📝 Chapter Notes & Revision

Redox Reactions

🏫 MP BoardClass 11Chemistry

📐 Formula & Cheat Sheet (English)

Class 11 Chemistry: Redox Reactions (ऑक्सीकरण एवं अपचयन अभिक्रियाएं)

Quick Revision Notes & Formula Sheet (MP Board)

1. Classical Concepts of Redox Reactions

ProcessClassical DefinitionExample
Oxidation (ऑक्सीकरण)Addition of Oxygen / Electronegative element OR Removal of Hydrogen / Electropositive element2Mg + O2 -> 2MgO<br>H2S + Cl2 -> 2HCl + S
Reduction (अपचयन)Addition of Hydrogen / Electropositive element OR Removal of Oxygen / Electronegative element2HgO -> 2Hg + O2<br>CH2=CH2 + H2 -> CH3-CH3

2. Modern Electronic Concept

  • Oxidation: Process involving loss of electron(s) (De-electronation).
    • Na -> Na+ + e-
  • Reduction: Process involving gain of electron(s) (Electronation).
    • Cl + e- -> Cl-
  • Oxidizing Agent (Oxidant / ऑक्सीकारक): Electron acceptor (gets reduced itself).
  • Reducing Agent (Reductant / अपचायक): Electron donor (gets oxidized itself).

Mnemonic: LEO says GER

  • Loss of Electrons is Oxidation | Gain of Electrons is Reduction.

3. Oxidation Number (Oxidation State / ऑक्सीकरण संख्या)

It is the imaginary or real charge developed on an atom when all other atoms attached to it are removed as ions.

Rules for Calculating Oxidation Number:

  1. Free / Uncombined State: Oxidation number of an element in its free state is always zero.
    • Examples: O2, H2, N2, P4, S8, Na, Fe = 0
  2. Monatomic Ions: Charge on the ion = Oxidation number.
    • Examples: Na+ = +1, Mg2+ = +2, Cl- = -1, Al3+ = +3
  3. Hydrogen:
    • Combined with non-metals = +1 (e.g., H2O, HCl)
    • Combined with metals (Metal Hydrides) = -1 (e.g., NaH, CaH2)
  4. Oxygen:
    • General compounds = -2 (e.g., H2O, CaO)
    • Peroxides (-O-O-) = -1 (e.g., H2O2, Na2O2)
    • Superoxides = -1/2 (e.g., KO2)
    • Oxygen Fluorides = +2 in OF2, +1 in O2F2
  5. Halogens:
    • Fluorine (F) is always -1 (most electronegative element).
    • Cl, Br, I are usually -1, but are positive when combined with more electronegative elements (e.g., in oxoacids like HClO4, Cl = +7).
  6. Alkali Metals (Group 1) = Always +1
  7. Alkaline Earth Metals (Group 2) = Always +2
  8. Sum Rule:
    • Neutral Molecule: Sum of oxidation numbers of all atoms = 0
    • Polyatomic Ion: Sum of oxidation numbers = Charge on the ion

4. Special/Exception Cases of Oxidation Number

These are frequently asked in MP Board examinations due to peroxide linkages or chemical structures:

  1. CrO5 (Chromium peroxide / Butterfly structure):

    • Has 2 peroxide bonds (-O-O-) and 1 oxide bond (=O).
    • Oxidation State of Cr = +6 (not +10).

  2. H2SO5 (Caro's Acid):

    • Contains 1 peroxide linkage.
    • Oxidation State of S = +6 (not +8).

  3. H2S2O8 (Marshall's Acid):

    • Contains 1 peroxide linkage.
    • Oxidation State of S = +6 (not +7).

  4. Br3O8 (Tribromooctaoxide):

    • Average Oxidation State of Br = +16/3
    • Structural Oxidation States: Two Br atoms are in +6 state and one Br atom is in +4 state.

  5. S4O6^2- (Tetrathionate Ion):

    • Average Oxidation State of S = +2.5
    • Structural Oxidation States: Two S atoms are in +5 state and two S atoms are in 0 state.

5. Types of Redox Reactions

  1. Combination Reaction: Two elements or compounds combine.
    • C + O2 -> CO2
  2. Decomposition Reaction: Breakdown of a compound into simpler substances (at least one must be an element).
    • 2KClO3 -> 2KCl + 3O2
  3. Displacement Reaction: An atom/ion in a compound is replaced by an atom/ion of another element.
    • Metal Displacement: Zn + CuSO4 -> ZnSO4 + Cu
    • Non-Metal Displacement: 2Na + 2H2O -> 2NaOH + H2
  4. Disproportionation Reaction (असमानुपातन अभिक्रिया): Same element undergoes simultaneous oxidation and reduction.
    • 2H2O2 -> 2H2O + O2 (O goes from -1 to -2 and 0)
    • P4 + 3NaOH + 3H2O -> PH3 + 3NaH2PO2 (P goes from 0 to -3 and +1)

6. Balancing Redox Reactions

A. Oxidation Number Method

  1. Write skeletal ionic equation.
  2. Assign oxidation numbers to all elements and identify elements undergoing oxidation and reduction.
  3. Calculate the change in oxidation number per atom and balance the total change in oxidation numbers by multiplying with suitable coefficients.
  4. Balance atoms other than H and O.
  5. Balance Charge & Hydrogen/Oxygen:
    • Acidic Medium: Add H2O to balance O-atoms, then add H+ to balance H-atoms.
    • Basic Medium: Add H2O to balance O-atoms, then balance H by adding H2O to deficient side and equal number of OH- to the opposite side.

B. Ion-Electron Method (Half-Reaction Method)

  1. Split equation into two half-reactions: Oxidation Half and Reduction Half.
  2. Balance all atoms except H and O.
  3. Balance O atoms by adding H2O to the oxygen-deficient side.
  4. Balance H atoms:
    • Acidic Medium: Add H+ to H-deficient side.
    • Basic Medium: Add equal number of OH- ions to both sides for every H+ added.
  5. Balance charge by adding electrons (e-).
  6. Equalize the number of electrons in both half-reactions by multiplying with suitable integers and add the two half-reactions.

7. Redox Reactions and Electrode Processes

  • Redox Couple (रेडॉक्स युग्म): Consists of oxidized and reduced forms of an element involved in an oxidation or reduction half-reaction (e.g., Zn2+/Zn, Cu2+/Cu).
  • Electrochemical Cell (Galvanic/Voltaic Cell): Device that converts chemical energy into electrical energy through redox reactions.

Key Cell Terminology (Mnemonic: LOAN)

  • Left side
  • Oxidation takes place
  • Anode
  • Negative electrode
TerminalElectrode ProcessSign (Galvanic Cell)Mnemonic
AnodeOxidation (Loss of electrons)Negative (-)AN OX
CathodeReduction (Gain of electrons)Positive (+)RED CAT

Standard Electrode Potential ($E^\circ$)

  • Measured relative to Standard Hydrogen Electrode (SHE) whose potential is arbitrarily taken as 0.00 V.
  • Electromotive Force (EMF) Formula: $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$ $$E^\circ_{\text{cell}} = E^\circ_{\text{Right}} - E^\circ_{\text{Left}}$$

Electrochemical Series (इलेक्ट्रोकेमिकल श्रेणी)

  • Arrangement of elements in increasing order of their standard reduction potentials ($E^\circ$).
  • Key Points:
    • Higher standard reduction potential ($E^\circ > 0$) $\rightarrow$ Stronger Oxidizing Agent (e.g., $F_2$, $E^\circ = +2.87\text{ V}$).
    • Lower standard reduction potential ($E^\circ < 0$) $\rightarrow$ Stronger Reducing Agent (e.g., $Li$, $E^\circ = -3.05\text{ V}$).
    • A metal with lower $E^\circ$ can displace a metal with higher $E^\circ$ from its salt solution.

8. Summary Formula Flash-Cards

  • Oxidation State Equation: $$\text{Sum of (Number of atoms } \times \text{ Oxidation state)} = \text{Total Charge}$$
  • Cell EMF: $$E^\circ_{\text{cell}} = E^\circ_{\text{reduction (cathode)}} - E^\circ_{\text{reduction (anode)}}$$
  • Feasibility of a Redox Reaction:
    • If $E^\circ_{\text{cell}} > 0$ (Positive) $\rightarrow$ Reaction is spontaneous / feasible.
    • If $E^\circ_{\text{cell}} < 0$ (Negative) $\rightarrow$ Reaction is non-spontaneous.