Electrochemistry: Galvanic and Electrolytic Cells, Concentration Cells, Fuel Cells, and Electrolysis

Electrochemistry

Overview

Electrochemistry is the study of chemical processes that cause electrons to move, which is the basis for the generation of electricity in batteries and the use of electricity to drive chemical reactions. This field is essential for understanding batteries, fuel cells, corrosion, and electrolysis.

Galvanic (Voltaic) Cells

Definition and Function

• Galvanic (Voltaic) cells are electrochemical cells that convert chemical energy into electrical energy through spontaneous redox reactions.

• These cells are the basis for batteries and many types of fuel cells.

Key Components

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

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

• Salt bridge: Maintains electrical neutrality by allowing ion flow between half-cells.

• Electron flow: From anode to cathode through an external circuit.

Example Reaction

• $\text{Zn}(s) + 2\text{H}^+(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{H}_2(g)$

• Standard cell notation: Zn(s) | Zn2+(aq) || H+(aq) | H2(g) | Pt(s)

Cell Potentials and the Nernst Equation

Standard Electrode Potentials

• Standard reduction potentials (E0): Measured under standard conditions (1 M, 1 atm, 25°C) relative to the standard hydrogen electrode (SHE).

• Cell potential (Ecell): $E_{cell} = E_{cathode}^0 - E_{anode}^0$

Nernst Equation

• Used to calculate cell potential under non-standard conditions.

$E_{cell} = E_{cell}^0 - \frac{0.0592\ V}{n} \log Q$

• n: Number of moles of electrons transferred.

• Q: Reaction quotient.

Concentration Cells

Definition and Operation

• A concentration cell is a voltaic cell in which the voltage is generated due to a difference in ion concentrations between two half-cells.

• Both electrodes are the same material, but the ion concentrations differ.

Example

• Cell notation: Cu(s) | Cu2+ (0.05 M) || Cu2+ (0.50 M) | Cu(s)

• Potential calculated using the Nernst equation:

$E_{cell} = 0 - \frac{0.0592}{2} \log \left( \frac{0.05}{0.50} \right) = 0.030\ V$

Types of Electrochemical Cells

Voltaic (Galvanic) vs. Electrolytic Cells

• Voltaic (Galvanic) Cell: Releases free energy from a spontaneous reaction to produce electricity (Ecell > 0, ΔG < 0).

• Electrolytic Cell: Absorbs free energy from an external source to drive a non-spontaneous reaction (Ecell < 0, ΔG > 0).

Comparison Table

Feature	Voltaic Cell	Electrolytic Cell
Spontaneity	Spontaneous	Non-spontaneous
Energy Flow	Produces electricity	Requires external power
Electron Flow	Anode to cathode	Anode to cathode
Sign of Ecell	Positive	Negative

Fuel Cells

Definition and Application

• Fuel cells are electrochemical cells that convert the chemical energy of a fuel (often hydrogen) and an oxidizing agent (often oxygen) into electricity through a pair of redox reactions.

• Example: Hydrogen fuel cell in vehicles, where H2 and O2 react to form water and release energy.

Electrolysis

Definition and Principles

• Electrolysis is the process of using electrical energy to drive a non-spontaneous chemical reaction.

• Commonly used for electroplating, metal extraction, and water splitting.

Electrolysis of Water

• Overall reaction: $2\text{H}_2\text{O}(l) \rightarrow 2\text{H}_2(g) + \text{O}_2(g)$

• At the anode (oxidation): $2\text{H}_2\text{O}(l) \rightarrow \text{O}_2(g) + 4\text{H}^+(aq) + 4e^-$

• At the cathode (reduction): $4\text{H}_2\text{O}(l) + 4e^- \rightarrow 2\text{H}_2(g) + 4\text{OH}^-(aq)$

Stoichiometry of Electrolysis and Faraday's Law

Faraday's Law of Electrolysis

• The amount of substance produced at each electrode is directly proportional to the quantity of charge passed through the cell.

• 1 Faraday (F) = 96,485 C/mol e-

Steps for Calculations

1. Balance half-reactions to find moles of electrons per mole of product.

2. Use Faraday's constant to convert charge to moles of electrons.

3. Use molar mass to find the mass of product formed.

Example Calculation

• How many moles of H2(g) will be produced if 1.5 V and a current of 100 A is applied for 5 minutes?

$5\ min \times \frac{60\ s}{1\ min} \times 100\ C/s = 30,000\ C$

$\frac{30,000\ C}{96,485\ C/mol\ e^-} = 0.31\ mol\ e^-$

$0.31\ mol\ e^- \times \frac{2\ mol\ H_2}{4\ mol\ e^-} = 0.155\ mol\ H_2$

Applications and Practice Problems

• Calculate the time required to produce 1 mole of O2 with a current of 500 A.

• Practice using the Nernst equation for concentration cells.

• Understand the operation and applications of fuel cells in vehicles.

Key Terms and Definitions

• Electrochemical cell: Device that generates electrical energy from chemical reactions or uses electrical energy to drive chemical reactions.

• Anode: Electrode where oxidation occurs.

• Cathode: Electrode where reduction occurs.

• Salt bridge: Device used to maintain electrical neutrality in electrochemical cells.

• Faraday's constant (F): The charge of one mole of electrons, 96,485 C/mol e-.

• Nernst equation: Equation used to calculate cell potential under non-standard conditions.

• Electrolysis: Use of electrical energy to drive a non-spontaneous chemical reaction.