Redox Reactions and Oxidation Numbers

Oxidation Numbers

  • Terminology

    • Redox Reactions (Oxidation-Reduction Reactions)
    • Key Chemical Concept: Net movement of electrons

  • Oxidation

    • Defined as the loss of electrons by a species accompanied by an increase in its oxidation number (getting more positive).
    • Example Reaction:
    • ext{Zn(s)} + 2 ext{H}^+ ightarrow ext{Zn}^{2+} + ext{H}_2(g)
      • Zinc (Zn) loses electrons, identified as the reducing agent.
      • Oxidation number increases from 0 to +2 for Zn.
    • Oxidizing Agent
      • Substance that accepts electrons in a reaction and undergoes a decrease in oxidation number (species is reduced).
      • Example: In the reaction, hydrogen ions (H+) are reduced as they gain electrons.

  • Reduction

    • Defined as the gain of electrons accompanied by a decrease in oxidation number (getting more negative).
    • Example:
    • In the reaction, the hydrogen ions (H+) accept electrons (reduction).
    • Reducing Agent
      • Substance that donates electrons in a reduction reaction and undergoes an increase in oxidation number (the species is oxidized).

  • Understanding Oxidation and Reduction

    • Transfer or shift of electrons is critical.
    • Oxidation number changes reflect the movement of electrons:
    • When oxidation number increases, it indicates an oxidation process.
    • When oxidation number decreases, it indicates a reduction process.

  • Example Calculations of Oxidation Numbers

    • Determine oxidation number for each element in the following compounds:
    • Zinc Chloride (ZnCl₂):
      • Zn: +2, Cl: -1 (Total = 0)
    • Sulfur Trioxide (SO₃):
      • S: +6, O: -2 (Total = 0)
    • Nitric Acid (HNO₃):
      • H: +1, N: +5, O: -2 (Total = 0)
    • Phosphoric Acid (H₃PO₄):
      • H: +1, P: +5, O: -2 (Total = 0)

  • Identifying Oxidized and Reduced Species

    • Review the following reactions to identify oxidizing agents and reducing agents:
    1. 2 ext{Al}(s) + 3 ext{H}_2 SO_4 (aq) ightarrow ext{Al}_2(SO_4)_3 + 3 ext{H}_2(g)
      • Al is oxidized (+3), H is reduced (+1).
      • Al is the reducing agent, H+ is the oxidizing agent.
    2. 2 ext{H}_2(g) + ext{O}_2(g) ightarrow 2 ext{H}_2O (g)
      • H is oxidized (0 to +1), O is reduced (0 to -2).
      • H2 is the reducing agent, O2 is the oxidizing agent.
    3. 2 ext{Fe}(s) + 3 ext{Cl}_2(g) ightarrow 2 ext{FeCl}_3(s)
      • Fe is oxidized (+3), Cl is reduced (-1).
      • Fe is the reducing agent, Cl2 is the oxidizing agent.

Types of Redox Reactions

  • Overview of Redox Reactions

    • Redox reactions involve the transfer of electrons and include various types:
    • Combination/Synthesis
      • Example: 2K(s) + Cl_2(g)
        ightarrow 2KCl(s)
    • Decomposition
      • Example: 2H_2O(l)
        ightarrow 2H_2(g) + O_2(g)
    • Single Replacement
      • Example: 2AgNO_3(aq) + Cu(s)
        ightarrow Cu(NO_3)_2(aq) + 2Ag(s)
    • Combustion
      • Example: Any hydrocarbon combusting in oxygen results in CO₂ and H₂O.

  • Single Replacement Reaction Insights

    • Reactivity Series:
    • Elements can replace others in reactions based on their reactivity.
    • Example Reaction:
      • Cu(s) + 2AgNO_3(aq)
        ightarrow Cu(NO_3)_2(aq) + 2Ag(s)
      • Cu replaces Ag due to higher reactivity.

  • Understanding Reactants and Products

    • Examine reactants and products to determine oxidation and reduction.
    • Tracking changes in oxidation states indicates which species are oxidized and reduced.