Redox Reactions revision notes

Key Concepts & Definitions

Oxidation (Classical):

Addition of oxygen or an electronegative element, or the removal of hydrogen or an electropositive element. Example of electropositive element removal: 2K4[Fe(CN)6]+H2O2→2K3[Fe(CN)6]+2KOH2K_4[Fe(CN)_6] + H_2O_2 \rightarrow 2K_3[Fe(CN)_6] + 2KOH2K4​[Fe(CN)6​]+H2​O2​→2K3​[Fe(CN)6​]+2KOH (Potassium is removed from ferrocyanide).

Reduction (Classical):

Removal of oxygen or an electronegative element, or the addition of hydrogen or an electropositive element. Example of electropositive element addition: 2HgCl2+SnCl2→Hg2Cl2+SnCl42HgCl_2 + SnCl_2 \rightarrow Hg_2Cl_2 + SnCl_42HgCl2​+SnCl2​→Hg2​Cl2​+SnCl4​ (Mercury is added to mercuric chloride).

Oxidation (Electron Transfer):

Loss of electron(s) by any species.

Reduction (Electron Transfer):

Gain of electron(s) by any species.

Oxidising Agent (Oxidant):

An electron acceptor that increases the oxidation number of another species.

Reducing Agent (Reductant):

An electron donor that lowers the oxidation number of another species.

Oxidation Number (O.N.):

The assigned oxidation state based on the assumption that an electron pair in a covalent bond entirely belongs to the more electronegative element.

Stock Notation:

Representing the O.N. of a metal in a compound using a Roman numeral in parentheses immediately after the metal's symbol (e.g., Au(III)Cl3Au(III)Cl_3Au(III)Cl3​, Hg2(I)Cl2Hg_2(I)Cl_2Hg2​(I)Cl2​).

Redox Couple:

Represents the oxidised and reduced forms of a substance taking part in a half-reaction, separated by a vertical line or slash representing an interface (e.g., Zn2+/ZnZn^{2+}/ZnZn2+/Zn).

Electrode Potential:

The potential difference between an electrode and its ions in solution.

Standard Electrode Potential (E0E^0E0):

The potential measured under standard conditions (unity concentration, 1 atm pressure for gases, 298 K) relative to the Standard Hydrogen Electrode (E0=0.00E^0 = 0.00E0=0.00 V).

Daniell Cell:

A galvanic cell where Zn acts as the anode (oxidation) and Cu as the cathode (reduction).

Salt Bridge:

A U-tube containing KClKClKCl or NH4NO3NH_4NO_3NH4​NO3​ in agar-agar that completes the electrical circuit and allows ion migration without physical mixing of solutions.

JEE TIPModern environmental phenomena like the "Hydrogen Economy" (liquid hydrogen as fuel) and "Ozone Hole" development are governed by redox principles.

JEE TIPOxidation is viewed as a decrease in electron density around an atom, while reduction is an increase in electron density.

Important Rules, Laws & Principles

• Rules for Assigning Oxidation Number:
  • In the free or elemental state, O.N. is always zero (e.g., $H_2, P_4, S_8, Na$).
  • For monoatomic ions, O.N. equals the charge.
  • Alkali metals (Group 1) are always +1; alkaline earth metals (Group 2) are always +2; Aluminium is always +3.
  • Oxygen is generally -2, Hydrogen is generally +1, and Fluorine is ALWAYS -1.
  • The algebraic sum of O.N. of all atoms in a neutral compound is zero; for polyatomic ions, it equals the net charge.
• Maximum Oxidation Number Rule: The highest O.N. of a representative element generally equals its group number (Groups 1 & 2) or group number minus 10 (p-block elements 13-17).
• JEE TIPIf an element is in its highest possible oxidation state (e.g., Cl in $ClO_4^-$ is +7), it can only act as an oxidant and cannot disproportionate. An element in an intermediate state can act as both an oxidant and a reductant.
• Electrochemical Activity Principle (Competitive Electron Transfer): A metal with a more negative $E^0$ is a stronger reducing agent and will displace a metal with a less negative/positive $E^0$ from its salt solution (Reducing power: $Zn > Cu > Ag$).
• JEE TIPWhen cobalt ($Co$) is placed in nickel sulphate ($NiSO_4$), an equilibrium is reached where neither reactants nor products are greatly favored, leaving moderate concentrations of both $Co^{2+}$ and $Ni^{2+}$.

Types of Redox Reactions

• Combination Reactions: $A + B \rightarrow C$, where at least one reactant is in elemental form (e.g., $3Mg + N_2 \rightarrow Mg_3N_2$).
• Decomposition Reactions: Breakdown of a compound where at least one product is in the elemental state (e.g., $2NaH \rightarrow 2Na + H_2$; $2KClO_3 \rightarrow 2KCl + 3O_2$).
• Displacement Reactions:
  • Metal Displacement: Reactive metals displace less reactive ones ($V_2O_5 + 5Ca \rightarrow 2V + 5CaO$).
  • Hydrogen Displacement: Very active metals (Na, Ca) displace $H_2$ from cold water; less active (Mg, Fe) from steam; others (Zn) from acids. Noble metals like Ag and Au do not react even with strong acids like HCl.
  • Halogen Displacement: A more reactive halogen displaces a heavier halide from solution ($Cl_2 + 2KI \rightarrow 2KCl + I_2$).
• Disproportionation Reactions: An element in an intermediate oxidation state is simultaneously oxidised and reduced.
  • $P_4 + 3OH^- + 3H_2O \rightarrow PH_3 + 3H_2PO_2^-$.
  • $S_8 + 12OH^- \rightarrow 4S^{2-} + 2S_2O_3^{2-} + 6H_2O$.
  • $Cl_2 + 2OH^- \rightarrow ClO^- + Cl^- + H_2O$ (Household bleach formation).
  • $(CN)_2 + 2OH^- \rightarrow CN^- + CNO^- + H_2O$.

Redox Titrations & Mechanisms

• Self-Indicators: Intensely coloured reagents like $MnO_4^-$ act as self-indicators. The visible end point is achieved after the last of the reductant is consumed, leaving the first lasting tinge of pink colour.
• External Indicators: $Cr_2O_7^{2-}$ requires an indicator like diphenylamine, which turns intense blue immediately after the equivalence point.
• Iodometric Titrations: Rely on reactions where $Cu^{2+}$ oxidises $I^-$ to $I_2$ (forming insoluble $Cu_2I_2$). The liberated $I_2$ (present as $KI_3$) is titrated with thiosulphate ($S_2O_3^{2-}$). Starch is used as an indicator, forming an intense blue complex with $I_2$ that disappears at the end point.
• Daniell Cell Circuitry: Electrons travel from the anode (Zn) to the cathode (Cu) through the external wire. The direction of current is opposite to the direction of electron flow.

Formulae & Equations

• Balancing by Oxidation Number Method:
  1. Identify atoms undergoing redox and assign O.N..
  2. Calculate total increase and decrease in O.N. and cross-multiply to equalize.
  3. Add $H^+$ (acidic) or $OH^-$ (basic) to balance ionic charges.
  4. Add $H_2O$ to balance hydrogen/oxygen atoms.
• Balancing by Half-Reaction (Ion-Electron) Method:
  1. Separate into oxidation and reduction half-reactions.
  2. Balance atoms other than O and H. Add $H_2O$ for oxygen, $H^+$ for hydrogen.
  3. For basic mediums: Add $OH^-$ to BOTH sides equal to the number of $H^+$ added, and combine $H^+$ and $OH^-$ into $H_2O$.
  4. Add electrons ($e^-$) to balance charges and equalize electrons between halves.

Trends & Comparisons

• Halogen Oxidising Power: Decreases down the group: $F_2 > Cl_2 > Br_2 > I_2$.
• Electron Releasing Tendency (Reducing Power of Metals): $Zn > Cu > Ag$.
• Standard Electrode Potentials ($E^0$):
  • $F_2 / F^-$ couple: Highest positive $E^0$ (+2.87 V) $\rightarrow$ $F_2$ is the strongest oxidising agent.
  • $Li^+ / Li$ couple: Highest negative $E^0$ (-3.05 V) $\rightarrow$ Lithium metal is the strongest reducing agent.

⚠️ EXCEPTIONS & ANOMALIES

• Oxygen O.N. Anomalies: Normally -2, BUT:
  • -1 in peroxides ($H_2O_2, Na_2O_2$).
  • -1/2 in superoxides ($KO_2, RbO_2$).
  • +2 in $OF_2$ and +1 in $O_2F_2$.
  • Fractional states: +1/2 in $O_2^+$ and -1/2 in $O_2^-$.
• Hydrogen O.N. Anomalies: Normally +1, BUT it is -1 in binary metal hydrides (e.g., $LiH$, $NaH$, $CaH_2$).
• The Paradox of Fractional Oxidation Numbers: Electrons are never shared in fractions. A fractional state is merely an average of different integer states based on structure.
  • Carbon Suboxide ($C_3O_2$): $O=C=C=C=O$. Terminal carbons are +2, central is 0 (Average = +4/3).
  • Tribromooctaoxide ($Br_3O_8$): Terminal bromines are +6, central is +4 (Average = +16/3).
  • Tetrathionate ($S_4O_6^{2-}$): Two terminal sulphurs are +5, two central sulphurs are 0 (Average = 2.5).
  • Mixed Oxides: $Fe_3O_4$, $Mn_3O_4$, and $Pb_3O_4$ are stoichiometric mixtures. E.g., $Pb_3O_4$ is a mixture of 2 moles of $PbO$ (+2) and 1 mole of $PbO_2$ (+4).
• Fluorine's Extreme Singular Behavior:
  • Being the most electronegative, it never exhibits a positive oxidation state.
  • Does NOT Disproportionate: Unlike $Cl_2$, fluorine cannot disproportionate in alkali. Instead, it forms $F^-$ and $OF_2$ ($2F_2 + 2OH^- \rightarrow 2F^- + OF_2 + H_2O$).
  • Aqueous Displacement Anomaly: $F_2$ is too reactive to displace other halogens in water; it attacks water directly ($2H_2O + 2F_2 \rightarrow 4HF + O_2$).
  • Oxidation of $F^-$: There is no standard chemical way to oxidise $F^-$ to $F_2$ (electrolysis is required). The single anomaly: $XeO_6^{4-}$ is powerful enough to chemically oxidize $F^-$ to $F_2$ ($XeO_6^{4-} + 2F^- + 6H^+ \rightarrow XeO_3 + F_2 + 3H_2O$).
• Reaction Path Anomalies ($Pb_3O_4$):
  • With $HCl$, $Pb_3O_4$ undergoes a redox reaction (yielding $PbCl_2 + Cl_2$) because $PbO_2$ oxidises $Cl^-$.
  • With $HNO_3$, it undergoes an acid-base reaction only (yielding $Pb(NO_3)_2 + PbO_2$) because $HNO_3$ is an oxidant itself and $PbO_2$ is passive against it.
• Decomposition Without Redox: Thermal decomposition of $CaCO_3 \rightarrow CaO + CO_2$ involves no change in oxidation numbers.
• Strictly Oxidants: While $SO_2$ and $H_2O_2$ can be both oxidants and reductants, Ozone ($O_3$) and Nitric Acid ($HNO_3$) can act ONLY as oxidants.
• Unstable Oxidants ($AgF_2$): Silver prefers a +1 state. The compound $AgF_2$ is highly unstable and acts as a very strong oxidising agent as it forces its way back to a stable state.

Previous Year JEE Topics

• Exact vs. Average Oxidation States: Deducing true integer states from structures of $C_3O_2$, $S_4O_6^{2-}$, and $Br_3O_8$.
• Disproportionation Rules: Identifying intermediate oxoanions ($ClO_2^-, ClO_3^-$) that can disproportionate vs. those that cannot ($ClO_4^-, F_2$).
• Balancing in Basic Medium: Mastery of adding $OH^-$ to neutralize $H^+$ during half-reaction balancing.
• Redox Titration Stoichiometry: Finding endpoints and reacting moles, especially in Iodometric titrations involving $Cu^{2+}$ and $S_2O_3^{2-}$.
• Fluorine's Anomalous Reactions: Reactivity with water, alkalis, and identifying the exceptionally rare reaction ($XeO_6^{4-}$) that oxidizes $F^-$.

Memory Aids & JEE Traps

• JEE TIP$Cl_2$ displaces $Br^-$ and $I^-$, liberating $Br_2$ and $I_2$. These dissolve in organic solvents like $CCl_4$ to form distinct coloured layers used for lab identification.
• JEE TIP
  • Misconception $\rightarrow$ Electrons can be shared or transferred in fractions, resulting in true fractional oxidation states.
  • Correct Understanding $\rightarrow$ A fractional O.N. is strictly an average of distinct integer states in the structure (e.g., $S_4O_6^{2-}$ has two +5 sulphurs and two 0 sulphurs, averaging to 2.5).
• JEE TIP
  • Misconception $\rightarrow$ All decomposition reactions are inherently redox reactions.
  • Correct Understanding $\rightarrow$ A decomposition is only redox if at least one product is in the elemental state. The decomposition of $CaCO_3$ is not a redox reaction.
• JEE TIP
  • Misconception $\rightarrow$ $F_2$ disproportionates in alkaline mediums just like chlorine and bromine.
  • Correct Understanding $\rightarrow$ Fluorine never exhibits a positive oxidation state, so it does not disproportionate. It yields $F^-$, $OF_2$, and $H_2O$ instead.
• JEE TIP
  • Misconception $\rightarrow$ All oxoanions of halogens ($ClO^-$, $ClO_2^-$, $ClO_3^-$, $ClO_4^-$) can undergo disproportionation.
  • Correct Understanding $\rightarrow$ Elements in their maximum oxidation state (like Cl at +7 in $ClO_4^-$) can only act as oxidants and cannot disproportionate.
• JEE TIP
  • Misconception $\rightarrow$ Because $F_2$ is the strongest oxidant, it is used to displace $Cl^-, Br^-$, and $I^-$ in aqueous solutions.
  • Correct Understanding $\rightarrow$ $F_2$ is so reactive it attacks water directly to produce HF and $O_2$; thus, halogen displacement with fluorine isn't carried out in aqueous solutions.
• JEE TIP
  • Misconception $\rightarrow$ The mixed oxide $Pb_3O_4$ yields the same type of products with all strong acids.
  • Correct Understanding $\rightarrow$ It gives a redox reaction with HCl (yielding $Cl_2$) but gives an acid-base reaction with $HNO_3$ (leaving solid $PbO_2$ unreacted).
• JEE TIP
  • Misconception $\rightarrow$ Thiosulphate ($S_2O_3^{2-}$) always oxidizes to the same product.
  • Correct Understanding $\rightarrow$ It reacts differently based on oxidant strength: it yields tetrathionate ($S_4O_6^{2-}$) with weak oxidants like $I_2$, but oxidizes completely to sulphate ($SO_4^{2-}$) with strong oxidants like $Br_2$ or $Cl_2$.
• JEE TIP
  • Misconception $\rightarrow$ A highly negative Standard Electrode Potential ($E^0$) means the substance is a weak reducing agent.
  • Correct Understanding $\rightarrow$ A negative $E^0$ means the redox couple is a stronger reducing agent than the $H^+/H_2$ couple. Lithium has $E^0$ of -3.05 V and is the strongest reducing agent.
• JEE TIP
  • Misconception $\rightarrow$ All oxygen and nitrogen compounds can act as both oxidising and reducing agents.
  • Correct Understanding $\rightarrow$ Ozone ($O_3$) and Nitric Acid ($HNO_3$) are locked in oxidation states that allow them to act only as oxidants.
• JEE TIP
  • Misconception $\rightarrow$ $F^-$ can be oxidized to $F_2$ using standard strong chemical oxidants like $KMnO_4$ or $K_2Cr_2O_7$.
  • Correct Understanding $\rightarrow$ $F^-$ can only be oxidized to $F_2$ electrolytically. The single chemical exception in JEE syllabus is its rare reaction with the $XeO_6^{4-}$ ion.