14 Redox reaction

Redox Overview

  • Redox (reduction-oxidation) reactions: involve the transfer of electrons between substances.

Electrochemistry

  • Definition: Branch of chemistry focusing on the conversion between chemical energy and electrical energy.

Electrochemical Processes

  • Involve redox reactions where:

    • Energy from a spontaneous reaction is converted to electricity.

    • Electrical energy induces a non-spontaneous reaction.

Redox Reaction Fundamentals

  • Simultaneous Processes: Reduction and oxidation always occur together.

  • Change in Oxidation States: At least two substances undergo a change in oxidation state; one is oxidized and the other is reduced.

Oxidation and Reduction Defined

Oxidation

  • Originally described as reactions with O2 addition.

  • Loss of Electrons: Increases oxidation state.

  • Example: S + O2 → SO2

Reduction

  • Initially defined as reactions involving O2 removal.

  • Gain of Electrons: Decreases oxidation state.

  • Example: 2Fe2O3 + 3C → 4Fe + CO2

Mnemonics for Redox Reactions

  • Common Mnemonics:

    • “oil rig”: Oxidation Is Loss, Reduction Is Gain of electrons.

    • “leo ger”: Loss of Electron is Oxidation, Gain of Electron is Reduction.

Redox Reaction Basis

Characteristics

  • Oxidation Process:

    • Oxygen: Gain of O2.

    • Loss of Electrons: Decrease in electrons.

    • Increase in oxidation state.

  • Reduction Process:

    • Loss of O2.

    • Gain Electrons.

    • Decrease in oxidation state.

Applications of Redox Reactions

  • Photosynthesis: Reaction where light energy converts CO2 and H2O into glucose and O2.

  • Cellular Respiration: Breakdown of glucose to harvest energy (ATP).

  • Combustion of Fuels: Oxidation of substances producing heat and light.

  • Corrosion Processes: Material degradation from environmental reactions (e.g., rusting of iron).

  • Bleaching: Chemical processes to remove color from textiles.

  • Batteries: Dry cell batteries convert chemical energy to electrical energy.

Photosynthesis Chemical Reaction

  • Equation:6CO2 + 6H2O → C6H12O6 + 6O2 + Light Energy

  • Significance: Plants convert glucose to energy, proteins, fats/lipids, starch, or cellulose.

Cellular Respiration Chemical Reaction

  • Equation:C6H12O6 + 6O2 → 6CO2 + 6H2O + Energy

  • Energy Yield: Produces 34-36 ATP for cellular metabolism.

Combustion Reaction Explained

  • Definition: A reaction with O2 that releases heat and light.

  • Example of Complete Combustion:C3H8 + 5O2 → 3CO2 + 4H2O + Heat

Corrosion Reactions

  • Definition: Degradation from environmental reactions.

  • Example: Rusting of Iron:4Fe + 3O2 → 2Fe2O3

Bleaching Agents

  • Definition: Used to remove color from materials.

  • Common Agents: Sodium hypochlorite (NaOCl) and Hydrogen peroxide (H2O2) act as oxidizing agents.

  • Reaction Example:OCl- + 2e- + HOH → Cl- + 2OH-

Dry-Cell Battery Construction

  • Components: Metal electrode, graphite rod, moist electrolyte paste, and metal cylinder.

Oxidizing and Reducing Agents

Oxidizing Agent

  • Function: Accepts electrons from another substance, causing oxidation.

  • Characteristics: Decrease in oxidation number, strong oxidizers have high electronegativity (e.g., F, O, Cl).

Reducing Agent

  • Function: Provides electrons to reduce another substance.

  • Characteristics: Increases oxidation number; often consists of active metals that easily lose electrons.

Identifying Redox Reactions

  • If there is a change in oxidation states, a redox reaction occurred.

  • Reducing Agent: Element whose oxidation number increases.

  • Oxidizing Agent: Element whose oxidation number decreases.

Common Oxidizing and Reducing Agents

  • Oxidizing Agents: O2, HNO3, KMnO4, Cr2O7^2-, F2, Cl2.

  • Reducing Agents: H2, metals, hydrocarbons.

Activity Series of Metals

  • Metals' reactivity indicates the likelihood of oxidation.

  • Order of Reactivity:

    1. Lithium (Li)

    2. Barium (Ba)

    3. Calcium (Ca)

    4. Aluminum (Al)

    5. Zinc (Zn)

    6. Iron (Fe)

    7. Copper (Cu)

    • Conclusion: The more active metal is more readily oxidized.

Balancing Redox Reactions

Principles

  • Ensure gain of electrons equals loss of electrons.

  • Method 1: Half reaction method.

  • Method 2: Oxidation number method.

Half Reaction Method

  • Split overall reaction into oxidation and reduction half-reactions; balance separately and combine.

Systematic Balancing Methods

  • Oxidation Number Method: Focus on changes in oxidation numbers.

  • Half-Reaction Method: Inspect movements of electron flow for balancing.

Balancing Redox Equations in Acidic Solutions

  • Steps:

    1. Determine oxidation numbers.

    2. Identify and separate “redox” atoms.

    3. Balance for redox atoms and charges with electrons.

    4. Balance O with water and H using acids/protons.

Complete Example: Iron and Dichromate Reaction

  • Balanced equation:6Fe2+(aq) + Cr2O7^2-(aq) + 14H+(aq) → 6Fe3+(aq) + 2Cr3+(aq) + 7H2O(l)

Balancing Redox Equations in Basic Solutions

  • Repeat similar steps as in acidic solutions.

  • Utilize hydroxide ions to balance oxygen.

Example Reaction in Basic Solutions

  • Example: MnO2 + ClO3- + OH- → MnO4- + Cl- + H2O

    • Complete balanced equation:2MnO2 + ClO3- + 2OH- → 2MnO4- + Cl- + H2O

Final Remarks on Balancing

  • Techniques include oxidation-number-change and half-reaction methods for systematic balancing of redox equations.

  • Assess and ensure all species charges and atoms balanced after adjustments.