Electrochemistry: Oxidation-Reduction, Galvanic Cells, and Applications

Oxidation and Reduction

Introduction to Redox Reactions

Redox (reduction-oxidation) reactions are fundamental chemical processes involving the transfer of electrons between substances. These reactions are essential in both natural and industrial processes, including corrosion, metabolism, and energy production in batteries.

• Oxidation: The loss of electrons by a substance.

• Reduction: The gain of electrons by a substance.

• Oxidizing agent: The substance that accepts electrons (is reduced).

• Reducing agent: The substance that donates electrons (is oxidized).

Example: The reaction of magnesium with oxygen:

$2\text{Mg}(s) + \text{O}_2(g) \rightarrow 2\text{MgO}(s)$

• Mg is oxidized (loses electrons), O2 is reduced (gains electrons).

Assigning Oxidation Numbers

Oxidation numbers are used to keep track of electron transfer in redox reactions. They are assigned based on a set of rules:

Rule	Description
1	The oxidation number of an atom in its elemental form is 0.
2	The oxidation number of a monatomic ion is equal to its charge.
3	Oxygen is usually -2, hydrogen is +1, fluorine is -1 in compounds.
4	The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion's charge.

Example: Assign oxidation numbers in $\text{K}_2\text{Cr}_2\text{O}_7$:

• K: +1, O: -2, Cr: +6 (by solving $2(+1) + 2(x) + 7(-2) = 0$)

Balancing Redox Equations

Net Ionic Equations for Redox Reactions

Redox equations must be balanced for both mass and charge. The half-reaction method is commonly used:

1. Write separate half-reactions for oxidation and reduction.

2. Balance all elements except H and O.

3. Balance O by adding H2O; balance H by adding H+ (in acidic solution).

4. Balance charge by adding electrons.

5. Multiply half-reactions to equalize electrons, then add together.

Example:

$\text{MnO}_4^- + 8\text{H}^+ + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O}$

Redox Reactions in Acidic and Basic Solutions

• In acidic solutions, use H+ and H2O to balance H and O.

• In basic solutions, after balancing as if acidic, add OH- to both sides to neutralize H+ and form water.

Galvanic (Voltaic) Cells

Structure and Function

Galvanic cells convert chemical energy from spontaneous redox reactions into electrical energy. They consist of two half-cells connected by a salt bridge, allowing ion flow to maintain charge balance.

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

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

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

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

Cell Potential and Standard Reduction Potentials

The cell potential (Ecell) is the voltage produced by a galvanic cell. It is calculated using standard reduction potentials (Eo) from tables:

$E_{cell}^o = E_{cathode}^o - E_{anode}^o$

Standard reduction potentials are measured under standard conditions (1 M, 1 atm, 25°C).

Table: Standard Reduction Potentials (Selected)

Half-Reaction	Eo (V)
F2 + 2e- → 2F-	+2.87
Cl2 + 2e- → 2Cl-	+1.36
Cu2+ + 2e- → Cu	+0.34
Zn2+ + 2e- → Zn	-0.76
Li+ + e- → Li	-3.04

Additional info: The full table includes many more half-reactions, ordered by their tendency to be reduced.

Corrosion

Process and Prevention

Corrosion is the deterioration of metals due to redox reactions with substances in the environment, such as oxygen and water. The most common example is the rusting of iron:

$4\text{Fe}(s) + 3\text{O}_2(g) + 6\text{H}_2\text{O}(l) \rightarrow 4\text{Fe}^{3+}(aq) + 6\text{OH}^-(aq)$

• Prevention methods include coating, galvanization (zinc coating), and cathodic protection.

Electrolysis

Principles and Applications

Electrolysis uses electrical energy to drive non-spontaneous chemical reactions. It is used for processes such as electroplating, extraction of metals, and the decomposition of compounds.

• At the cathode, reduction occurs (gain of electrons).

• At the anode, oxidation occurs (loss of electrons).

Example: Electrolysis of molten sodium chloride:

$2\text{NaCl}(l) \rightarrow 2\text{Na}(l) + \text{Cl}_2(g)$

Comparison of Electrolytic and Galvanic Cells

• Galvanic cells: Spontaneous reactions produce electricity.

• Electrolytic cells: Electricity is used to drive non-spontaneous reactions.

Batteries

Overview

Batteries are practical applications of galvanic cells, providing portable sources of electrical energy. They consist of one or more electrochemical cells and are classified as primary (non-rechargeable) or secondary (rechargeable) batteries.

• Common types: Dry cell, lead-acid battery, lithium-ion battery.

Essential Biological Redox Reactions

Photosynthesis and Cellular Respiration

Redox reactions are central to energy conversion in living organisms:

• Photosynthesis: Converts CO2 and H2O into glucose and O2 using sunlight.

• Cellular respiration: Converts glucose and O2 into CO2, H2O, and energy.

$6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$

$\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy}$