JEE: Redox Reactions

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1. Fundamental Concepts & Definitions 1.1. Oxidation & Reduction Oxidation: Classical Concept: Addition of Oxygen (e.g., $C + O_2 \rightarrow CO_2$) Removal of Hydrogen (e.g., $H_2S + Cl_2 \rightarrow 2HCl + S$) Addition of Electronegative element (e.g., $Fe + S \rightarrow FeS$) Removal of Electropositive element (e.g., $2KI + Cl_2 \rightarrow 2KCl + I_2$) Electronic Concept: Loss of one or more electrons by an atom, ion, or molecule. (e.g., $Na \rightarrow Na^+ + e^-$) Oxidation Number Concept: Increase in the oxidation number (O.N.) of an element. Reduction: Classical Concept: Removal of Oxygen (e.g., $CuO + H_2 \rightarrow Cu + H_2O$) Addition of Hydrogen (e.g., $Cl_2 + H_2S \rightarrow 2HCl + S$) Removal of Electronegative element (e.g., $2FeCl_3 + H_2 \rightarrow 2FeCl_2 + 2HCl$) Addition of Electropositive element (e.g., $HgCl_2 + SnCl_2 \rightarrow Hg_2Cl_2 + SnCl_4$) Electronic Concept: Gain of one or more electrons by an atom, ion, or molecule. (e.g., $Cl_2 + 2e^- \rightarrow 2Cl^-$) Oxidation Number Concept: Decrease in the oxidation number (O.N.) of an element. Limitations of Classical Definitions: Not applicable when no oxygen/hydrogen is involved, or for organic/complex compounds. Electronic and O.N. concepts are more universal. 1.2. Redox Reaction, Agents, and Couple Redox Reaction: Any chemical reaction in which the oxidation numbers of atoms are changed; involves simultaneous oxidation and reduction. It is fundamentally an electron transfer process. Oxidising Agent (Oxidant): A substance that causes oxidation in another substance and is itself reduced. It accepts electrons. (e.g., $KMnO_4, K_2Cr_2O_7, O_2, Cl_2$). Reducing Agent (Reductant): A substance that causes reduction in another substance and is itself oxidised. It donates electrons. (e.g., $H_2, C, Na, H_2S, SnCl_2$). Redox Couple: Consists of the oxidised and reduced forms of a substance taking part in a half-reaction. It is represented as $Oxidised form/Reduced form$ (e.g., $Zn^{2+}/Zn$, $Fe^{3+}/Fe^{2+}$). 1.3. Oxidation State vs. Formal Charge vs. Valency Oxidation State (O.N.): A hypothetical charge assigned to an atom in a molecule or ion, assuming that all bonds are ionic and electrons in a bond are completely transferred to the more electronegative atom. It's a tool for electron accounting in redox. Formal Charge: The charge assigned to an atom in a molecule, assuming that electrons in a chemical bond are shared equally between the atoms, regardless of electronegativity. It helps in determining the most plausible Lewis structure. Valency: The combining capacity of an element, typically equal to the number of bonds an atom can form. It is always a positive integer. 2. Electronic Concept & Types of Redox Reactions 2.1. Electron Transfer & Half-Reactions Redox reactions are seen as two separate processes: an oxidation half-reaction (electron loss) and a reduction half-reaction (electron gain). Oxidation: $Zn \rightarrow Zn^{2+} + 2e^-$ Reduction: $Cu^{2+} + 2e^- \rightarrow Cu$ Overall Redox: $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$ 2.2. Special Types of Redox Reactions based on O.N. Change Internal Redox: When an element within the same compound undergoes both oxidation and reduction, but the atoms are different (e.g., $NH_4NO_3 \xrightarrow{\Delta} N_2O + 2H_2O$. Here $N$ from $NH_4^+$ is oxidised, $N$ from $NO_3^-$ is reduced). Disproportionation (Dismutation) Reaction: A single element in one reactant is simultaneously oxidised and reduced, resulting in products containing the same element in at least two different oxidation states. Condition: The element must be in an intermediate oxidation state. Examples: $2H_2O_2 \rightarrow 2H_2O + O_2$ (O: $-1 \rightarrow -2$ and $0$) $3Cl_2 + 6NaOH \rightarrow 5NaCl + NaClO_3 + 3H_2O$ (Cl: $0 \rightarrow -1$ and $+5$) Comproportionation (Synproportionation) Reaction: The reverse of disproportionation. Two compounds containing the same element in different oxidation states react to form a product in which the element is in an intermediate oxidation state. Examples: $Ag^{2+} + Ag \rightarrow 2Ag^+$ (Ag: $+2$ and $0 \rightarrow +1$) $SO_2 + 2H_2S \rightarrow 3S + 2H_2O$ (S: $+4$ and $-2 \rightarrow 0$) 3. Oxidation Number: Rules and Applications 3.1. Rules for Assigning Oxidation Numbers (O.N.) Free Elements: The O.N. of an atom in its elemental form (uncombined state) is zero. (e.g., $H_2, O_2, S_8, Na, Fe$). Monoatomic Ions: The O.N. of a monoatomic ion is equal to its charge. (e.g., $Na^+ = +1$, $Cl^- = -1$, $Ca^{2+} = +2$). Group 1 Elements (Alkali Metals): Always $+1$ in compounds. Group 2 Elements (Alkaline Earth Metals): Always $+2$ in compounds. Fluorine: Always $-1$ in compounds, as it is the most electronegative element. Hydrogen: $+1$ when bonded to non-metals (e.g., $HCl, H_2O, NH_3$). $-1$ when bonded to metals (metal hydrides, e.g., $NaH, CaH_2$). Oxygen: Usually $-2$ in compounds. Exceptions: Peroxides (e.g., $H_2O_2, Na_2O_2$): $-1$ Superoxides (e.g., $KO_2, RbO_2$): $-1/2$ Ozonides (e.g., $KO_3$): $-1/3$ When bonded to fluorine (e.g., $OF_2$): $+2$; ($O_2F_2$): $+1$. Sum of O.N.: In a neutral molecule, the sum of O.N. of all atoms is zero. In a polyatomic ion, the sum of O.N. of all atoms is equal to the charge on the ion. 3.2. Applications and Nuances of O.N. Stock Notation: Used for compounds of metals with variable oxidation states to indicate the O.N. of the metal using Roman numerals in parentheses (e.g., $Fe(II)Cl_2$ for $FeCl_2$, $Fe(III)Cl_3$ for $FeCl_3$). Fractional Oxidation Numbers: These arise when an element exists in different oxidation states within the same compound, and the calculated O.N. is an average. Example: $Fe_3O_4$. Calculated O.N. for $Fe$ is $+8/3$. This is an average, as $Fe_3O_4$ is a mixed oxide ($FeO \cdot Fe_2O_3$) containing $Fe^{2+}$ and $Fe^{3+}$ ions. Example: $C_3O_2$ (carbon suboxide). Average O.N. for $C$ is $+4/3$. Structure $O=C=C=C=O$ reveals O.N.s of $0, +2, +2$ for carbons. Variable Oxidation States: Many d-block and p-block elements exhibit variable O.N. (e.g., $Mn$ from $-3$ to $+7$, $Cr$ from $0$ to $+6$, $S$ from $-2$ to $+6$, $N$ from $-3$ to $+5$). 3.3. Identifying Redox vs. Non-Redox Reactions A reaction is a redox reaction if there is a change in the oxidation number of at least one element from reactants to products. Non-redox reactions, such as acid-base reactions, precipitation reactions, and some double displacement/metathesis reactions, do not involve a change in O.N. for any element. Redox in Everyday Processes: Combustion, respiration, photosynthesis, corrosion (rusting), bleaching, photography, operation of batteries and fuel cells. 4. Balancing Redox Equations 4.1. General Principles Total increase in oxidation number must equal total decrease in oxidation number. Number of electrons lost in oxidation must equal number of electrons gained in reduction. Mass balance (equal number of atoms of each element on both sides). Charge balance (total charge on both sides must be equal). 4.2. Oxidation Number Method Write the unbalanced equation: Identify reactants and products. Assign O.N.s: Determine the oxidation number for all atoms in the reaction. Identify changes: Note which atoms undergo changes in O.N. (oxidised and reduced species). Calculate total change: Determine the total increase in O.N. for the oxidised species and the total decrease in O.N. for the reduced species per formula unit. Equalise O.N. changes: Multiply the oxidised and reduced species by appropriate integers to make the total increase in O.N. equal to the total decrease in O.N. Balance other atoms: Balance all other atoms (except H and O) by inspection. Balance Oxygen (O): Add $H_2O$ molecules to the side deficient in oxygen. Balance Hydrogen (H): In acidic medium: Add $H^+$ ions to the side deficient in hydrogen. In basic medium: Add $H_2O$ molecules to the side deficient in hydrogen, and an equal number of $OH^-$ ions to the opposite side. (Alternatively, balance as if in acidic medium, then add $OH^-$ to both sides equal to the number of $H^+$ ions to neutralise them to $H_2O$). Verify: Check mass and charge balance. 4.3. Ion-Electron (Half-Reaction) Method Write the unbalanced ionic equation: Split the reaction into two half-reactions (oxidation and reduction). Balance atoms (excluding O and H): Balance all atoms other than oxygen and hydrogen in each half-reaction. Balance Oxygen (O): Add $H_2O$ molecules to the side deficient in oxygen. Balance Hydrogen (H): In acidic medium: Add $H^+$ ions to the side deficient in hydrogen. In basic medium: Add $H_2O$ to the side deficient in H, and an equal number of $OH^-$ ions to the opposite side. (Or, balance with $H^+$ as in acidic, then add $OH^-$ to both sides to cancel $H^+$ ions, forming $H_2O$). Balance Charge: Add electrons ($e^-$) to the more positive side of each half-reaction to balance the charge. Equalise Electrons: Multiply each half-reaction by an appropriate integer so that the number of electrons lost in the oxidation half-reaction equals the number of electrons gained in the reduction half-reaction. Combine Half-Reactions: Add the two balanced half-reactions together and cancel out any common species (electrons, $H_2O, H^+/OH^-$). Verify: Check mass and charge balance. Practice Patterns: Oxidations involving $Cr_2O_7^{2-}$ (to $Cr^{3+}$), $MnO_4^-$ (to $Mn^{2+}$ in acid, $MnO_2$ in neutral, $MnO_4^{2-}$ in strong base), $SO_3^{2-}$ to $SO_4^{2-}$, $NO_2^-$ to $NO_3^-$, $Fe^{2+}$ to $Fe^{3+}$. 5. Types of Redox Reactions (by O.N. change) Combination Reactions: Two or more substances combine to form a single new substance. $A + B \rightarrow AB$. Often redox (e.g., $C + O_2 \rightarrow CO_2$). Decomposition Reactions: A single compound breaks down into two or more simpler substances. $AB \rightarrow A + B$. Can be redox (e.g., $2KClO_3 \rightarrow 2KCl + 3O_2$). Displacement Reactions: An atom or ion in a compound is replaced by an atom or ion of another element. Metal Displacement: A more reactive metal displaces a less reactive metal from its salt solution (e.g., $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$). Non-metal Displacement: A more reactive non-metal (usually halogen) displaces a less reactive non-metal (e.g., $Cl_2(g) + 2KI(aq) \rightarrow 2KCl(aq) + I_2(aq)$). Hydrogen Displacement: Active metals react with acids or water to displace hydrogen (e.g., $Zn + 2HCl \rightarrow ZnCl_2 + H_2$). Disproportionation & Comproportionation: (See Section 2.2). 6. Competitive Electron Transfer Reactions & Activity Series Redox Series / Activity Series: A list of metals (or non-metals) arranged in order of their decreasing tendency to lose (or gain) electrons. More reactive metals (higher in the series) have a greater tendency to lose electrons and act as stronger reducing agents. They can displace less reactive metals from their salt solutions. Example: $Zn$ is more reactive than $Cu$. So, $Zn$ displaces $Cu$ from $CuSO_4$ solution ($Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$). This concept is qualitatively linked to standard electrode potentials ($E^\circ$), where a more negative $E^\circ$ for reduction indicates a stronger reducing agent (tendency to get oxidised). 7. Redox Titrations 7.1. Basic Concept & Indicators Redox Titration: A volumetric analysis technique based on a redox reaction between the analyte and the titrant. The endpoint is detected when the reaction is complete. Indicators: Self-Indicators: The titrant itself changes color upon completion of the reaction (e.g., $KMnO_4$, which is purple and becomes colorless when reduced to $Mn^{2+}$). External Indicators: A separate substance added to the reaction mixture that changes color at or near the equivalence point (e.g., starch for $I_2$, diphenylamine for $Fe^{2+}$ titration with $Cr_2O_7^{2-}$). Primary Standard: A substance of high purity, known composition, and stability, used to prepare solutions of accurately known concentration. (e.g., oxalic acid, Mohr's salt for redox). Secondary Standard: A solution whose concentration is determined by titration against a primary standard (e.g., $KMnO_4$, $NaOH$). 7.2. Common Redox Titration Systems Permanganometry ($KMnO_4$): Strong oxidising agent. Usually performed in acidic medium ($MnO_4^- \rightarrow Mn^{2+}$). Self-indicator (purple $MnO_4^-$ to colorless $Mn^{2+}$). Typical reactions: Oxalic acid, $Fe^{2+}$ ions, $C_2O_4^{2-}$, $H_2S$, $SO_2$. $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$ ($n$-factor = 5). Dichromatometry ($K_2Cr_2O_7$): Strong oxidising agent. Used in acidic medium ($Cr_2O_7^{2-} \rightarrow 2Cr^{3+}$). Color change: Orange ($Cr_2O_7^{2-}$) to Green ($Cr^{3+}$). External indicator like diphenylamine sulfonate is often used. Typical reactions: $Fe^{2+}$ ions, $I^-$. $Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O$ ($n$-factor = 6). Iodometry & Iodimetry: Involve $I_2/I^-$ redox couple. Iodimetry: Direct titration with a standard $I_2$ solution (oxidant) to estimate reducing agents (e.g., thiosulphate, arsenite). Starch is used as an indicator (blue-black with $I_2$). Iodometry: Indirect titration. An oxidising agent (analyte) is made to react with excess $KI$ to liberate $I_2$. The liberated $I_2$ is then titrated with a standard solution of sodium thiosulphate ($Na_2S_2O_3$). $Oxidant + 2I^- \rightarrow Oxidised~Product + I_2$ $I_2 + 2S_2O_3^{2-} \rightarrow 2I^- + S_4O_6^{2-}$ Cerimetry ($Ce^{4+}/Ce^{3+}$): $Ce(SO_4)_2$ or $Ce(NO_3)_4$ are strong oxidants. $Ce^{4+}$ (yellow) reduces to $Ce^{3+}$ (colorless). Ferroin is a common indicator. ($Ce^{4+} + e^- \rightarrow Ce^{3+}$). 7.3. Equivalent Concept in Redox Titrations Equivalent Mass ($E$): Molar Mass / $n$-factor. $n$-factor (Valency Factor): The number of moles of electrons transferred per mole of the substance in the redox reaction. For an oxidising agent: $n$-factor = total decrease in O.N. per molecule/ion. For a reducing agent: $n$-factor = total increase in O.N. per molecule/ion. Normality ($N$): $N = Molarity (M) \times n$-factor. At the equivalence point of a titration: $N_1V_1 = N_2V_2$ or $(M_1 \times n_1)V_1 = (M_2 \times n_2)V_2$. 8. JEE-Focused Micro-Keywords & Advanced Concepts Disproportionation Conditions: An element can disproportionate if it exists in at least three oxidation states and the intermediate O.N. is less stable than the two extreme O.N.s it forms. Strength of Oxidising/Reducing Agents (Qualitative Order): Strong Oxidants: $F_2 > O_3 > MnO_4^- > Cr_2O_7^{2-} > Cl_2 > H_2O_2$. Strong Reductants: Alkali metals > Alkaline earth metals > $Al > Zn > Fe$. The stronger the oxidising agent, the weaker its conjugate reducing agent. Internal Redox in Same Molecule/Ion: Some compounds contain an element in an intermediate O.N. which can act as both an oxidising and a reducing agent. Example: Nitrous acid ($HNO_2$, N is $+3$). Can be oxidised to $HNO_3$ (N is $+5$) or reduced to $NO$ (N is $+2$), $N_2O$ (N is $+1$), $N_2$ (N is $0$), $NH_3$ (N is $-3$). Example: Hypochlorites ($ClO^-$, Cl is $+1$). Can be reduced to $Cl^-$ (Cl is $-1$) or disproportionate to $Cl^-$ and $ClO_3^-$ (Cl is $+5$). Redox in Inorganic Qualitative Analysis: Many tests involve characteristic redox reactions. $Fe^{2+}$ (green) oxidised to $Fe^{3+}$ (yellow/brown) by $H_2O_2$ in acidic medium. $Sn^{2+}$ (reducing agent) reducing $HgCl_2$ to $Hg_2Cl_2$ (white ppt) and then to $Hg$ (black ppt). $S_2O_3^{2-}$ (thiosulphate) as a reducing agent in iodometric titrations. Electrode Processes: In an electrochemical cell, oxidation always occurs at the anode, and reduction always occurs at the cathode. In galvanic cells, anode is negative and cathode is positive. In electrolytic cells, anode is positive and cathode is negative. EMF and Nernst Equation: These are crucial for quantitative analysis of redox processes and are covered in detail in Electrochemistry, building upon the qualitative understanding of redox.