Oxidation-Reduction Reactions Study Notes

CHAPTER 7: OXIDATION-REDUCTION REACTIONS (REDOX REACTIONS)

Overview of Redox Reactions

• Reducing Agent: A substance that loses electrons, causing the reduction of another species.

• Oxidizing Agent: A substance that gains electrons, causing the oxidation of another species.

• Electrons (e-): The transferred particles in oxidation-reduction reactions.

Definitions of Key Terms

• Oxidation: Defined as the loss of electrons, which corresponds to an increase in oxidation state.

• Reduction: Defined as the gain of electrons, which corresponds to a decrease in oxidation state.

• Key Principle: There cannot be reduction without oxidation; the number of electrons gained by one species equals the number lost by another species.

Example Reactions

1. Oxidation Reaction:
   $ext{Zn}(s) <br>ightarrow ext{Zn}^{2+}(aq) + 2e^-$

   • Zinc atom loses 2 valence electrons and is oxidized.

2. Reduction Reaction:
   $ext{Cu}^{2+}(aq) + 2e^- <br>ightarrow ext{Cu}(s)$

   • Copper ion gains 2 electrons and is reduced.

Overall Redox Reaction

• Comprehensive Redox Reaction: $ext{Zn}(s) + ext{Cu}^{2+}(aq) ightarrow ext{Zn}^{2+}(aq) + ext{Cu}(s)$

  • Observation: Electrons are not included in the overall reaction.

  • Oxidation State Changes:

    • Zinc: 0 to +2 (loss of 2 electrons; Zn(s) is oxidized).

    • Copper: +2 to 0 (gain of 2 electrons; Cu2+ is reduced).

  • Role Designation:

    • Reducing Agent: Zinc (causes reduction of Cu2+).

    • Oxidizing Agent: Copper (causes oxidation of Zn).

  • Agent Changes:

    • The reducing agent itself gets oxidized.

    • The oxidizing agent itself gets reduced.

Identifying Redox Reactions via Oxidation States

Guidelines for Determining Oxidation States

1. In neutral compounds, the sum of all oxidation states must equal zero.

2. In ions, the sum must equal the charge on the ion.

3. Free elements have an oxidation state of zero (e.g., Na, H2, P4).

4. The oxidation state of Fluorine (F) in compounds is -1.

5. Oxidation states for metals:

   • Group 1A: +1

   • Group 2A: +2

6. Hydrogen's oxidation states depend on context:

   • With nonmetals: +1

   • With metals: -1

7. The oxidation state of Oxygen in compounds is -2 except in peroxides or with Fluorine.

8. Typical oxidation states for nonmetals:

   • Group 17: -1

   • Group 16: -2

   • Group 15: -3

Determining Oxidation States in a Reaction

Sample Problem with Reaction

• Given Reaction: $ext{Mg}(s) + 2 ext{H}<em>2 ext{O}(l) ightarrow ext{Mg(OH)}</em>2(aq) + ext{H}_2(g)$

  • Statements to evaluate:

    • Mg(s) as the oxidizing agent.

    • Mg(s) being oxidized.

    • H atoms in H2O undergoing reduction.

    • H2O acting as the oxidizing agent.

Types of Redox Reactions

Typical Redox Reaction Types

1. Single Replacement Reactions: An element replaces another in a compound.

   • Example:
     $ext{CuSO}<em>4 + ext{Zn} ightarrow ext{Cu} + ext{ZnSO}</em>4$

2. Decomposition Reactions: A compound breaks down into simpler substances.

   • Example:
     $2 ext{HgO}(s) <br>ightarrow 2 ext{Hg}(l) + ext{O}_2(g)$

3. Combination Reactions: Two or more substances combine to form a single substance.

   • Example:
     $2 ext{Na}(s) + ext{Cl}_2(g) <br>ightarrow 2 ext{NaCl}(s)$

4. Combustion Reactions: Substance reacts with oxygen, producing heat.

   • Example:
     $2 ext{H}<em>2(g) + ext{O}</em>2(g) <br>ightarrow 2 ext{H}_2 ext{O}(g)$

   • Example:
     $ext{C}<em>3 ext{H}</em>8(g) + 5 ext{O}<em>2(g) ightarrow 3 ext{CO}</em>2(g) + 4 ext{H}_2 ext{O}(g)$

5. Disproportionation Reactions: A single reactant is both oxidized and reduced.

   • Example:
     $3 ext{Cl}<em>2(g) + 6 ext{OH}^-(aq) ightarrow 5 ext{Cl}^-(aq) + ext{ClO}</em>3^-(aq) + 3 ext{H}_2 ext{O}(l)$

Non-redox Reactions

Typical Non-redox Reactions

1. Double Replacement Reactions: Exchange of cations or anions between two ionic compounds.

   • Example:
     $ext{HCl}(aq) + ext{NaOH}(aq) <br>ightarrow ext{NaCl}(aq) + ext{H}_2 ext{O}(l)$

2. Precipitation Reactions: Formation of a solid from the solution.

   • Example:
     $ext{CuCl}<em>2(aq) + 2 ext{AgNO}</em>3(aq) <br>ightarrow ext{Cu(NO}<em>3)</em>2(aq) + 2 ext{AgCl}(s)$

Oxidation States and Organic Compounds

• C-C Bonds: Do not alter the oxidation state (remains zero).

• C-H Bonds: Hydrogen has an oxidation state of +1. Each H attached decreases the oxidation state of C by 1.

• C with Electronegativity: Bonds to more electronegative elements (like O, N, Cl) increase oxidation state of C by 1.

• Electronegativity Trend: F > O > N, Cl > Br > C > H

Guidelines for Oxidation States in Organic Compounds

1. C in C-C bonds has an oxidation state of zero.

2. In C-H bonds, H is treated as having an oxidation state of +1 leading to a decrease in oxidation state of C.

3. Bonds to oxygen or more electronegative elements increase the oxidation state of C.

Identifying Redox Reactions in Organic Compounds

• Reduction of Organic Compounds:

  • Decrease in oxidation state for carbon.

  • Decrease in oxygen content.

  • Increase in hydrogen content.

• Oxidation of Organic Compounds:

  • Increase in oxidation state for carbon.

  • Increase in oxygen content.

  • Decrease in hydrogen content.

  • Broader Definition: Oxidation is any process that increases the content of more electronegative atoms than carbon.

Example Reaction with Methane

• Query: Is methane (CH4) being reduced or oxidized in the following reaction? $ext{CH}<em>4 + ext{Cl}</em>2 ightarrow ext{CH}_3 ext{Cl} + ext{HCl}$

  • Presence of UV light indicates a radical substitution, leading to oxidation of methane.

Identifying Redox Reactions

• To discern redox:

  • Evaluate reactions based on:

  1. $ext{HBr}(aq) + ext{KOH}(aq) ightarrow ext{KBr}(aq) + ext{H}_2 ext{O}(l)$

     • Not a redox reaction.

  2. $2 ext{Al}(s) + 3 ext{Cl}<em>2(g) ightarrow 2 ext{AlCl}</em>3(s)$

     • Is a redox reaction.

  3. $14 ext{H}^+(aq) + ext{Cr}<em>2 ext{O}</em>7^{2-}(aq) + 3 ext{Ni}(s) ightarrow 3 ext{Ni}^{2+}(aq) + 2 ext{Cr}^{3+}(aq) + 7 ext{H}_2 ext{O}(l)$

     • Is a redox reaction.

  4. $ext{AgNO}<em>3(aq) + ext{HCl}(aq) ightarrow ext{HNO}</em>3(aq) + ext{AgCl}(s)$

     • Not a redox reaction.

Identifying Reducing and Oxidizing Agents

• Example Reaction:
  $ext{Fe}<em>2 ext{S}</em>3 + 12 ext{HNO}<em>3 ightarrow 2 ext{Fe(NO}</em>3)<em>3 + 3 ext{S} + 6 ext{NO}</em>2 + 6 ext{H}_2 ext{O}$

  • Identify reducing agent.

• Example Reaction:
  $2 ext{HClO}<em>2 + ext{MnO}</em>2 <br>ightarrow ext{Cl}<em>2 + 2 ext{MnO}</em>4^- + 2 ext{H}^+$

  • Identify oxidizing agent.

Oxidation State vs Formal Charge

• Oxidation State (OS):

  • A hypothetical charge that an atom would have if all bonds were ionic (100% electron transfer).

  • Calculation:

    • For atom X:
      OS = ext{# of valence electrons} - rac{1}{2} ext{(bonding electrons)} - ext{(electrons in lone pairs)}

• Formal Charge (FC):

  • Charge assigned to an atom in a molecule under the assumption that electrons are equally shared in bonds.

  • Calculation:

    • For atom X:
      FC = ext{# of valence electrons} - rac{1}{2} ext{(bonding electrons)} - ext{(lone pair electrons)}

  • Example Calculation for O:

    • $FC = 6 - rac{1}{2} imes 6 - 2 = 6 - 5 = +1$

Worksheet for Practice

• 
  

• Determine formal charge and oxidation state for components in NH3:
  

  Atom	Formal Charge	Oxidation State
  N
  H
  General Guidelines for Determining Oxidation States of Atoms

  1. Neutral compounds: sum = 0.

  2. Ions: sum = charge on the ion.

  3. Free elements: oxidation state = 0.

  4. Fluorine: oxidation state = -1.

  5. Group 1A (alkali metals): +1; Group 2A (alkaline earth metals): +2.

  6. Hydrogen with nonmetals: +1; with metals: -1.

  7. Oxygen in compounds: -2, except in peroxides or with fluorine.

  8. Nonmetals: various oxidation states depending on group: Group 17 = -1, Group 16 = -2, Group 15 = -3.