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Introductory Chemistry - 1st Canadian Edition: Oxidation-Reduction Reactions

Introductory Chemistry - 1st Canadian Edition
Oxidation-Reduction Reactions
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table of contents
  1. Cover
  2. Title Page
  3. Copyright
  4. Table Of Contents
  5. Acknowledgments
  6. Dedication
  7. About BCcampus Open Education
  8. Chapter 1. What is Chemistry
    1. Some Basic Definitions
    2. Chemistry as a Science
  9. Chapter 2. Measurements
    1. Expressing Numbers
    2. Significant Figures
    3. Converting Units
    4. Other Units: Temperature and Density
    5. Expressing Units
    6. End-of-Chapter Material
  10. Chapter 3. Atoms, Molecules, and Ions
    1. Acids
    2. Ions and Ionic Compounds
    3. Masses of Atoms and Molecules
    4. Molecules and Chemical Nomenclature
    5. Atomic Theory
    6. End-of-Chapter Material
  11. Chapter 4. Chemical Reactions and Equations
    1. The Chemical Equation
    2. Types of Chemical Reactions: Single- and Double-Displacement Reactions
    3. Ionic Equations: A Closer Look
    4. Composition, Decomposition, and Combustion Reactions
    5. Oxidation-Reduction Reactions
    6. Neutralization Reactions
    7. End-of-Chapter Material
  12. Chapter 5. Stoichiometry and the Mole
    1. Stoichiometry
    2. The Mole
    3. Mole-Mass and Mass-Mass Calculations
    4. Limiting Reagents
    5. The Mole in Chemical Reactions
    6. Yields
    7. End-of-Chapter Material
  13. Chapter 6. Gases
    1. Pressure
    2. Gas Laws
    3. Other Gas Laws
    4. The Ideal Gas Law and Some Applications
    5. Gas Mixtures
    6. Kinetic Molecular Theory of Gases
    7. Molecular Effusion and Diffusion
    8. Real Gases
    9. End-of-Chapter Material
  14. Chapter 7. Energy and Chemistry
    1. Formation Reactions
    2. Energy
    3. Stoichiometry Calculations Using Enthalpy
    4. Enthalpy and Chemical Reactions
    5. Work and Heat
    6. Hess’s Law
    7. End-of-Chapter Material
  15. Chapter 8. Electronic Structure
    1. Light
    2. Quantum Numbers for Electrons
    3. Organization of Electrons in Atoms
    4. Electronic Structure and the Periodic Table
    5. Periodic Trends
    6. End-of-Chapter Material
  16. Chapter 9. Chemical Bonds
    1. Lewis Electron Dot Diagrams
    2. Electron Transfer: Ionic Bonds
    3. Covalent Bonds
    4. Other Aspects of Covalent Bonds
    5. Violations of the Octet Rule
    6. Molecular Shapes and Polarity
    7. Valence Bond Theory and Hybrid Orbitals
    8. Molecular Orbitals
    9. End-of-Chapter Material
  17. Chapter 10. Solids and Liquids
    1. Properties of Liquids
    2. Solids
    3. Phase Transitions: Melting, Boiling, and Subliming
    4. Intermolecular Forces
    5. End-of-Chapter Material
  18. Chapter 11. Solutions
    1. Colligative Properties of Solutions
    2. Concentrations as Conversion Factors
    3. Quantitative Units of Concentration
    4. Colligative Properties of Ionic Solutes
    5. Some Definitions
    6. Dilutions and Concentrations
    7. End-of-Chapter Material
  19. Chapter 12. Acids and Bases
    1. Acid-Base Titrations
    2. Strong and Weak Acids and Bases and Their Salts
    3. Brønsted-Lowry Acids and Bases
    4. Arrhenius Acids and Bases
    5. Autoionization of Water
    6. Buffers
    7. The pH Scale
    8. End-of-Chapter Material
  20. Chapter 13. Chemical Equilibrium
    1. Chemical Equilibrium
    2. The Equilibrium Constant
    3. Shifting Equilibria: Le Chatelier’s Principle
    4. Calculating Equilibrium Constant Values
    5. Some Special Types of Equilibria
    6. End-of-Chapter Material
  21. Chapter 14. Oxidation and Reduction
    1. Oxidation-Reduction Reactions
    2. Balancing Redox Reactions
    3. Applications of Redox Reactions: Voltaic Cells
    4. Electrolysis
    5. End-of-Chapter Material
  22. Chapter 15. Nuclear Chemistry
    1. Units of Radioactivity
    2. Uses of Radioactive Isotopes
    3. Half-Life
    4. Radioactivity
    5. Nuclear Energy
    6. End-of-Chapter Material
  23. Chapter 16. Organic Chemistry
    1. Hydrocarbons
    2. Branched Hydrocarbons
    3. Alkyl Halides and Alcohols
    4. Other Oxygen-Containing Functional Groups
    5. Other Functional Groups
    6. Polymers
    7. End-of-Chapter Material
  24. Chapter 17. Kinetics
    1. Factors that Affect the Rate of Reactions
    2. Reaction Rates
    3. Rate Laws
    4. Concentration–Time Relationships: Integrated Rate Laws
    5. Activation Energy and the Arrhenius Equation
    6. Reaction Mechanisms
    7. Catalysis
    8. End-of-Chapter Material
  25. Chapter 18. Chemical Thermodynamics
    1. Spontaneous Change
    2. Entropy and the Second Law of Thermodynamics
    3. Measuring Entropy and Entropy Changes
    4. Gibbs Free Energy
    5. Spontaneity: Free Energy and Temperature
    6. Free Energy under Nonstandard Conditions
    7. End-of-Chapter Material
  26. Appendix A: Periodic Table of the Elements
  27. Appendix B: Selected Acid Dissociation Constants at 25°C
  28. Appendix C: Solubility Constants for Compounds at 25°C
  29. Appendix D: Standard Thermodynamic Quantities for Chemical Substances at 25°C
  30. Appendix E: Standard Reduction Potentials by Value
  31. Glossary
  32. About the Authors
  33. Versioning History

Oxidation-Reduction Reactions

Learning Objectives

  1. Define oxidation and reduction.
  2. Assign oxidation numbers to atoms in simple compounds.
  3. Recognize a reaction as an oxidation-reduction reaction.

Consider this chemical reaction:

\ce{Mg(s)}+\ce{Cl2(g)}\rightarrow \ce{MgCl2}

The reactants are two electrically neutral elements; they have the same number of electrons as protons. The product, however, is ionic; it is composed of Mg2+ and Cl− ions. Somehow, the individual Mg atoms lose two electrons to make the Mg2+ ion, while the Cl atoms gain an electron to become Cl− ions. This reaction involves the transfer of electrons between atoms.

The process of losing and gaining electrons occurs simultaneously. However, mentally we can separate the two processes. Oxidation is defined as the loss of one or more electrons by an atom. Reduction is defined as the gain of one or more electrons by an atom. So oxidation and reduction always occur together; it is only mentally that we can separate them. Chemical reactions that involve the transfer of electrons are called oxidation-reduction (or redox) reactions.

Redox reactions require that we keep track of the electrons assigned to each atom in a chemical reaction. How do we do that? We use oxidation numbers to keep track of electrons in atoms. Oxidation numbers are assigned to atoms based on four rules. Oxidation numbers are not necessarily equal to the charge on the atom (although sometimes they can be); we must keep the concepts of charge and oxidation numbers separate.

The rules for assigning oxidation numbers to atoms are as follows:

  1. Atoms in their elemental state are assigned an oxidation number of 0.
  2. Atoms in monatomic (i.e., one-atom) ions are assigned an oxidation number equal to their charge. Oxidation numbers are usually written with the sign first, then the magnitude, to differentiate them from charges.
  3. In compounds, fluorine is assigned a −1 oxidation number; oxygen is usually assigned a −2 oxidation number [except in peroxide compounds (where it is −1) and in binary compounds with fluorine (where it is positive)]; and hydrogen is usually assigned a +1 oxidation number [except when it exists as the hydride ion (H−), in which case rule 2 prevails].
  4. In compounds, all other atoms are assigned an oxidation number so that the sum of the oxidation numbers on all the atoms in the species equals the charge on the species (which is zero if the species is neutral).

Here are some examples for practice. In H2, both H atoms have an oxidation number of 0 by rule 1. In MgCl2, magnesium has an oxidation number of +2, while chlorine has an oxidation number of −1 by rule 2. In H2O, the H atoms each have an oxidation number of +1, while the O atom has an oxidation number of −2, even though hydrogen and oxygen do not exist as ions in this compound (rule 3). By contrast, by rule 3, each H atom in hydrogen peroxide (H2O2) has an oxidation number of +1, while each O atom has an oxidation number of −1. We can use rule 4 to determine oxidation numbers for the atoms in SO2. Each O atom has an oxidation number of −2; for the sum of the oxidation numbers to equal the charge on the species (which is zero), the S atom is assigned an oxidation number of +4. Does this mean that the sulfur atom has a 4+ charge on it? No, it means only that the S atom is assigned a +4 oxidation number by our rules of apportioning electrons among the atoms in a compound.

Example 14.1

Assign oxidation numbers to the atoms in each substance.

  1. Cl2
  2. GeO2
  3. Ca(NO3)2

Solution

  1. Cl2 is the elemental form of chlorine. Rule 1 states each atom has an oxidation number of 0.
  2. By rule 3, oxygen is normally assigned an oxidation number of −2. For the sum of the oxidation numbers to equal the charge on the species (zero), the Ge atom is assigned an oxidation number of +4.
  3. Ca(NO3)2 can be separated into two parts: the Ca2+ ion and the NO3− ion. Considering these separately, the Ca2+ ion has an oxidation number of +2 by rule 2. Now consider the NO3− ion. Oxygen is assigned an oxidation number of −2, and there are three of them. According to rule 4, the sum of the oxidation numbers on all atoms must equal the charge on the species, so we have the simple algebraic equation:

    x+3(-2)=-1

    where x is the oxidation number of the N atom and the −1 represents the charge on the species. Evaluating for x:

    \begin{array}{rrrrr} x&+&(-6)&=&-1 \\ &&x&=&5 \end{array}

    Thus the oxidation number on the N atom in the NO3− ion is +5.

Test Yourself
Assign oxidation numbers to the atoms in H3PO4.

Answer
H: +1; O: −2; P: +5

All redox reactions occur with a simultaneous change in the oxidation numbers of some atoms. At least two elements must change their oxidation numbers. When an oxidation number of an atom is increased in the course of a redox reaction, that atom is being oxidized. When an oxidation number of an atom is decreased in the course of a redox reaction, that atom is being reduced. Thus oxidation and reduction can also be defined in terms of increasing or decreasing oxidation numbers, respectively.

Oxidation reactions can become quite complex, as attested by the following redox reaction:

\begin{array}{ccccccccccc} \ce{6H+(aq)}&+&\ce{2MnO4-(aq)}&+&\ce{5H2O2(\ell)}&\rightarrow&\ce{2Mn^{2+}(aq)}&+&\ce{5O2(g)}&+&\ce{8H2O(\ell)} \\ &&(+7)&&(-1)&&(+2)&&(0)&& \end{array}

To demonstrate that this is a redox reaction, the oxidation numbers of the species being oxidized and reduced are listed; can you determine what is being oxidized and what is being reduced? This is also an example of a net ionic reaction; spectator ions that do not change oxidation numbers are not displayed in the equation. Eventually, we will need to learn techniques for writing correct (i.e., balanced) redox reactions.

Example 14.2

Identify what is being oxidized and reduced in this redox reaction.

2Na + Br2 → 2NaBr

Solution
Both reactants are the elemental forms of their atoms, so the Na and Br atoms have oxidation numbers of 0. In the ionic product, the Na+ ions have an oxidation number of +1, while the Br− ions have an oxidation number of −1.

Sodium is increasing its oxidation number from 0 to +1, so it is being oxidized; bromine is decreasing its oxidation number from 0 to −1, so it is being reduced:

NaBr Redox

Because oxidation numbers are changing, this is a redox reaction. The total number of electrons being lost by sodium (two, one lost from each Na atom) is gained by bromine (two, one gained for each Br atom).

Test Yourself
Identify what is being oxidized and reduced in this redox reaction.

C + O2 → CO2

Answer
C is being oxidized from 0 to +4; O is being reduced from 0 to −2.

Food and Drink App: Fortifying Food with Iron

Iron is an essential mineral in our diet; iron-containing compounds like the heme protein in hemoglobin could not function without it. Most biological iron has the form of the Fe2+ ion; iron with other oxidation numbers is almost inconsequential in human biology (although the body does contain an enzyme to reduce Fe3+ to Fe2+, so Fe3+ must have some biological significance, albeit minor). To ensure that we ingest enough iron, many foods are enriched with iron. Although Fe2+ compounds are the most logical substances to use, some foods — bread and breakfast cereals are the most well-known examples — use “reduced iron” as an ingredient. Reduced iron is simply iron metal; iron is added as a fine metallic powder. The metallic iron is oxidized to Fe2+ in the digestive system and then absorbed by the body, but the question remains: Why are we ingesting metallic iron? Why not just use Fe2+ salts as an additive?

Ingredients
Many prepared foods list reduced iron in their ingredients list.

Although it is difficult to establish conclusive reasons, a search of scientific and medical literature suggests a few reasons. One reason is that fine iron filings do not affect the taste of the product. The size of the iron powder (several dozen micrometers) is not noticeable when chewing iron-supplemented foods, and the tongue does not detect any changes in flavour that can be detected when using Fe2+ salts. Fe2+ compounds can affect other properties of foodstuffs during preparation and cooking, like dough pliability, yeast growth, and colour. Finally, of the common iron substances that might be used, metallic iron is the least expensive. These factors appear to be among the reasons why metallic iron is the supplement of choice in some foods.

Key Takeaways

  • Oxidation-reduction (redox) reactions involve the transfer of electrons from one atom to another.
  • Oxidation numbers are used to keep track of electrons in atoms.
  • There are rules for assigning oxidation numbers to atoms.
  • Oxidation is an increase in oxidation number (loss of electrons); reduction is a decrease in oxidation number (gain of electrons).

Exercises

Questions

  1. Is this reaction a redox reaction? Explain your answer.

    2K(s) + Br2(ℓ) → 2KBr(s)

  2. Is this reaction a redox reaction? Explain your answer.

    2NaCl(aq) + Pb(NO3)2(aq) → 2NaNO3(aq) + PbCl2(s)

  3. Which substance loses electrons and which substance gains electrons in this reaction?

    2Mg(s) + O2(g) → 2MgO

  4. Which substance loses electrons and which substance gains electrons in this reaction?

    16Fe(s) + 3S8(s) → 8Fe2S3(s)

  5. Which substance is oxidized and which substance is reduced in this reaction?

    2Li(s) + O2(g) → Li2O2(s)

  6. Which substance is oxidized and which substance is reduced in this reaction?

    2Fe(s) + 3I2(s) → 2FeI3(s)

  7. What are two different definitions of oxidation?
  8. What are two different definitions of reduction?
  9. Assign oxidation numbers to the atoms in each substance.
    1. P4
    2. SO3
    3. SO32−
    4. Ca3(PO3)2
  10. Assign oxidation numbers to the atoms in each substance.
    1. PCl5
    2. (NH4)2Se
    3. Ag
    4. Li2O2
  11. Assign oxidation numbers to the atoms in each substance.
    1. NO
    2. NO2
    3. CrCl2
    4. CrCl3
  12. Assign oxidation numbers to the atoms in each substance.
    1. NaH
    2. N2O3
    3. NO2−
    4. CuNO3
  13. Assign oxidation numbers to the atoms in each substance.
    1. CH2O
    2. NH3
    3. Rb2SO4
    4. Zn(C2H3O2)2
  14. Assign oxidation numbers to the atoms in each substance.
    1. C6H6
    2. B(OH)3
    3. Li2S
    4. Au
  15. Identify what is being oxidized and reduced in this redox reaction by assigning oxidation numbers to the atoms.

    2NO + Cl2 → 2NOCl

  16. Identify what is being oxidized and reduced in this redox reaction by assigning oxidation numbers to the atoms.

    Sr + SO3 → SrSO3

  17. Identify what is being oxidized and reduced in this redox reaction by assigning oxidation numbers to the atoms.

    2KrF2 + 2H2O → 2Kr + 4HF + O2

  18. Identify what is being oxidized and reduced in this redox reaction by assigning oxidation numbers to the atoms.

    SO3 + SCl2 → SOCl2 + SO2

  19. Identify what is being oxidized and reduced in this redox reaction by assigning oxidation numbers to the atoms.

    2Rb + MgCl2 → 2RbCl + Mg

  20. Identify what is being oxidized and reduced in this redox reaction by assigning oxidation numbers to the atoms.

    2C8H18 + 25O2 → 16CO2 + 18H2O

Answers

  1. Yes, because oxidation numbers are changing.
  1. lose: Mg; gain: O
  1. oxidized: Li; reduced: O
  1. increase in oxidation number; loss of electrons
    1. P: 0
    2. S: +6; O: −2
    3. S: +4; O: −2
    4. Ca: +2; P: +3; O: −2
    1. N: +2; O: −2
    2. N: +4; O: −2
    3. Cr: +2; Cl: −1
    4. Cr: +3; Cl: −1
    1. C: 0; H: +1; O: −2
    2. N: −3; H: +1
    3. Rb: +1; S: +6; O: −2
    4. Zn: +2; C: 0; H: +1; O: −2
  1. oxidized: N; reduced: Cl
  1. oxidized: O; reduced: Kr
  1. oxidized: Rb; reduced: Mg

Annotate

Next Chapter
Balancing Redox Reactions
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Chemistry

Copyright © 2014

                                by Jessie A. Key

            Introductory Chemistry - 1st Canadian Edition by Jessie A. Key is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.
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