Comprehensive CBSE Class 12 Chemistry Master Study Guide

Solutions and Colligative Properties

• Azeotropic Mixtures and Raoult's Law Deviations:

  • Azeotropes: Binary mixtures having the same composition in both liquid and vapour phases and boiling at a constant temperature. Components of an azeotrope cannot be separated by fractional distillation.
  • Negative Deviation: Occurs when solute-solvent intermolecular attractive forces are stronger than solute-solute and solvent-solvent interactions ($A-B > A-A, B-B$). Example: A mixture of phenol and aniline exhibits strong intermolecular hydrogen bonding between the phenolic hydroxyl proton and the nitrogen lone pair of aniline.
  • Boiling Point Behavior: Solutions showing negative deviation form maximum boiling azeotropes, causing an increase in the overall boiling point of the solution upon mixing.
  • Positive Deviation: Occurs when solute-solvent intermolecular interactions are weaker than pure component interactions ($A-B < A-A, B-B$). Example: A mixture of ethanol and water. The boiling point of a positive deviation azeotrope is lower than that of either pure component.
  • Vapor Pressure & Boiling Point Relationship: Addition of a non-volatile solute lowers vapor pressure and elevates boiling point ($\Delta T_b = i \times K_b \times m$). Addition of a volatile solute with high vapor pressure (e.g., methanol to water) increases total vapor pressure and lowers the boiling point.

• Henry's Law and Gas Solubility:

  • Statement: The partial pressure of a gas in the vapour phase ($p$) is directly proportional to the mole fraction of the gas ($x$) in the solution: $p = K_H \times x$, where $K_H$ is Henry's law constant.
  • Physiological Applications:
  • Bends (Decompression Sickness): Scuba divers breathe air at high pressure underwater, increasing nitrogen solubility in blood. Upon ascending rapidly, pressure decreases, releasing dissolved nitrogen as gas bubbles in blood vessels, blocking capillaries and causing painful, dangerous conditions called bends. Diluting breathing gas with helium mitigates this.
  • Anoxia: At high altitudes, the partial pressure of oxygen is lower than at sea level. Low atmospheric pressure leads to lower concentrations of dissolved oxygen in the blood and tissues of climbers, causing weakness and inability to think clearly, a condition known as anoxia.
  • Calculation Example: Mole fraction of $\text{CO}_2$ in water at $298\,K$ under $760\,mm\,Hg$ with $K_H = 1.25 \times 10^6\,mm\,Hg$:
  • $x = \frac{p}{K_H} = \frac{760\,mm\,Hg}{1.25 \times 10^6\,mm\,Hg} = 6.08 \times 10^{-4}$.

• Colligative Properties & van 't Hoff Factor ($i$):

  • Freezing Point Depression: $\Delta T_f = i \times K_f \times m$
  • Boiling Point Elevation: $\Delta T_b = i \times K_b \times m$
  • Osmotic Pressure: $\Pi = i \times C \times R \times T = i \times \left(\frac{w}{M \times V}\right) \times R \times T$
  • van 't Hoff Factor Formula: $i = 1 + (n - 1)\alpha$ for dissociation; $i = 1 + \left(\frac{1}{n} - 1\right)\alpha$ for association.
  • Isotonic, Hypertonic, and Hypotonic Solutions:
  • Isotonic: Two solutions having equal osmotic pressure across a semipermeable membrane. Normal saline solution ($0.9\%\,\text{w/v}\text{ NaCl}$) is isotonic with fluid inside human red blood cells.
  • Hypertonic: A solution with higher osmotic pressure relative to cellular fluid. RBCs placed in hypertonic solution shrink (crenate).
  • Hypotonic: A solution with lower osmotic pressure relative to cellular fluid. RBCs placed in hypotonic solution swell and burst (hemolysis). Swelling of raisins in water is an example of osmosis in hypotonic media.
  • Advantages of Osmotic Pressure Method: Preferred for determining molar masses of biomolecules and polymers because measurements are made at room temperature and osmotic pressure magnitude is significantly larger and easily readable compared to minute changes in freezing or boiling points.

• Ideal vs Non-Ideal Solutions:

  • Ideal Solutions: Obey Raoult's law across the entire concentration range; $\Delta_{mix} H = 0$; $\Delta_{mix} V = 0$; solute-solvent interactions equal solute-solute and solvent-solvent interactions.
  • Non-Ideal Solutions: Do not obey Raoult's law; $\Delta_{mix} H \neq 0$; $\Delta_{mix} V \neq 0$; show positive or negative deviations.

Electrochemistry and Electrochemical Cells

• Electromotive Force (EMF) and Nernst Equation:

  • Nernst Equation Formulation: For a cell reaction $aA + bB \rightarrow cC + dD$:
  • $E_{cell} = E^\circ_{cell} - \frac{2.303 \times R \times T}{n \times F} \log\left(\frac{[C]^c [D]^d}{[A]^a [B]^b}\right)$
  • At $298\,K$: $E_{cell} = E^\circ_{cell} - \frac{0.0591}{n} \log\left(\frac{\text{[Products]}}{\text{[Reactants]}}\right)$
  • Standard Free Energy Change: $\Delta_r G^\circ = -n \times F \times E^\circ_{cell}$
  • Equilibrium Constant Relationship: $\Delta_r G^\circ = -2.303 \times R \times T \log(K_c)$, hence $E^\circ_{cell} = \frac{0.0591}{n} \log(K_c)$. $K_c$ is related to $E^\circ_{cell}$ and not $E_{cell}$ because at equilibrium $E_{cell} = 0\,V$.

• Conductivity and Molar Conductivity:

  • Molar Conductivity Formula: $\Lambda_m = \frac{\kappa \times 1000}{c}$, where $\kappa$ is conductivity in $\text{S}\,\text{cm}^{-1}$ and $c$ is molar concentration in $\text{mol}\,\text{L}^{-1}$.
  • Kohlrausch's Law of Independent Migration of Ions: The limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation:
  • $\Lambda_m^\circ = A \times \lambda_ +^\circ + B \times \lambda_ -^\circ$
  • Extrapolation Behavior: For strong electrolytes, $\Lambda_m$ decreases linearly with $\sqrt{c}$ (Debye-Hückel-Onsager equation) allowing direct extrapolation to zero concentration. For weak electrolytes, $\Lambda_m$ increases steeply at low concentrations due to dissociation, making direct graphical extrapolation to $c = 0$ impossible.
  • Measurement Restrictions: Direct Current (DC) is not used to measure ionic solution resistance because DC alters the chemical composition of the solution via electrolysis. Alternating Current (AC) is used instead.

• Batteries, Fuel Cells, and Corrosion:

  • Primary vs Secondary Batteries: Primary batteries (e.g., Dry cell, Mercury cell) cannot be recharged as reaction occurs only once. Secondary batteries (e.g., Lead storage battery) are rechargeable via reverse current application.
  • Dry Cell (Leclanché Cell): Anode: $\text{Zn}(s) \rightarrow \text{Zn}^{2+} + 2e^-$; Cathode: $2\text{MnO}_2 + 2\text{NH}_4^+ + 2e^- \rightarrow \text{Mn}_2\text{O}_3 + 2\text{NH}_3 + \text{H}_2\text{O}$. Electrolyte: paste of $\text{NH}_4\text{Cl}$ and $\text{ZnCl}_2$.
  • Mercury Cell: Anode: $\text{Zn(Hg)} + 2\text{OH}^- \rightarrow \text{ZnO}(s) + \text{H}_2\text{O} + 2e^-$; Cathode: $\text{HgO}(s) + \text{H}_2\text{O} + 2e^- \rightarrow \text{Hg}(l) + 2\text{OH}^-$. Cell potential remains constant ($\sim 1.35\,V$) throughout its life because the overall cell reaction does not involve any ion in solution whose concentration can change.
  • Lead Storage Battery (Discharge vs Charge):
  • Discharge Anode: $\text{Pb}(s) + \text{SO}_4^{2-} \rightarrow \text{PbSO}_4(s) + 2e^-$
  • Discharge Cathode: $\text{PbO}_2(s) + \text{SO}_4^{2-} + 4\text{H}^+ + 2e^- \rightarrow \text{PbSO}_4(s) + 2\text{H}_2\text{O}$
  • Recharging Overall Reaction: $2\text{PbSO}_4(s) + 2\text{H}_2\text{O}(l) \rightarrow \text{Pb}(s) + \text{PbO}_2(s) + 2\text{H}_2\text{SO}_4(aq)$
  • Hydrogen-Oxygen Fuel Cell: Used in Apollo space programs. Anode: $2\text{H}_2(g) + 4\text{OH}^-(aq) \rightarrow 4\text{H}_2\text{O}(l) + 4e^-$; Cathode: $\text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4e^- \rightarrow 4\text{OH}^-(aq)$. Advantages: high efficiency ($\sim 70\%$), continuous power supply, pollution-free produce (water vapor used as drinking supply).
  • Cathodic Protection Against Corrosion: Magnesium blocks are connected to underground iron pipelines because Mg has a more negative standard electrode potential than Fe, acting as a sacrificial anode and preventing Fe oxidation.

• Faraday's Laws of Electrolysis:

  • First Law: Mass of substance deposited/liberated at an electrode is proportional to quantity of electricity: $m = Z \times Q = Z \times I \times t$
  • Second Law: When the same quantity of electricity passes through different electrolytes, masses of substances liberated are proportional to their chemical equivalent weights ($\text{E}_1 / \text{E}_2 = m_1 / m_2$).
  • Electrolysis Products:
  • Dilute $\text{NaCl}(aq)$ with Pt electrodes: $\text{H}_2(g)$ at cathode, $\text{O}_2(g)$ at anode.
  • Concentrated $\text{CuCl}_2(aq)$ with Pt electrodes: $\text{Cu}(s)$ at cathode, $\text{Cl}_2(g)$ at anode.
  • Concentrated $\text{H}_2\text{SO}_4$ at high current density: Peroxodisulfate ion ($\text{S}_2\text{O}_8^{2-}$) forms at anode.

Chemical Kinetics and Reaction Dynamics

• Order vs Molecularity of Reaction:

  • Order: Sum of powers of concentration terms in rate law expression. Determined experimentally; can be zero, fractional, or integer; applicable to elementary and complex reactions.
  • Molecularity: Number of reacting species taking part in an elementary reaction that collide simultaneously. Always a positive integer; cannot be zero or fractional; meaningless for complex reactions.
  • Pseudo-First Order Reactions: A bimolecular reaction made kinetically first-order by presence of one reactant in large excess (e.g., hydrolysis of ethyl acetate in excess water).

• First Order Kinetics Equations:

  • Integrated Rate Law: $k = \frac{2.303}{t} \log\left(\frac{[A]_0}{[A]}\right)$
  • Half-Life: $t_{1/2} = \frac{0.693}{k}$
  • Thermal Decomposition in Gas Phase: For $A(g) \rightarrow B(g) + C(g)$:
  • $k = \frac{2.303}{t} \log\left(\frac{P_0}{2P_0 - P_t}\right)$

• Temperature Dependence and Activation Energy ($E_a$):

  • Activation Energy: Minimum excess energy required by reactant molecules to form an activated complex and result in chemical transformation.
  • Arrhenius Equation: $k = A e^{-E_a/RT} \implies \log\left(\frac{k_2}{k_1}\right) = \frac{E_a}{2.303 \times R} \left(\frac{T_2 - T_1}{T_1 T_2}\right)$
  • Fraction of Molecules: Fraction of molecules having energy $\ge E_a$ is represented by $e^{-E_a/RT}$.

d-Block and f-Block Transition Elements

• General Properties of d-Block Elements:

  • Irregular Trend in $E^\circ_{M^{2+}/M}$: Due to irregular variations in sublimation enthalpies, ionization enthalpies ($I_1 + I_2$), and hydration enthalpies across 3d transition metals.
  • Positive Value for $E^\circ_{\text{Cu}^{2+}/\text{Cu}}$ ($+0.34\,V$): High enthalpy of atomization and high ionization enthalpy of copper are not compensated by its hydration enthalpy.
  • Highly Negative $E^\circ_{\text{Mn}^{2+}/\text{Mn}}$ ($-0.91\,V$): Due to extra stability of half-filled $d^5$ electronic configuration in $\text{Mn}^{2+}$.
  • Non-Transition Classification: Zinc, Cadmium, and Mercury are not considered transition elements because they have completely filled $d^{10}$ orbitals in their ground state as well as in common oxidation states.
  • Catalytic Activity: Transition metals act as excellent catalysts due to variable oxidation states, ability to form reaction intermediates, and provision of large surface area for reactant adsorption.
  • Melting Points: Chromium has a higher melting point than Manganese because Chromium ($3d^5 4s^1$) has 6 unpaired electrons available for strong metallic bonding, whereas Manganese ($3d^5 4s^2$) has a stable $s^2$ pair resulting in weaker metallic binding.

• Potassium Permanganate ($\text{KMnO}_4$) & Dichromate ($\text{K}_2\text{Cr}_2\text{O}_7$) Reactions:

  • Dichromate Preparation:
  • $4\text{FeCr}_2\text{O}_4 + 8\text{Na}_2\text{CO}_3 + 7\text{O}_2 \rightarrow 8\text{Na}_2\text{CrO}_4 + 2\text{Fe}_2\text{O}_3 + 8\text{CO}_2$
  • $2\text{Na}_2\text{CrO}_4 + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{Cr}_2\text{O}_7 + \text{Na}_2\text{SO}_4 + \text{H}_2\text{O}$
  • $\text{Na}_2\text{Cr}_2\text{O}_7 + 2\text{KCl} \rightarrow \text{K}_2\text{Cr}_2\text{O}_7 + 2\text{NaCl}$
  • Permanganate Preparation Sequence: Dark brown $\text{MnO}_2$ fused with KOH and $\text{O}_2$ yields dark green manganate $\text{K}_2\text{MnO}_4$, which upon electrolytic oxidation in alkaline solution yields dark purple permanganate $\text{KMnO}_4$.
  • Disproportionation of Manganate: In acidic medium, green manganate ion undergoes disproportionation:
  • $3\text{MnO}_4^{2-} + 4\text{H}^+ \rightarrow 2\text{MnO}_4^- + \text{MnO}_2 + 2\text{H}_2\text{O}$
  • Oxidizing Ionic Equations for $\text{MnO}_4^-$:
  • Acidic medium with $\text{Fe}^{2+}$: $5\text{Fe}^{2+} + \text{MnO}_4^- + 8\text{H}^+ \rightarrow 5\text{Fe}^{3+} + \text{Mn}^{2+} + 4\text{H}_2\text{O}$
  • Acidic medium with $\text{I}^-$: $10\text{I}^- + 2\text{MnO}_4^- + 16\text{H}^+ \rightarrow 5\text{I}_2 + 2\text{Mn}^{2+} + 8\text{H}_2\text{O}$
  • Neutral/Alkaline medium with $\text{I}^-$: $\text{I}^- + 2\text{MnO}_4^- + \text{H}_2\text{O} \rightarrow \text{IO}_3^- + 2\text{MnO}_2 + 2\text{OH}^-$
  • Acidic medium with Oxalate: $5\text{C}_2\text{O}_4^{2-} + 2\text{MnO}_4^- + 16\text{H}^+ \rightarrow 10\text{CO}_2 + 2\text{Mn}^{2+} + 8\text{H}_2\text{O}$

• f-Block Elements (Lanthanoids and Actinoids):

  • Lanthanoid Contraction: Continuous decrease in atomic and ionic radii of lanthanoid elements with increasing atomic number due to poor shielding effect of $4f$ electrons.
  • Consequence: Zr ($4d$) and Hf ($5d$) have almost identical atomic radii ($160\,pm$ and $159\,pm$), making their chemical separation extremely difficult.
  • Oxidation States in Lanthanoids: Cerium exhibits +4 oxidation state ($\text{Ce}^{4+}$ has noble gas configuration $[\text{Xe}]$, making it a strong oxidant). Europium exhibits +2 oxidation state ($\text{Eu}^{2+}$ has stable $[\text{Xe}] 4f^7$ configuration).
  • Actinoids: Actinoid chemistry is complex due to radioactivity and comparable energy levels of $5f$, $6d$, and $7s$ orbitals, leading to a wider range of oxidation states.

Coordination Chemistry and Bonding Theories

• Terminology and Nomenclature:

  • IUPAC Examples:
  • $[\text{Ag}(\text{NH}_3)_2][\text{Ag}(\text{CN})_2]$: diamminesilver(I) dicyanidoargentate(I)
  • $\text{K}_3[\text{Fe}(\text{C}_2\text{O}_4)_3]$: potassium trioxalatoferrate(III)
  • $[\text{PtCl}_2(\text{en})_2]\text{SO}_4$: dichloridobis(ethane-1,2-diamine)platinum(IV) sulfate
  • $(\text{NH}_4)_2[\text{CoF}_4]$: ammonium tetrafluoridocobaltate(II)
  • $[\text{Cr}(\text{NH}_3)_4(\text{ONO})\text{Cl}]\text{NO}_3$: tetraamminechloridonitritochromium(III) nitrate
  • Ambidentate Ligand: A ligand possessing two different donor atoms that can coordinate to a central metal ion through either atom (e.g., $\text{NO}_2^-$ via N or O; $\text{SCN}^-$ via S or N).
  • Chelate Effect: Increased stability of coordination complexes formed by poly- or bidentate ligands compared to unidentate analogs due to entropy gain upon chelate ring formation (e.g., $[\text{Co}(\text{en})_3]^{3+}$).

• Valence Bond Theory (VBT) & Crystal Field Theory (CFT):

  • Octahedral Splitting ($\Delta_o$): In an octahedral crystal field, ligands approach along the cartesian axes. $d_{x^2-y^2}$ and $d_{z^2}$ orbitals ($e_g$ set) experience greater electrostatic repulsion and rise in energy, while $d_{xy}, d_{yz}, d_{xz}$ orbitals ($t_{2g}$ set) decrease in energy.
  • Spectrochemical Series & High/Low Spin Configurations:
  • $[\text{Co}(\text{NH}_3)_6]^{3+}$: Strong field ligand $\text{NH}_3$, $\Delta_o > P$, inner orbital complex ($d^2sp^3$), low spin, diamagnetic ($t_{2g}^6 e_g^0$).
  • $[\text{CoF}_6]^{3-}$: Weak field ligand $\text{F}^-$, $\Delta_o < P$, outer orbital complex ($sp^3d^2$), high spin, paramagnetic ($t_{2g}^4 e_g^2$) with 4 unpaired electrons.
  • $[\text{NiCl}_4]^{2-}$: Tetrahedral, $sp^3$ hybridised, paramagnetic (2 unpaired electrons).
  • $[\text{Ni}(\text{CO})_4]$: Tetrahedral, $sp^3$ hybridised, strong field CO causes pairing, diamagnetic ($d^{10}$ configuration).
  • $[\text{Fe}(\text{CN})_6]^{3-}$: $d^2sp^3$ hybridised, inner orbital, paramagnetic (1 unpaired electron).

Haloalkanes and Haloarenes

• Reactivity Trends ($S_N1$ vs $S_N2$):

  • $S_N1$ Reactivity: Governed by carbocation stability ($3^\circ > 2^\circ > 1^\circ$). Tertiary alkyl halides like 2-bromo-2-methylpropane ($(\text{CH}_3)_3\text{CBr}$) undergo $S_N1$ rapidly with racemisation.
  • $S_N2$ Reactivity: Governed by steric hindrance ($1^\circ > 2^\circ > 3^\circ$). Primary alkyl halides undergo $S_N2$ rapidly with inversion of configuration (Walden inversion).
  • Low Reactivity of Haloarenes: Less reactive towards nucleophilic substitution due to resonance stabilization (partial double bond character of C-X bond), $sp^2$ hybridisation of C atom, instability of phenyl cation, and electron repulsion.

• Directing Effects and Reactions:

  • Chlorine Directing Effect: Chlorine is electron-withdrawing via $-I$ effect, yet ortho-, para-directing in electrophilic substitution because $\text{+R}$ resonance effect stabilizes carbocation at ortho and para positions.
  • Wurtz-Fittig and Finkelstein Reactions:
  • Finkelstein: $\text{CH}_3\text{Cl} + \text{NaI} \xrightarrow{\text{dry acetone}} \text{CH}_3\text{I} + \text{NaCl}\downarrow$
  • Wurtz: $2\text{C}_2\text{H}_5\text{Cl} + 2\text{Na} \xrightarrow{\text{dry ether}} \text{C}_4\text{H}_{10} + 2\text{NaCl}$
  • Phosgene Formation: Chloroform ($\text{CHCl}_3$) slowly oxidized by air in presence of light to form toxic phosgene gas ($\text{COCl}_2$):
  • $2\text{CHCl}_3 + \text{O}_2 \xrightarrow{\text{sunlight}} 2\text{COCl}_2 + 2\text{HCl}$
  • Stored in closed, dark-coloured bottles filled completely to prevent air exposure.

Alcohols, Phenols, and Ethers

• Acidity Trends:

  • Acidity of Phenols vs Alcohols: Phenols are significantly more acidic than alcohols because phenoxide ion is resonance stabilized by delocalization of negative charge over the aromatic ring, whereas alkoxide ion is not.
  • Substituent Effects on Phenol Acidity:
  • Electron-withdrawing groups (e.g., $\text{-NO}_2$) increase acidity by stabilizing phenoxide ion (ortho-nitrophenol > phenol).
  • Electron-releasing groups (e.g., $\text{-CH}_3$, $\text{-OCH}_3$) decrease acidity (phenol > 4-methylphenol).

• Distinguishing Tests and Named Reactions:

  • Lucas Test: Distinguishes $1^\circ, 2^\circ, 3^\circ$ alcohols using anhydrous $\text{ZnCl}_2 + \text{conc. HCl}$. $3^\circ$ alcohols give immediate turbidity; $2^\circ$ give turbidity in 5 minutes; $1^\circ$ do not give turbidity at room temperature.
  • Reimer-Tiemann Reaction: Phenol reacts with $\text{CHCl}_3$ and aqueous NaOH to form salicylaldehyde (2-hydroxybenzaldehyde).
  • Kolbe's Reaction: Phenol reacts with NaOH to give phenoxide, which reacts with $\text{CO}_2$ followed by acidification to yield salicylic acid (2-hydroxybenzoic acid).

• Ethers & Cleavage Reactions:

  • Williamson Ether Synthesis: $\text{R-X} + \text{R'-ONa} \rightarrow \text{R-O-R'} + \text{NaX}$ via $S_N2$ mechanism. Alkyl halide must be primary; tertiary alkyl halides undergo elimination to form alkenes.
  • Cleavage by HI: Unsymmetrical ether $\text{R-O-R'}$ with HI cleaves to give alkyl iodide from smaller alkyl group. If one group is tertiary, iodide forms at tertiary carbon via stable $3^\circ$ carbocation. Anisole reacts with HI to form phenol and methyl iodide ($\text{CH}_3\text{I}$) due to high stability of $sp^2$ aryl-oxygen bond.

Aldehydes, Ketones, and Carboxylic Acids

• Reactivity Order Towards Nucleophilic Addition:

  • Reactivity order towards HCN: Methanal > Ethanal > Propanone > Di-tert-butyl ketone.
  • Aldehydes are more reactive than ketones due to steric hindrance in ketones and electron-donating inductive effect of two alkyl groups reducing carbonyl carbon electrophilicity.

• Acidity of Carboxylic Acids & High Boiling Points:

  • Higher Boiling Points: Carboxylic acids have higher boiling points than alcohols of comparable molecular mass because they form stable hydrogen-bonded cyclic dimers.
  • Acidity Trend: $\text{CF}_3\text{COOH} > \text{O}_2\text{N-CH}_2\text{COOH} > \text{FCOCH}_2\text{COOH} > \text{HCOOH} > \text{C}_6\text{H}_5\text{COOH} > \text{CH}_3\text{COOH}$.
  • $\alpha$-Hydrogens: Aldehyde and ketone $\alpha$-hydrogens are acidic due to strong electron-withdrawing carbonyl group and resonance stabilization of enolate ion.

• Named Reactions and Conversions:

  • Etard Reaction: Toluene to benzaldehyde using $\text{CrO}_2\text{Cl}_2$ in $\text{CS}_2$ followed by hydrolysis.
  • Rosenmund Reduction: Benzoyl chloride to benzaldehyde using $\text{H}_2 / \text{Pd-BaSO}_4$
  • Stephen Reaction: Nitrile to aldehyde using $\text{SnCl}_2 + \text{HCl}$ followed by $\text{H}_3\text{O}^+$.
  • Cannizzaro Reaction: Aldehydes without $\alpha$-hydrogen (e.g., HCHO, $\text{C}_6\text{H}_5\text{CHO}$) undergo self-oxidation-reduction with conc. NaOH to yield alcohol and carboxylate salt.
  • Wolff-Kishner Reduction: Carbonyl group to $\text{-CH}_2-$ using $\text{NH}_2\text{NH}_2 / \text{KOH, ethylene glycol, } \Delta$
  • Clemmensen Reduction: Carbonyl group to $\text{-CH}_2-$ using $\text{Zn(Hg)} / \text{conc. HCl}$
  • Hell-Volhard-Zelinsky (HVZ) Reaction: Carboxylic acid with $\alpha$--hydrogen reacts with $\text{X}_2 / \text{Red Phosphorus}$ to form $\alpha$-halocarboxylic acid.

Organic Compounds Containing Nitrogen (Amines & Diazonium Salts)

• Basicity Order of Amines:
  • Gas Phase: $3^\circ > 2^\circ > 1^\circ > \text{NH}_3\n - **Aqueous Phase (Methyl Substituted):** $2^\circ > 1^\circ > 3^\circ > \text{NH}_3
  • $\left((\text{CH}_3)_2\text{NH} > \text{CH}_3\text{NH}_2 > (\text{CH}_3)_3\text{N} > \text{NH}_3\right)$
  • Aqueous Phase (Ethyl Substituted): $2^\circ > 3^\circ > 1^\circ > \text{NH}_3\n - \left((\text{C}_2\text{H}_5)_2\text{NH} > (\text{C}_2\text{H}_5)_3\text{N} > \text{C}_2\text{H}_5\text{NH}_2 > \text{NH}_3\right)\n - **Aromatic Amines:** Weaker bases than \text{NH}_3 because lone pair on N is delocalized into benzene ring via resonance.\n\n- **Distinguishing Tests & Chemical Reactions:**\n - **Hinsberg Test (Benzenesulfonyl chloride, \text{C}_6\text{H}_5\text{SO}_2\text{Cl}):**\n - Primary amine: forms sulfonamide soluble in alkali.\n - Secondary amine: forms sulfonamide insoluble in alkali.\n - Tertiary amine: does not react.\n - **Carbylamine Test:** Primary aliphatic/aromatic amine cooked with \text{CHCl}_3 + \text{KOH}$gives foul-smelling isocyanide ($\text{R-NC}).\n - **Hoffmann Bromamide Degradation:** Primary amide + \text{Br}_2 + 4\text{NaOH} \rightarrow$Primary amine ($\text{R-NH}_2) with one carbon less.\n - **Gabriel Phthalimide Synthesis:** Used for primary aliphatic amines. Cannot prepare aromatic primary amines because aryl halides do not undergo nucleophilic substitution with phthalimide anion.\n - **Aniline Acetylation Before Nitration:** Direct nitration of aniline produces $51\% para, $47\% meta, $2\% ortho derivatives and causes oxidation. Acetylation protects $\text{-NH}_2$ group as acetanilide, controlling reactivity and yielding exclusively para-nitroaniline after hydrolysis.

Biomolecules and Biological Macromolecules

• Glucose Structure Elucidation Reactions:

  • Prolonged heating with HI gives $n$-hexane $\rightarrow$ shows straight chain of 6 carbons.
  • Reaction with hydroxylamine ($\text{NH}_2\text{OH}$) gives oxime and with HCN gives cyanohydrin $\rightarrow$ confirms presence of carbonyl group.
  • Oxidation with $\text{Br}_2$ water gives gluconic acid $\rightarrow$ confirms carbonyl group is aldehyde.
  • Acetylation with acetic anhydride ($(\text{CH}_3\text{CO})_2\text{O}$) gives pentaacetate $\rightarrow$ confirms presence of five $\text{-OH}$ groups attached to different carbon atoms.
  • Oxidation with conc. $\text{HNO}_3$ gives saccharic acid (dicarboxylic acid) $\rightarrow$ confirms presence of primary alcohol group.
  • Limitations of Open Chain Structure: Glucose does not give Schiff's test, does not form hydrogen sulfite adduct with $\text{NaHSO}_3$, pentaacetate does not react with $\text{NH}_2\text{OH}$, exists in $\alpha$ and $\beta$ cyclic hemiacetal forms.

• Proteins and Nucleic Acids:

  • Peptide Linkage: Amide bond ($\text{-CO-NH-}$) formed between $\text{-COOH}$ group of one $\alpha$-amino acid and $\text{-NH}_2$ group of another.
  • Fibrous vs Globular Proteins:
  • Fibrous: Polypeptide chains run parallel, held by hydrogen and disulfide bonds; insoluble in water (e.g., keratin, myosin).
  • Globular: Chains coil around into spherical shapes; soluble in water (e.g., insulin, albumin).
  • Nucleic Acids: Polymers of nucleotides made of pentose sugar, nitrogenous base, and phosphate moiety.
  • DNA vs RNA: DNA contains 2-deoxyribose sugar and Thymine base (Thymine, Adenine, Guanine, Cytosine). RNA contains Ribose sugar and Uracil base instead of Thymine.

• Vitamins:

  • Fat-Soluble: Vitamins A, D, E, K. (Vitamin E deficiency increases fragility of RBCs and muscular weakness).
  • Water-Soluble: Vitamins B and C.

Comprehensive Exam Practice Problems, Case Studies, and Worked Solutions

• Problem 1 (Kinetics Gas Phase Decomposition):
  • Reaction: $\text{C}_2\text{H}_5\text{Cl}(g) \rightarrow \text{C}_2\text{H}_4(g) + \text{HCl}(g)$
  • Data: At $t = 0\,s$, $P_{total} = 0.30\,atm$; At $t = 30\,s$, $P_{total} = 0.50\,atm$. Given \log(3) = 0.48$.\n - **Solution Steps:**\n - P_0 = 0.30\,atm\n - P_{total} = P_0 + x \implies 0.50 = 0.30 + x \implies x = 0.20\,atm\n - P_{\text{C}2\text{H}_5\text{Cl}} = P_0 - x = 0.30 - 0.20 = 0.10\,atm\n - k = \frac{2.303}{t} \log\left(\frac{P_0}{P_t}\right) = \frac{2.303}{30} \log\left(\frac{0.30}{0.10}\right) = \frac{2.303}{30} \log(3)\n - k = \frac{2.303 \times 0.48}{30} = 3.68 \times 10^{-2}\,s^{-1}\n\n- **Problem 2 (Boiling Point Elevation & Dimerization):**\n - **Data:** 0.61\,g$benzoic acid ($M = 122\,g\,mol^{-1}$) in$5\,g$$\text{CS}_2$, dimerises$88\%$.$T_b^\circ(\text{CS}_2) = 46.2\,^\circ C$,$K_b = 2.3\,K\,kg\,mol^{-1}.\n - **Solution Steps:**\n - \alpha = 0.88\n - i = 1 - \alpha + \frac{\alpha}{2} = 1 - 0.88 + 0.44 = 0.56\n - Molality m = \frac{0.61 / 122}{5 / 1000} = \frac{0.005}{0.005} = 1.0\,m\n - \Delta T_b = i \times K_b \times m = 0.56 \times 2.3 \times 1.0 = 1.288\,K\n - Boiling Point T_b = 46.2 + 1.288 = 47.488\,^\circ C\n\n- **Problem 3 (Cell EMF and Gibbs Free Energy Calculation):**\n - **Cell:** \text{Mg}(s) | \text{Mg}^{2+}(0.01\,M) || \text{Ag}^+(0.001\,M) | \text{Ag}(s)\n - **Data:** E^\circ{\text{Mg}^{2+}/\text{Mg}} = -2.37\,V$,$E^\circ_{\text{Ag}^+/\text{Ag}} = +0.80\,V$,$1\,F = 96500\,C\,mol^{-1}.\n - **Solution Steps:**\n - E^\circ_{cell} = 0.80 - (-2.37) = +3.17\,V\n - Cell reaction: \text{Mg}(s) + 2\text{Ag}^+(aq) \rightarrow \text{Mg}^{2+}(aq) + 2\text{Ag}(s)$,$n = 2\n - E_{cell} = E^\circ_{cell} - \frac{0.0591}{2} \log\left(\frac{[\text{Mg}^{2+}]}{[\text{Ag}^+]^2}\right) = 3.17 - 0.02955 \log\left(\frac{10^{-2}}{(10^{-3})^2}\right)\n - E_{cell} = 3.17 - 0.02955 \log(10^4) = 3.17 - 0.02955 \times 4 = 3.17 - 0.1182 = 3.0518\,V\n - \Delta_r G = -n F E_{cell} = -2 \times 96500 \times 3.0518 = -589000\,J\,mol^{-1} = -589\,kJ\,mol^{-1}\n\n- **Problem 4 (Osmotic Pressure Calculation):**\n - **Data:** \text{CaCl}2$dissolved in$2.46\,L$water,$\Pi = 0.70\,atm$at$27\,^\circ C$($300\,K$),$i = 2.59$,$M = 111\,g\,mol^{-1}$,$R = 0.082\,L\,atm\,K^{-1}\,mol^{-1}.\n - **Solution Steps:**\n - \Pi = i \times \left(\frac{w}{M \times V}\right) \times R \times T\n - 0.70 = 2.59 \times \left(\frac{w}{111 \times 2.46}\right) \times 0.082 \times 300\n - 0.70 = w \times 0.2333 \implies w = \frac{0.70}{0.2333} = 3.0\,g\n\n- **Problem 5 (Molar Conductivity and Degree of Dissociation):**\n - **Data:** Concentration c = 0.1\,M\,\text{NaCl}$,$\kappa = 1.06 \times 10^{-2}\,S\,cm^{-1}$,$\lambda^\circ{\text{Na}^+} = 50.1\,S\,cm^2\,mol^{-1}$,$\lambda^\circ_{\text{Cl}^-} = 76.5\,S\,cm^2\,mol^{-1}$.
  • Solution Steps:
  • $\Lambda_m = \frac{\kappa \times 1000}{c} = \frac{1.06 \times 10^{-2} \times 1000}{0.1} = 106\,S\,cm^2\,mol^{-1}$
  • $\Lambda_m^\circ = 50.1 + 76.5 = 126.6\,S\,cm^2\,mol^{-1}$
  • $\alpha = \frac{\Lambda_m}{\Lambda_m^\circ} = \frac{106}{126.6} = 0.837$ or $83.7\%$