Electrochemistry: Redox Reactions, Galvanic and Electrolytic Cells, and Applications

Electrochemistry

Introduction to Electrochemistry

Electrochemistry is the branch of chemistry that studies the relationship between electricity and chemical reactions, particularly those involving electron transfer (redox reactions). It is fundamental to understanding batteries, corrosion, electrolysis, and many industrial processes.

Redox Basics & Agents

Oxidation and Reduction

Redox (reduction-oxidation) reactions involve the transfer of electrons between chemical species. The substance that loses electrons is oxidized, while the substance that gains electrons is reduced.

• Oxidation: Loss of electrons; oxidation state becomes more positive or less negative.

• Reduction: Gain of electrons; oxidation state becomes more negative or less positive.

• Oxidizing Agent: The species that causes oxidation by accepting electrons (itself is reduced).

• Reducing Agent: The species that causes reduction by donating electrons (itself is oxidized).

Examples of Redox Reactions in Industry

• Claus Process: Used to recover sulfur from natural gas and crude oil. The overall reaction is: $2\mathrm{H}_2\mathrm{S} + \mathrm{O}_2 \rightarrow 2\mathrm{S} + 2\mathrm{H}_2\mathrm{O}$

• Haber Process: Synthesis of ammonia for fertilizers and explosives: $\mathrm{N}_2(g) + 3\mathrm{H}_2(g) \rightarrow 2\mathrm{NH}_3(g)$

Identifying Oxidation and Reduction in Reactions

To determine which element is oxidized or reduced, compare oxidation states before and after the reaction. The oxidizing agent is reduced, and the reducing agent is oxidized.

• Example: $\mathrm{MnO}_4^- + 5\mathrm{Fe}^{2+} + 8\mathrm{H}^+ \rightarrow \mathrm{Mn}^{2+} + 5\mathrm{Fe}^{3+} + 4\mathrm{H}_2\mathrm{O}$ - Iron is oxidized (Fe2+ to Fe3+), manganese is reduced (MnO4- to Mn2+).

Oxidizing and Reducing Agents in Everyday Substances

Common oxidizing agents include O2, Cl2, and KMnO4. Reducing agents include H2, metals like Zn, and organic compounds like gasoline.

Changes in Oxidation State

Recognizing Redox Reactions

Redox reactions are characterized by changes in oxidation states. Combustion, rusting, and reactions involving pure elements are always redox processes.

• Combustion: $\mathrm{CH}_3\mathrm{OH}(l) + \mathrm{O}_2(g) \rightarrow \mathrm{CO}_2(g) + \mathrm{H}_2\mathrm{O}(g)$

• Rusting: $\mathrm{Fe}(s) \rightarrow \mathrm{Fe}^{2+}(aq) + 2e^-$

Balancing Aqueous-phase Redox Equations

Steps for Balancing Redox Reactions

Redox reactions in aqueous solution are balanced by separating them into half-reactions, balancing atoms and charges, and then combining them.

1. Write oxidation and reduction half-reactions.

2. Balance all elements except H and O.

3. Balance O by adding H2O, H by adding H+ (in acidic solution), and charge by adding electrons.

4. Multiply half-reactions to equalize electrons and add together.

Galvanic (Voltaic) Cells

Structure and Function

Galvanic cells convert chemical energy into electrical energy through spontaneous redox reactions. They consist of two half-cells connected by a salt bridge and an external circuit.

• Anode: Site of oxidation (loses electrons).

• Cathode: Site of reduction (gains electrons).

• Salt Bridge: Maintains electrical neutrality by allowing ion flow.

• Electron Flow: From anode to cathode through the external wire.

Cell Notation and Calculations

Cell notation summarizes the components of a galvanic cell. The standard cell potential ($E^\circ_{cell}$) is calculated from standard reduction potentials:

• $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

Standard Reduction Potentials

Definition and Use

Standard reduction potentials ($E^\circ$) measure the tendency of a species to gain electrons under standard conditions (1 M, 1 atm, 298 K). The more positive the $E^\circ$, the greater the tendency to be reduced.

• Standard hydrogen electrode (SHE) is assigned $E^\circ = 0.00$ V.

• Tables of $E^\circ$ values are used to predict cell voltages and spontaneity.

Relationship Between ΔG°, K, and E°cell

Thermodynamic Connections

The spontaneity of electrochemical reactions is related to the standard cell potential, equilibrium constant, and standard free energy change:

• $\Delta G^\circ = -n F E^\circ_{cell}$

• $\Delta G^\circ = -RT \ln K$

• $E^\circ_{cell} = \frac{0.0592\,\mathrm{V}}{n} \log K$ (at 25°C)

Nernst Equation and Non-Standard Cells

Nernst Equation

The Nernst equation allows calculation of cell potential under non-standard conditions:

• $E = E^\circ - \frac{RT}{nF} \ln Q$

• At 25°C: $E = E^\circ - \frac{0.0592\,\mathrm{V}}{n} \log Q$

• Q is the reaction quotient.

Corrosion

Mechanism of Corrosion

Corrosion is the deterioration of metals by redox reactions with environmental agents, such as oxygen and water. Iron rusting is a common example, involving the oxidation of Fe to Fe2+ and subsequent reactions to form hydrated iron(III) oxide (rust).

Electrolytic Cells and Faraday’s Law

Electrolytic Cells

Electrolytic cells use electrical energy to drive non-spontaneous redox reactions. They are used in electroplating, metal purification, and industrial processes like the chloralkali process.

• Cathode: Site of reduction (gains electrons).

• Anode: Site of oxidation (loses electrons).

Faraday’s Law of Electrolysis

Faraday’s Law relates the amount of substance produced at an electrode to the quantity of electricity passed through the cell:

• $\text{moles plated} = \frac{I \cdot t}{n \cdot \mathcal{F}}$

• I = current (A), t = time (s), n = number of electrons, $\mathcal{F}$ = Faraday’s constant (96500 C/mol e-).

Applications and Industrial Processes

Industrial Electrochemistry

• Chloralkali Process: Electrolysis of brine to produce Cl2, H2, and NaOH.

• Aluminum Production: Electrolysis of molten Al2O3 to obtain pure aluminum metal.

Summary Table: Key Electrochemical Concepts

Concept	Definition	Key Equation
Oxidation	Loss of electrons	—
Reduction	Gain of electrons	—
Cell Potential ($E^\circ_{cell}$)	Voltage produced by a cell	$E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$
Gibbs Free Energy ($\Delta G^\circ$)	Spontaneity of reaction	$\Delta G^\circ = -n F E^\circ_{cell}$
Nernst Equation	Cell potential under non-standard conditions	$E = E^\circ - \frac{0.0592}{n} \log Q$
Faraday’s Law	Amount of substance produced in electrolysis	$\text{moles} = \frac{I t}{n \mathcal{F}}$