrons).

• Electrons flow from Ni to Ag+.

Summary Table: Galvanic Cell Key Features

Component	Process	Charge	Electron Flow
Anode	Oxidation	Negative	Releases electrons
Cathode	Reduction	Positive	Accepts electrons

Additional info: In a galvanic cell, the anode is negative because it is the source of electrons, while the cathode is positive. The salt bridge is essential for maintaining charge balance by allowing ion migration. The cell potential can be calculated using standard reduction potentials from a reference table.