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Chemistry 2e

17.1 Review of Redox Chemistry

Chemistry 2e17.1 Review of Redox Chemistry
  1. Preface
  2. 1 Essential Ideas
    1. Introduction
    2. 1.1 Chemistry in Context
    3. 1.2 Phases and Classification of Matter
    4. 1.3 Physical and Chemical Properties
    5. 1.4 Measurements
    6. 1.5 Measurement Uncertainty, Accuracy, and Precision
    7. 1.6 Mathematical Treatment of Measurement Results
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  3. 2 Atoms, Molecules, and Ions
    1. Introduction
    2. 2.1 Early Ideas in Atomic Theory
    3. 2.2 Evolution of Atomic Theory
    4. 2.3 Atomic Structure and Symbolism
    5. 2.4 Chemical Formulas
    6. 2.5 The Periodic Table
    7. 2.6 Molecular and Ionic Compounds
    8. 2.7 Chemical Nomenclature
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  4. 3 Composition of Substances and Solutions
    1. Introduction
    2. 3.1 Formula Mass and the Mole Concept
    3. 3.2 Determining Empirical and Molecular Formulas
    4. 3.3 Molarity
    5. 3.4 Other Units for Solution Concentrations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  5. 4 Stoichiometry of Chemical Reactions
    1. Introduction
    2. 4.1 Writing and Balancing Chemical Equations
    3. 4.2 Classifying Chemical Reactions
    4. 4.3 Reaction Stoichiometry
    5. 4.4 Reaction Yields
    6. 4.5 Quantitative Chemical Analysis
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  6. 5 Thermochemistry
    1. Introduction
    2. 5.1 Energy Basics
    3. 5.2 Calorimetry
    4. 5.3 Enthalpy
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  7. 6 Electronic Structure and Periodic Properties of Elements
    1. Introduction
    2. 6.1 Electromagnetic Energy
    3. 6.2 The Bohr Model
    4. 6.3 Development of Quantum Theory
    5. 6.4 Electronic Structure of Atoms (Electron Configurations)
    6. 6.5 Periodic Variations in Element Properties
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  8. 7 Chemical Bonding and Molecular Geometry
    1. Introduction
    2. 7.1 Ionic Bonding
    3. 7.2 Covalent Bonding
    4. 7.3 Lewis Symbols and Structures
    5. 7.4 Formal Charges and Resonance
    6. 7.5 Strengths of Ionic and Covalent Bonds
    7. 7.6 Molecular Structure and Polarity
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  9. 8 Advanced Theories of Covalent Bonding
    1. Introduction
    2. 8.1 Valence Bond Theory
    3. 8.2 Hybrid Atomic Orbitals
    4. 8.3 Multiple Bonds
    5. 8.4 Molecular Orbital Theory
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  10. 9 Gases
    1. Introduction
    2. 9.1 Gas Pressure
    3. 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
    4. 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
    5. 9.4 Effusion and Diffusion of Gases
    6. 9.5 The Kinetic-Molecular Theory
    7. 9.6 Non-Ideal Gas Behavior
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  11. 10 Liquids and Solids
    1. Introduction
    2. 10.1 Intermolecular Forces
    3. 10.2 Properties of Liquids
    4. 10.3 Phase Transitions
    5. 10.4 Phase Diagrams
    6. 10.5 The Solid State of Matter
    7. 10.6 Lattice Structures in Crystalline Solids
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  12. 11 Solutions and Colloids
    1. Introduction
    2. 11.1 The Dissolution Process
    3. 11.2 Electrolytes
    4. 11.3 Solubility
    5. 11.4 Colligative Properties
    6. 11.5 Colloids
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  13. 12 Kinetics
    1. Introduction
    2. 12.1 Chemical Reaction Rates
    3. 12.2 Factors Affecting Reaction Rates
    4. 12.3 Rate Laws
    5. 12.4 Integrated Rate Laws
    6. 12.5 Collision Theory
    7. 12.6 Reaction Mechanisms
    8. 12.7 Catalysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  14. 13 Fundamental Equilibrium Concepts
    1. Introduction
    2. 13.1 Chemical Equilibria
    3. 13.2 Equilibrium Constants
    4. 13.3 Shifting Equilibria: Le Châtelier’s Principle
    5. 13.4 Equilibrium Calculations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  15. 14 Acid-Base Equilibria
    1. Introduction
    2. 14.1 Brønsted-Lowry Acids and Bases
    3. 14.2 pH and pOH
    4. 14.3 Relative Strengths of Acids and Bases
    5. 14.4 Hydrolysis of Salts
    6. 14.5 Polyprotic Acids
    7. 14.6 Buffers
    8. 14.7 Acid-Base Titrations
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  16. 15 Equilibria of Other Reaction Classes
    1. Introduction
    2. 15.1 Precipitation and Dissolution
    3. 15.2 Lewis Acids and Bases
    4. 15.3 Coupled Equilibria
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  17. 16 Thermodynamics
    1. Introduction
    2. 16.1 Spontaneity
    3. 16.2 Entropy
    4. 16.3 The Second and Third Laws of Thermodynamics
    5. 16.4 Free Energy
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  18. 17 Electrochemistry
    1. Introduction
    2. 17.1 Review of Redox Chemistry
    3. 17.2 Galvanic Cells
    4. 17.3 Electrode and Cell Potentials
    5. 17.4 Potential, Free Energy, and Equilibrium
    6. 17.5 Batteries and Fuel Cells
    7. 17.6 Corrosion
    8. 17.7 Electrolysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  19. 18 Representative Metals, Metalloids, and Nonmetals
    1. Introduction
    2. 18.1 Periodicity
    3. 18.2 Occurrence and Preparation of the Representative Metals
    4. 18.3 Structure and General Properties of the Metalloids
    5. 18.4 Structure and General Properties of the Nonmetals
    6. 18.5 Occurrence, Preparation, and Compounds of Hydrogen
    7. 18.6 Occurrence, Preparation, and Properties of Carbonates
    8. 18.7 Occurrence, Preparation, and Properties of Nitrogen
    9. 18.8 Occurrence, Preparation, and Properties of Phosphorus
    10. 18.9 Occurrence, Preparation, and Compounds of Oxygen
    11. 18.10 Occurrence, Preparation, and Properties of Sulfur
    12. 18.11 Occurrence, Preparation, and Properties of Halogens
    13. 18.12 Occurrence, Preparation, and Properties of the Noble Gases
    14. Key Terms
    15. Summary
    16. Exercises
  20. 19 Transition Metals and Coordination Chemistry
    1. Introduction
    2. 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
    3. 19.2 Coordination Chemistry of Transition Metals
    4. 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
    5. Key Terms
    6. Summary
    7. Exercises
  21. 20 Organic Chemistry
    1. Introduction
    2. 20.1 Hydrocarbons
    3. 20.2 Alcohols and Ethers
    4. 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
    5. 20.4 Amines and Amides
    6. Key Terms
    7. Summary
    8. Exercises
  22. 21 Nuclear Chemistry
    1. Introduction
    2. 21.1 Nuclear Structure and Stability
    3. 21.2 Nuclear Equations
    4. 21.3 Radioactive Decay
    5. 21.4 Transmutation and Nuclear Energy
    6. 21.5 Uses of Radioisotopes
    7. 21.6 Biological Effects of Radiation
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  23. A | The Periodic Table
  24. B | Essential Mathematics
  25. C | Units and Conversion Factors
  26. D | Fundamental Physical Constants
  27. E | Water Properties
  28. F | Composition of Commercial Acids and Bases
  29. G | Standard Thermodynamic Properties for Selected Substances
  30. H | Ionization Constants of Weak Acids
  31. I | Ionization Constants of Weak Bases
  32. J | Solubility Products
  33. K | Formation Constants for Complex Ions
  34. L | Standard Electrode (Half-Cell) Potentials
  35. M | Half-Lives for Several Radioactive Isotopes
  36. Answer Key
    1. Chapter 1
    2. Chapter 2
    3. Chapter 3
    4. Chapter 4
    5. Chapter 5
    6. Chapter 6
    7. Chapter 7
    8. Chapter 8
    9. Chapter 9
    10. Chapter 10
    11. Chapter 11
    12. Chapter 12
    13. Chapter 13
    14. Chapter 14
    15. Chapter 15
    16. Chapter 16
    17. Chapter 17
    18. Chapter 18
    19. Chapter 19
    20. Chapter 20
    21. Chapter 21
  37. Index
By the end of this section, you will be able to:
  • Describe defining traits of redox chemistry
  • Identify the oxidant and reductant of a redox reaction
  • Balance chemical equations for redox reactions using the half-reaction method

Since reactions involving electron transfer are essential to the topic of electrochemistry, a brief review of redox chemistry is provided here that summarizes and extends the content of an earlier text chapter (see chapter on reaction stoichiometry). Readers wishing additional review are referred to the text chapter on reaction stoichiometry.

Oxidation Numbers

By definition, a redox reaction is one that entails changes in oxidation number (or oxidation state) for one or more of the elements involved. The oxidation number of an element in a compound is essentially an assessment of how the electronic environment of its atoms is different in comparison to atoms of the pure element. By this description, the oxidation number of an atom in an element is equal to zero. For an atom in a compound, the oxidation number is equal to the charge the atom would have in the compound if the compound were ionic. Consequential to these rules, the sum of oxidation numbers for all atoms in a molecule is equal to the charge on the molecule. To illustrate this formalism, examples from the two compound classes, ionic and covalent, will be considered.

Simple ionic compounds present the simplest examples to illustrate this formalism, since by definition the elements’ oxidation numbers are numerically equivalent to ionic charges. Sodium chloride, NaCl, is comprised of Na+ cations and Cl anions, and so oxidation numbers for sodium and chlorine are, +1 and −1, respectively. Calcium fluoride, CaF2, is comprised of Ca2+ cations and F anions, and so oxidation numbers for calcium and fluorine are, +2 and −1, respectively.

Covalent compounds require a more challenging use of the formalism. Water is a covalent compound whose molecules consist of two H atoms bonded separately to a central O atom via polar covalent O−H bonds. The shared electrons comprising an O−H bond are more strongly attracted to the more electronegative O atom, and so it acquires a partial negative charge in the water molecule (relative to an O atom in elemental oxygen). Consequently, H atoms in a water molecule exhibit partial positive charges compared to H atoms in elemental hydrogen. The sum of the partial negative and partial positive charges for each water molecule is zero, and the water molecule is neutral.

Imagine that the polarization of shared electrons within the O−H bonds of water were 100% complete—the result would be transfer of electrons from H to O, and water would be an ionic compound comprised of O2− anions and H+ cations. And so, the oxidations numbers for oxygen and hydrogen in water are −2 and +1, respectively. Applying this same logic to carbon tetrachloride, CCl4, yields oxidation numbers of +4 for carbon and −1 for chlorine. In the nitrate ion, NO3NO3, the oxidation number for nitrogen is +5 and that for oxygen is −2, summing to equal the 1− charge on the molecule:

(1Natom)(+5Natom)+(3Oatoms)(−2Oatom)=+5+−6=−1(1Natom)(+5Natom)+(3Oatoms)(−2Oatom)=+5+−6=−1

Balancing Redox Equations

The unbalanced equation below describes the decomposition of molten sodium chloride:

NaCl(l)Na(l)+Cl2(g)unbalancedNaCl(l)Na(l)+Cl2(g)unbalanced

This reaction satisfies the criterion for redox classification, since the oxidation number for Na is decreased from +1 to 0 (it undergoes reduction) and that for Cl is increased from −1 to 0 (it undergoes oxidation). The equation in this case is easily balanced by inspection, requiring stoichiometric coefficients of 2 for the NaCl and Na:

2NaCl(l)2Na(l)+Cl2(g)balanced2NaCl(l)2Na(l)+Cl2(g)balanced

Redox reactions that take place in aqueous solutions are commonly encountered in electrochemistry, and many involve water or its characteristic ions, H+(aq) and OH(aq), as reactants or products. In these cases, equations representing the redox reaction can be very challenging to balance by inspection, and the use of a systematic approach called the half-reaction method is helpful. This approach involves the following steps:

  1. Write skeletal equations for the oxidation and reduction half-reactions.
  2. Balance each half-reaction for all elements except H and O.
  3. Balance each half-reaction for O by adding H2O.
  4. Balance each half-reaction for H by adding H+.
  5. Balance each half-reaction for charge by adding electrons.
  6. If necessary, multiply one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.
  7. Add the two half-reactions and simplify.
  8. If the reaction takes place in a basic medium, add OH ions the equation obtained in step 7 to neutralize the H+ ions (add in equal numbers to both sides of the equation) and simplify.

The examples below demonstrate the application of this method to balancing equations for aqueous redox reactions.

Example 17.1

Balancing Equations for Redox Reactions in Acidic Solutions Write the balanced equation representing reaction between solid copper and nitric acid to yield aqueous copper(II) ions and nitrogen monoxide gas.

Solution Following the steps of the half-reaction method:

  1. Write skeletal equations for the oxidation and reduction half-reactions.
    oxidation:Cu(s)Cu2+(aq)oxidation:Cu(s)Cu2+(aq)
    reduction:HNO3(aq)NO(g)reduction:HNO3(aq)NO(g)
  2. Balance each half-reaction for all elements except H and O.
    oxidation:Cu(s)Cu2+(aq)oxidation:Cu(s)Cu2+(aq)
    reduction:HNO3(aq)NO(g)reduction:HNO3(aq)NO(g)
  3. Balance each half-reaction for O by adding H2O.
    oxidation:Cu(s)Cu2+(aq)oxidation:Cu(s)Cu2+(aq)
    reduction:HNO3(aq)NO(g)+2H2O(l)reduction:HNO3(aq)NO(g)+2H2O(l)
  4. Balance each half-reaction for H by adding H+.
    oxidation:Cu(s)Cu2+(aq)oxidation:Cu(s)Cu2+(aq)
    reduction:3H+(aq)+HNO3(aq)NO(g)+2H2O(l)reduction:3H+(aq)+HNO3(aq)NO(g)+2H2O(l)
  5. Balance each half-reaction for charge by adding electrons.
    oxidation:Cu(s)Cu2+(aq)+2eoxidation:Cu(s)Cu2+(aq)+2e
    reduction:3e+3H+(aq)+HNO3(aq)NO(g)+2H2O(l)reduction:3e+3H+(aq)+HNO3(aq)NO(g)+2H2O(l)
  6. If necessary, multiply one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.
    oxidation (×3):3Cu(s)3Cu2+(aq)+62eoxidation (×3):3Cu(s)3Cu2+(aq)+62e
    reduction (×2):63e+63H+(aq)+2HNO3(aq)2NO(g)+42H2O(l)reduction (×2):63e+63H+(aq)+2HNO3(aq)2NO(g)+42H2O(l)
  7. Add the two half-reactions and simplify.
    3Cu(s)+6e+6H+(aq)+2HNO3(aq)3Cu2+(aq)+6e+2NO(g)4H2O(l)3Cu(s)+6e+6H+(aq)+2HNO3(aq)3Cu2+(aq)+6e+2NO(g)4H2O(l)
    3Cu(s)+6H+(aq)+2HNO3(aq)3Cu2+(aq)+2NO(g)+4H2O(l)3Cu(s)+6H+(aq)+2HNO3(aq)3Cu2+(aq)+2NO(g)+4H2O(l)
  8. If the reaction takes place in a basic medium, add OH ions the equation obtained in step 7 to neutralize the H+ ions (add in equal numbers to both sides of the equation) and simplify.
    This step not necessary since the solution is stipulated to be acidic.

The balanced equation for the reaction in an acidic solution is then

3Cu(s)+6H+(aq)+2HNO3(aq)3Cu2+(aq)+2NO(g)+4H2O(l)3Cu(s)+6H+(aq)+2HNO3(aq)3Cu2+(aq)+2NO(g)+4H2O(l)

Check Your Learning The reaction above results when using relatively diluted nitric acid. If concentrated nitric acid is used, nitrogen dioxide is produced instead of nitrogen monoxide. Write a balanced equation for this reaction.

Answer:

Cu(s)+2H+(aq)+2HNO3(aq)Cu2+(aq)+2NO2(g)+2H2O(l)Cu(s)+2H+(aq)+2HNO3(aq)Cu2+(aq)+2NO2(g)+2H2O(l)

Example 17.2

Balancing Equations for Redox Reactions in Basic Solutions Write the balanced equation representing reaction between aqueous permanganate ion, MnO4MnO4, and solid chromium(III) hydroxide, Cr(OH)3, to yield solid manganese(IV) oxide, MnO2, and aqueous chromate ion, CrO42−CrO42− The reaction takes place in a basic solution.

Solution Following the steps of the half-reaction method:

  1. Write skeletal equations for the oxidation and reduction half-reactions.
    oxidation:Cr(OH)3(s)CrO42−(aq)oxidation:Cr(OH)3(s)CrO42−(aq)
    reduction:MnO4(aq)MnO2(s)reduction:MnO4(aq)MnO2(s)
  2. Balance each half-reaction for all elements except H and O.
    oxidation:Cr(OH)3(s)CrO42−(aq)oxidation:Cr(OH)3(s)CrO42−(aq)
    reduction:MnO4(aq)MnO2(s)reduction:MnO4(aq)MnO2(s)
  3. Balance each half-reaction for O by adding H2O.
    oxidation:H2O(l)+Cr(OH)3(s)CrO42−(aq)oxidation:H2O(l)+Cr(OH)3(s)CrO42−(aq)
    reduction:MnO4(aq)MnO2(s)+2H2O(l)reduction:MnO4(aq)MnO2(s)+2H2O(l)
  4. Balance each half-reaction for H by adding H+.
    oxidation:H2O(l)+Cr(OH)3(s)CrO42−(aq)+5H+(aq)oxidation:H2O(l)+Cr(OH)3(s)CrO42−(aq)+5H+(aq)
    reduction:4H+(aq)+MnO4(aq)MnO2(s)+2H2O(l)reduction:4H+(aq)+MnO4(aq)MnO2(s)+2H2O(l)
  5. Balance each half-reaction for charge by adding electrons.
    oxidation:H2O(l)+Cr(OH)3(s)CrO42−(aq)+5H+(aq)+3eoxidation:H2O(l)+Cr(OH)3(s)CrO42−(aq)+5H+(aq)+3e
    reduction:3e+4H+(aq)+MnO4(aq)MnO2(s)+2H2O(l)reduction:3e+4H+(aq)+MnO4(aq)MnO2(s)+2H2O(l)
  6. If necessary, multiply one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.
    This step is not necessary since the number of electrons is already in balance.
  7. Add the two half-reactions and simplify.
    H2O(l)+Cr(OH)3(s)+3e+4H+(aq)+MnO4(aq)CrO42−(aq)+5H+(aq) +3e+MnO2(s)+2H2O(l)H2O(l)+Cr(OH)3(s)+3e+4H+(aq)+MnO4(aq)CrO42−(aq)+5H+(aq) +3e+MnO2(s)+2H2O(l)
    Cr(OH)3(s)+MnO4(aq)CrO42−(aq)H+(aq)+MnO2(s)+H2O(l)Cr(OH)3(s)+MnO4(aq)CrO42−(aq)H+(aq)+MnO2(s)+H2O(l)
  8. If the reaction takes place in a basic medium, add OH ions the equation obtained in step 7 to neutralize the H+ ions (add in equal numbers to both sides of the equation) and simplify.
    OH(aq)+Cr(OH)3(s)+MnO4(aq)CrO42−(aq)+H+(aq)+OH(aq)+MnO2(s)+H2O(l)OH(aq)+Cr(OH)3(s)+MnO4(aq)CrO42−(aq)+H+(aq)+OH(aq)+MnO2(s)+H2O(l)
    OH(aq)+Cr(OH)3(s)+MnO4(aq)CrO42(aq)+MnO2(s)+2H2O(l)OH(aq)+Cr(OH)3(s)+MnO4(aq)CrO42(aq)+MnO2(s)+2H2O(l)

Check Your Learning Aqueous permanganate ion may also be reduced using aqueous bromide ion, Br, the products of this reaction being solid manganese(IV) oxide and aqueous bromate ion, BrO3. Write the balanced equation for this reaction occurring in a basic medium.

Answer:

H2O(l)+2MnO4(aq)+Br(aq)2MnO2(s)+BrO3(aq)+2OH(aq)H2O(l)+2MnO4(aq)+Br(aq)2MnO2(s)+BrO3(aq)+2OH(aq)

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