REDOX REACTION- OXIDATIO AND REDUCTION-

  • Redox reaction- In this reaction, oxidation and reduction occur simultaneously.
  • Oxidation-
    • adding of oxygen or removal of hydrogen.
    • Adding of electronegative component or removal of electropositive component.
  • Reduction-
    • adding of hydrogen or removal of oxygen.
    • Adding of electropositive component or removal of electronegative component.

REDOX REACTION IN TERMS OF TRANSFER OF ELECTRONS-

  • In this process, 2 factor involving, first is loss of el– and second is gain of el–.
  • Each step is called half reaction, some of half reactions give full reaction.
  • Half reaction called oxidation, similarly other half reaction is called reduction.
  • Oxidizing agent(or Oxidant)– It gain electrons and cause the oxidation of another substance(lose of electron). It accept electrons.
  • Reducing agent(or Reductant)- It loss electron and cause of reduction of another substance to gain electron. It donate electrons.

IMPORTANT REACTIONS OF ELECTRON TRANSFER-

Reaction between Zinc and Copper Nitrate:

When metallic zinc is placed in an aqueous copper nitrate solution, zinc gets oxidized (loses electrons) to form Zn²⁺ ions, and copper ions (Cu²⁺) get reduced (gain electrons) to form copper metal.

The reaction is:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

This is a redox reaction: Zinc is oxidized, and copper is reduced.

Zinc’s electrons reduce copper ions, forming metallic copper.

State of Equilibrium:

If a copper strip is placed in a zinc sulfate solution, no reaction occurs, indicating that the equilibrium favors the products of the first reaction (Zn²⁺ and Cu).

Electron Transfer in Copper and Silver Nitrate Solution:

Copper reacts with silver nitrate to form Cu²⁺ ions and metallic silver.

The reaction shows that copper is oxidized, and silver ions are reduced.

Cobalt and Nickel Sulfate Reaction:

In a reaction between cobalt metal and nickel sulfate solution, equilibrium shows the formation of both Ni²⁺ and Co²⁺ ions, indicating a less favorable redox reaction.

Electron Release Tendencies:

The tendency of metals to lose electrons (oxidize) follows a certain order:
Zn > Cu > Ag.

This is used to design Galvanic cells, which convert chemical energy into electrical energy.

Oxidation Number:

  1. Definition:
    • The oxidation number (or state) represents the charge an atom would have if the compound were ionic, assuming the electrons in a covalent bond belong entirely to the more electronegative atom.
  2. Basic Rules for Determining Oxidation Numbers:
    • Elements in free state (e.g., H₂, O₂, Cl₂) have an oxidation number of 0.
    • Monatomic ions have oxidation numbers equal to their charge: Na⁺ (+1), Mg²⁺ (+2), Cl⁻ (-1).
    • Oxygen generally has an oxidation number of -2, except in peroxides (-1) and superoxides (½).
    • Hydrogen has an oxidation number of +1, except when bonded to metals, where it is -1.
    • Fluorine always has an oxidation number of -1 in compounds.
    • Sum of oxidation numbers in a neutral compound is zero; in polyatomic ions, it equals the ion’s charge.
  3. Stock Notation:
    • The oxidation state of an element in a compound is sometimes shown using Roman numerals in parentheses.
      Example: Cu(I) for Cu⁺ and Cu(II) for Cu²⁺.
  4. Example Redox Reactions:
    • Hydrogen and Oxygen Reaction (H₂ + O₂ → H₂O):
      • H goes from 0 to +1 (oxidized), O goes from 0 to -2 (reduced).
    • Cobalt and Nickel Reaction:
      • Neither cobalt nor nickel undergoes significant oxidation/reduction, hence equilibrium favors both reactants and products.

Types of Redox Reactions:

  1. Combination Reactions:
    • Two or more substances combine to form one product. Examples:
      • C + O₂ → CO₂
      • Mg + N₂ → Mg₃N₂
  2. Decomposition Reactions:
    • A compound breaks down into two or more simpler substances. Examples:
      • 2H₂O → 2H₂ + O₂
      • 2NaH → 2Na + H₂
  3. Displacement Reactions:
    • Involves the replacement of an ion or atom in a compound with an ion or atom from another substance.
      • Metal Displacement: A metal displaces another from a compound. Example:
        Zn + CuSO₄ → Cu + ZnSO₄
      • Non-metal Displacement: Includes reactions where non-metals like hydrogen are displaced. Example:
        Na + H₂O → NaOH + H₂

Summary of Redox Reactions:

Redox Reactions: Involve both oxidation and reduction processes.

Oxidation: Increase in oxidation number (loss of electrons).

Reduction: Decrease in oxidation number (gain of electrons).

Oxidizing Agent: The substance that gains electrons (causes oxidation).

Reducing Agent: The substance that loses electrons (causes reduction).

Disproportionation Reactions

  • Definition: Disproportionation reactions are a type of redox reaction where an element in one oxidation state is both oxidized and reduced.
  • Key Feature: In these reactions, the reacting substance contains an element that can exist in at least three oxidation states. The element is in an intermediate oxidation state, which is then converted to both higher and lower oxidation states.
  • Oxygen in H₂O₂ is in the -1 oxidation state. It is reduced to -2 in water (H₂O) and oxidized to 0 in O₂ gas.

Fluorine’s Deviation:

  • Fluorine, the most electronegative element, does not show a disproportionation tendency because it cannot exhibit a positive oxidation state. When it reacts with alkali, it produces: 2F2+2OH−→2F−+OF2+H2O

Disproportionation of Chlorine Oxyanions:

  • Chlorine oxyanions can undergo disproportionation, except for ClO4​ (where chlorine is already in its highest oxidation state of +7)

Fractional Oxidation Numbers:

  • In certain compounds, elements can have fractional oxidation numbers due to averaging between multiple oxidation states.
    • Examples:
      • C₃O₂: Carbon’s oxidation number is an average of +2 and 0.
      • Br₃O₈: Bromine has oxidation states of +6 and +4.
      • S₄O₆²⁻: Sulfur’s oxidation numbers range from +5 to 0.

Types of Redox Reactions:

  1. Combination Redox Reactions: Two elements or compounds combine to form a product
  2. Decomposition Redox Reactions: A compound breaks down into simpler substances
  3. Displacement Redox Reactions: An element replaces another element in a compound
  4. Disproportionation Redox Reactions: An element undergoes both oxidation and reduction in the same reaction

Balancing Redox Reactions:

  • Oxidation Number Method:
    • Identify atoms changing oxidation states.
    • Calculate the change in oxidation number and balance them by multiplying if necessary.
    • Add H⁺ or OH⁻ ions to balance charges in acidic or basic solutions.
    • Add H₂O to balance oxygen atoms and check the atom and charge balance.
  • Half-Reaction Method:
    • Split the reaction into two half-reactions: one for oxidation and one for reduction.
    • Balance atoms and charges in each half-reaction.
    • Combine the half-reactions, cancel out electrons, and check the balance.

Conclusion:

Understanding disproportionation reactions, fractional oxidation states, and how to balance redox reactions is essential in both conceptual chemistry and for solving competitive exams. The method of assigning oxidation numbers and using half-reactions helps in systematically approaching complex reactions.

  1. Redox Reaction Overview:
    • A redox reaction involves oxidation (loss of electrons) and reduction (gain of electrons) occurring simultaneously.
    • In the example with zinc and copper sulphate, zinc undergoes oxidation (Zn → Zn²⁺ + 2e⁻) and copper ions are reduced (Cu²⁺ + 2e⁻ → Cu).
  2. Modification of the Experiment:
    • The experiment is modified to separate the zinc metal from the copper sulphate solution, allowing indirect electron transfer.
    • Zinc and copper are placed in separate beakers containing their respective salt solutions (zinc sulphate and copper sulphate).
    • A salt bridge (U-tube with potassium chloride or ammonium nitrate gel) connects the two solutions, allowing ions to flow but preventing the mixing of the solutions.
  3. Redox Couples:
    • A redox couple is the combination of oxidized and reduced forms of a species involved in redox reactions.
    • Example: The redox couples in the experiment are Zn²⁺/Zn and Cu²⁺/Cu. In notation, the oxidized form is placed before the reduced form, separated by a vertical line (e.g., Zn²⁺|Zn).
  4. Electrodes and Electrodes Potential:
    • Electrodes are conductive materials where oxidation and reduction occur. The electrode potential is the ability of a substance to gain or lose electrons at the electrode.
    • Standard Electrode Potential: The electrode potential when ion concentrations are 1 mol/L, and the temperature is 298K. By convention, the standard electrode potential of the hydrogen electrode (H₂/H⁺) is taken as 0.00V.
    • A positive E₀ indicates a weaker reducing agent than H₂/H⁺, while a negative E₀ suggests a stronger reducing agent than H₂/H⁺.
  5. Working of the Daniell Cell:
    • In the Daniell Cell, two electrodes (zinc and copper) are connected via a salt bridge and a wire with an ammeter.
    • When the circuit is completed, electrons flow from zinc (the anode, where oxidation occurs) to copper (the cathode, where reduction occurs).
    • Ions migrate through the salt bridge to maintain electrical neutrality, ensuring the current flows.
  6. Current Flow:
    • The current flows opposite to the electron flow in the external circuit. Electrons flow from the zinc rod to the copper rod through the wire.
  7. Standard Electrode Potentials Table:
    • The table of standard electrode potentials (E₀) for various redox couples helps in determining the strength of reducing and oxidizing agents.
    • Higher positive E₀ values represent stronger oxidizing agents, and lower (more negative) E₀ values represent stronger reducing agents.
  8. Important Concepts:
    • Oxidizing Agent (Oxidant): Gains electrons and is reduced.
    • Reducing Agent (Reductant): Loses electrons and is oxidized.

These all are the notes of Redox Reaction. After some time you get important and NCERT solutions HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊