Redox Rxn

. Basics of Redox Reactions

  • Definition:

    • Redox reactions involve the transfer of electrons between two species.

    • Oxidation: Loss of electrons or increase in oxidation state.

    • Reduction: Gain of electrons or decrease in oxidation state.

  • Key Terms:

    • Oxidizing Agent: Gains electrons (undergoes reduction).

    • Reducing Agent: Loses electrons (undergoes oxidation).

2. Oxidation Number (ON)

  • Rules for Assigning Oxidation Numbers:

    1. Free elements (e.g., O2,H2O_2, H_2O2​,H2​) have ON = 0.

    2. Monoatomic ions: ON = charge of the ion (e.g., Na+Na^+Na+: ON = +1).

    3. Hydrogen:

      • With non-metals: +1 (e.g., HClHClHCl).

      • With metals (hydrides): -1 (e.g., NaHNaHNaH).

    4. Oxygen:

      • Usually -2.

      • In peroxides (e.g., H2O2H_2O_2H2​O2​): -1.

      • In superoxides (e.g., KO2KO_2KO2​): -1/2.

    5. Sum of ONs in a neutral compound = 0.

    6. Sum of ONs in a polyatomic ion = charge of the ion.

  • Examples:

    • H2SO4H_2SO_4H2​SO4​: H=+1,O=−2H = +1, O = -2H=+1,O=−2, S = +6 (from 2(+1)+4(−2)+x=02(+1) + 4(-2) + x = 02(+1)+4(−2)+x=0).

3. Balancing Redox Reactions

  • Ion-Electron Method (Half-Reaction Method):

    1. Separate into oxidation and reduction half-reactions.

    2. Balance all atoms except OOO and HHH.

    3. Balance OOO by adding H2OH_2OH2​O; balance HHH by adding H+H^+H+.

    4. Balance charge by adding e−e^-e−.

    5. Equalize electrons in both half-reactions and combine.

  • Oxidation Number Method:

    1. Assign oxidation numbers to all atoms.

    2. Identify species undergoing oxidation and reduction.

    3. Balance the total increase and decrease in ON by multiplying species appropriately.

    4. Balance other atoms and charges as needed.

4. Types of Redox Reactions

  • Combination Reaction: A+B→ABA + B → ABA+B→AB (both oxidation and reduction occur). Example: 2H2+O2→2H2O2H_2 + O_2 → 2H_2O2H2​+O2​→2H2​O.

  • Decomposition Reaction: AB→A+BAB → A + BAB→A+B. Example: 2KClO3→2KCl+3O22KClO_3 → 2KCl + 3O_22KClO3​→2KCl+3O2​.

  • Displacement Reaction:

    • Metal Displacement: Zn+CuSO4→ZnSO4+CuZn + CuSO_4 → ZnSO_4 + CuZn+CuSO4​→ZnSO4​+Cu.

    • Non-metal Displacement: Cl2+2KBr→2KCl+Br2Cl_2 + 2KBr → 2KCl + Br_2Cl2​+2KBr→2KCl+Br2​.

  • Disproportionation Reaction:

    • Single element undergoes both oxidation and reduction. Example: Cl2+H2O→HCl+HClOCl_2 + H_2O → HCl + HClOCl2​+H2​O→HCl+HClO.

5. Redox Titrations

  • Key Types:

    • KMnO4KMnO_4KMnO4​ Titration: Acts as an oxidizing agent in acidic medium (reduced to Mn2+Mn^{2+}Mn2+). Example: MnO4−+8H++5e−→Mn2++4H2OMnO_4^- + 8H^+ + 5e^- → Mn^{2+} + 4H_2OMnO4−​+8H++5e−→Mn2++4H2​O.

    • K2Cr2O7K_2Cr_2O_7K2​Cr2​O7​ Titration: Cr2O72−+14H++6e−→2Cr3++7H2OCr_2O_7^{2-} + 14H^+ + 6e^- → 2Cr^{3+} + 7H_2OCr2​O72−​+14H++6e−→2Cr3++7H2​O.

  • Indicators:

    • Self-indicating (e.g., KMnO4KMnO_4KMnO4​).

    • External (e.g., Starch in iodometric titration).

6. Electrochemical Aspects

  • Electrochemical Cells:

    • Oxidation at Anode (Negative).

    • Reduction at Cathode (Positive).

  • EMF Calculation (Nernst Equation):

    Ecell=Ecell∘−0.0591nlog⁡QE_{cell} = E^\circ_{cell} - \frac{0.0591}{n} \log QEcell​=Ecell∘​−n0.0591​logQ

    Where Q=[products][reactants]Q = \frac{[\text{products}]}{[\text{reactants}]}Q=[reactants][products]​.

  • Standard Electrode Potential:

    • High E∘E^\circE∘: Strong oxidizing agent.

    • Low E∘E^\circE∘: Strong reducing agent.

7. Practical Applications

  • Industrial Applications:

    • Electrolysis (e.g., extraction of aluminum).

    • Batteries: Lithium-ion, Lead-acid, etc.

  • Biological Processes:

    • Cellular respiration (involves redox reactions).