JEE Chemistry 2026 Cheatsheet

Cheatsheet Content

1. Basic Concepts of Chemistry Mole Concept: $N_A = 6.022 \times 10^{23}$ particles/mole. $$ \text{Moles} = \frac{\text{Given Mass}}{\text{Molar Mass}} = \frac{\text{Number of Particles}}{N_A} = \frac{\text{Volume of Gas at STP}}{22.4 \text{ L}} $$ Molar Mass: Sum of atomic masses of all atoms in a molecule. Percentage Composition: $$ \text{% Element} = \frac{\text{Mass of Element}}{\text{Molar Mass of Compound}} \times 100 $$ Empirical Formula: Simplest whole number ratio of atoms. Molecular Formula: Actual number of atoms. Stoichiometry: Quantitative relations in balanced reactions. Example: $2H_2 + O_2 \to 2H_2O$. 2 moles $H_2$ react with 1 mole $O_2$ to give 2 moles $H_2O$. Limiting Reagent (LR): Reactant consumed first, dictates product yield. Concentration Terms: Molarity (M): Moles of solute / Volume of solution (L). $(M_1V_1 = M_2V_2 \text{ for dilution})$ Molality (m): Moles of solute / Mass of solvent (kg). (Temperature independent) Mole Fraction ($X$): Moles of component / Total moles. ($X_A + X_B = 1$) Mass Percentage: (Mass of solute / Mass of solution) $\times 100$. ppm (parts per million): (Mass of solute / Mass of solution) $\times 10^6$. 2. Atomic Structure Electromagnetic Spectrum: $E = h\nu = hc/\lambda$. Bohr's Model: For H-atom/H-like species. Radius: $r_n = 0.529 \frac{n^2}{Z}$ Å. Energy: $E_n = -13.6 \frac{Z^2}{n^2}$ eV. Velocity: $v_n = 2.18 \times 10^6 \frac{Z}{n}$ m/s. Frequency of revolution $\propto Z^2/n^3$. Rydberg Formula: For spectral lines (emission): $$ \frac{1}{\lambda} = R_H Z^2 \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right) $$ where $R_H = 109677 \text{ cm}^{-1}$ (Rydberg constant), $n_2 > n_1$. De Broglie Wavelength: Particle-wave duality. $$ \lambda = \frac{h}{mv} = \frac{h}{p} $$ Heisenberg's Uncertainty Principle: $$ \Delta x \cdot \Delta p \ge \frac{h}{4\pi} \quad \text{or} \quad \Delta x \cdot m\Delta v \ge \frac{h}{4\pi} $$ Quantum Numbers: Describe atomic orbitals and electrons. Principal (n): $1, 2, 3...$ (Shell, size, energy). Azimuthal (l): $0, 1, 2... (n-1)$ (Subshell, shape; s, p, d, f). Magnetic ($m_l$): $-l, ..., 0, ..., +l$ (Orientation). Spin ($m_s$): $+1/2, -1/2$ (Electron spin). Rules for Filling Orbitals: Aufbau Principle: Fill from lowest to highest energy ($n+l$ rule). Pauli's Exclusion Principle: Max 2 electrons per orbital, with opposite spins. Hund's Rule of Maximum Multiplicity: Maximize unpaired electrons in degenerate orbitals before pairing. 3. Chemical Bonding and Molecular Structure Ionic Bond: Electrostatic attraction between ions formed by electron transfer. High EN difference. Covalent Bond: Sharing of electrons. Low EN difference. Bond Length: Average distance between nuclei. Bond Energy: Energy required to break a bond. Bond Order: Number of bonds between two atoms. Dipole Moment ($\mu$): $\mu = q \times d$. (Charge $\times$ distance). Vector quantity. Example: $H_2O$ has dipole moment, $CO_2$ has zero dipole moment (linear structure). Fajan's Rules (Covalent character in Ionic Bonds): Small cation, large anion $\implies$ more covalent character. High charge on cation/anion $\implies$ more covalent character. Cations with pseudo-noble gas configuration ($ns^2np^6nd^{10}$) are more polarizing than noble gas configuration ($ns^2np^6$). (e.g., $Cu^+$ vs $Na^+$). Valence Shell Electron Pair Repulsion (VSEPR) Theory: Predicts geometry based on minimizing electron pair repulsion. Order of Repulsion: Lone Pair-Lone Pair > Lone Pair-Bond Pair > Bond Pair-Bond Pair. Steric No. Hybridization Geometry Example 2 sp Linear $BeCl_2, CO_2$ 3 $sp^2$ Trigonal Planar $BF_3, SO_3$ 4 $sp^3$ Tetrahedral $CH_4$ 4 (3 BP, 1 LP) $sp^3$ Pyramidal $NH_3$ 4 (2 BP, 2 LP) $sp^3$ Bent $H_2O$ 5 $sp^3d$ Trigonal Bipyramidal $PCl_5$ 5 (4 BP, 1 LP) $sp^3d$ See-saw $SF_4$ 5 (3 BP, 2 LP) $sp^3d$ T-shaped $ClF_3$ 5 (2 BP, 3 LP) $sp^3d$ Linear $XeF_2$ 6 $sp^3d^2$ Octahedral $SF_6$ 6 (5 BP, 1 LP) $sp^3d^2$ Square Pyramidal $BrF_5$ 6 (4 BP, 2 LP) $sp^3d^2$ Square Planar $XeF_4$ Hybridization: Mixing of atomic orbitals to form new hybrid orbitals. $$ \text{Steric No.} = (\text{No. of } \sigma \text{ bonds}) + (\text{No. of lone pairs}) $$ Molecular Orbital Theory (MOT): Atomic orbitals combine to form Bonding Molecular Orbitals (BMO) and Anti-bonding Molecular Orbitals (ABMO). Bond Order: $$ BO = \frac{1}{2} (\text{Number of electrons in BMO} - \text{Number of electrons in ABMO}) $$ Magnetic properties: Paramagnetic (unpaired electrons), Diamagnetic (all electrons paired). Example: $O_2$ (BO=2, paramagnetic), $N_2$ (BO=3, diamagnetic). Hydrogen Bonding: Strong dipole-dipole interaction between H (bonded to F, O, N) and F, O, N of another molecule. Intermolecular H-bonding: Increases boiling point, solubility. (e.g., $H_2O$, alcohols). Intramolecular H-bonding: Decreases boiling point, solubility. (e.g., o-nitrophenol). 4. States of Matter Gas Laws: Boyle's Law: $P_1V_1 = P_2V_2$ (T, n constant). $P \propto 1/V$. Charles' Law: $V_1/T_1 = V_2/T_2$ (P, n constant). $V \propto T$. Gay-Lussac's Law: $P_1/T_1 = P_2/T_2$ (V, n constant). $P \propto T$. Avogadro's Law: $V_1/n_1 = V_2/n_2$ (P, T constant). $V \propto n$. Ideal Gas Equation: $$ PV = nRT $$ $R = 0.0821 \text{ L atm mol}^{-1} \text{ K}^{-1} = 8.314 \text{ J mol}^{-1} \text{ K}^{-1}$. Dalton's Law of Partial Pressures: $P_{total} = P_A + P_B + ...$. $P_A = X_A P_{total}$. Graham's Law of Diffusion/Effusion: $$ \frac{r_1}{r_2} = \sqrt{\frac{M_2}{M_1}} $$ Kinetic Theory of Gases (K.T.G.): Root mean square speed: $u_{rms} = \sqrt{3RT/M}$. Average speed: $u_{avg} = \sqrt{8RT/(\pi M)}$. Most probable speed: $u_{mp} = \sqrt{2RT/M}$. Real Gases (Van der Waals Equation): $$ \left(P + \frac{an^2}{V^2}\right)(V - nb) = nRT $$ 'a' accounts for intermolecular forces, 'b' for molecular volume. Compressibility Factor (Z): $Z = PV/(nRT)$. $Z=1$ for ideal gas. $Z 1$ (repulsive forces dominate). Solid State: Unit Cells: Simple Cubic (SC): Atoms at corners. $Z=1$. $r=a/2$. Packing efficiency $52.4\%$. Body Centered Cubic (BCC): Atoms at corners + center. $Z=2$. $r=\sqrt{3}a/4$. Packing efficiency $68\%$. Face Centered Cubic (FCC): Atoms at corners + face centers. $Z=4$. $r=\sqrt{2}a/4$. Packing efficiency $74\%$. Defects: Stoichiometric (Schottky, Frenkel), Non-stoichiometric (metal excess/deficiency). 5. Thermodynamics First Law of Thermodynamics: (Conservation of Energy) $$ \Delta U = q + w $$ $\Delta U$ = change in internal energy, $q$ = heat, $w$ = work. Work done by gas: $w = -P_{ext}\Delta V$. Work done on gas: $w = -P_{int}\Delta V$. Enthalpy ($\Delta H$): Heat change at constant pressure. $\Delta H = \Delta U + P\Delta V$. $$ \Delta H = q_p $$ Relation between $\Delta H$ and $\Delta U$: $$ \Delta H = \Delta U + \Delta n_g RT $$ $\Delta n_g$ = moles of gaseous products - moles of gaseous reactants. Hess's Law: Enthalpy change for a reaction is independent of the path. $$ \Delta H_{reaction}^\circ = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants}) $$ $$ \Delta H_{reaction}^\circ = \sum \text{Bond Energies}(\text{reactants}) - \sum \text{Bond Energies}(\text{products}) $$ Second Law of Thermodynamics: For a spontaneous process, total entropy of universe increases. $$ \Delta S_{univ} = \Delta S_{sys} + \Delta S_{surr} > 0 $$ Entropy ($\Delta S$): Measure of disorder/randomness. $$ \Delta S = \frac{q_{rev}}{T} $$ Third Law of Thermodynamics: Entropy of a perfect crystalline substance at absolute zero (0 K) is zero. Gibbs Free Energy ($\Delta G$): Predicts spontaneity. $$ \Delta G = \Delta H - T\Delta S $$ $\Delta G $\Delta G > 0$: Non-spontaneous. $\Delta G = 0$: At Equilibrium. Relation between $\Delta G$ and equilibrium constant (K): $$ \Delta G^\circ = -RT \ln K_{eq} $$ $$ \Delta G = \Delta G^\circ + RT \ln Q $$ 6. Chemical Equilibrium Law of Mass Action: For $aA + bB \rightleftharpoons cC + dD$: $$ K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b} $$ $$ K_p = \frac{(P_C)^c(P_D)^d}{(P_A)^a(P_B)^b} $$ Relation between $K_p$ and $K_c$: $$ K_p = K_c (RT)^{\Delta n_g} $$ Reaction Quotient (Q): Same expression as K, but for non-equilibrium concentrations. $Q $Q > K$: Net reaction proceeds backward. $Q = K$: Equilibrium. Le Chatelier's Principle: If a change is applied to a system at equilibrium, the system shifts in a way that relieves the stress. Concentration: Add reactant, shift right; add product, shift left. Pressure: Increase P, shift to side with fewer gas moles. Decrease P, shift to side with more gas moles. Temperature: Increase T, shift in endothermic direction. Decrease T, shift in exothermic direction. Catalyst: No effect on equilibrium position, only speeds up attainment of equilibrium. 7. Ionic Equilibrium Acids and Bases: Arrhenius: Acid ($H^+$ donor), Base ($OH^-$ donor). Brønsted-Lowry: Acid (proton donor), Base (proton acceptor). Conjugate acid-base pairs. Lewis: Acid (electron pair acceptor), Base (electron pair donor). pH Scale: $$ pH = -\log[H^+] \quad pOH = -\log[OH^-] $$ $$ pH + pOH = 14 \quad (\text{at } 25^\circ C) $$ $$ K_w = [H^+][OH^-] = 10^{-14} \quad (\text{at } 25^\circ C) $$ Acid/Base Dissociation Constants: $$ K_a = \frac{[H^+][A^-]}{[HA]} \quad pK_a = -\log K_a $$ $$ K_b = \frac{[BH^+][OH^-]}{[B]} \quad pK_b = -\log K_b $$ For conjugate acid-base pair: $K_a \cdot K_b = K_w$. Hydrolysis of Salts: Strong Acid + Strong Base: pH = 7. Strong Acid + Weak Base: Acidic, $pH Weak Acid + Strong Base: Basic, $pH > 7$. (e.g., $CH_3COONa$) Weak Acid + Weak Base: $pH$ depends on $K_a$ and $K_b$. Buffer Solutions: Resist change in pH upon addition of small amounts of acid/base. Acidic Buffer: Weak acid + its conjugate base (salt). $$ pH = pK_a + \log \frac{[\text{Salt}]}{[\text{Acid}]} \quad (\text{Henderson-Hasselbalch Equation}) $$ Basic Buffer: Weak base + its conjugate acid (salt). $$ pOH = pK_b + \log \frac{[\text{Salt}]}{[\text{Base}]} $$ Solubility Product ($K_{sp}$): For a sparingly soluble salt $A_xB_y \rightleftharpoons xA^{y+} + yB^{x-}$: $$ K_{sp} = [A^{y+}]^x [B^{x-}]^y $$ If Ionic Product (IP) $ If IP $= K_{sp}$, solution is saturated, equilibrium. If IP $> K_{sp}$, precipitation occurs. Common Ion Effect: Solubility of a sparingly soluble salt decreases in the presence of a common ion. 8. Redox Reactions and Electrochemistry Redox Definitions: Oxidation: Loss of electrons, increase in oxidation number. Reduction: Gain of electrons, decrease in oxidation number. Oxidizing Agent: Itself reduced, oxidizes others. Reducing Agent: Itself oxidized, reduces others. Balancing Redox Reactions: Oxidation number method or Ion-electron method. Electrochemical Cells: Galvanic (Voltaic) Cell: Chemical energy $\to$ Electrical energy. Spontaneous. Anode: Oxidation (negative electrode). Cathode: Reduction (positive electrode). Salt bridge: Maintains electrical neutrality. Cell Notation: $Anode | Anode \text{ ion} || Cathode \text{ ion} | Cathode$. Electrolytic Cell: Electrical energy $\to$ Chemical energy. Non-spontaneous. Standard Electrode Potential ($E^\circ$): Reduction potential measured against Standard Hydrogen Electrode (SHE, $E^\circ=0$). Standard Cell Potential ($E^\circ_{cell}$): $$ E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode} $$ (Both are standard reduction potentials). Nernst Equation: Relates cell potential to concentrations. $$ E_{cell} = E^\circ_{cell} - \frac{RT}{nF} \ln Q $$ At $25^\circ C$ (298 K): $$ E_{cell} = E^\circ_{cell} - \frac{0.0591}{n} \log Q $$ Relationship between $\Delta G^\circ$, $E^\circ_{cell}$, and $K_{eq}$: $$ \Delta G^\circ = -nFE^\circ_{cell} $$ $$ \Delta G^\circ = -RT \ln K_{eq} $$ $$ E^\circ_{cell} = \frac{RT}{nF} \ln K_{eq} = \frac{0.0591}{n} \log K_{eq} \quad (\text{at } 25^\circ C) $$ (where F = 96485 C/mol, Faraday's constant). Faraday's Laws of Electrolysis: First Law: Mass deposited/liberated is proportional to charge passed. $$ W = ZIt \quad (\text{where } Z = \text{Electrochemical equivalent} = \text{Equivalent weight}/F) $$ Second Law: If same charge is passed through different electrolytes, masses deposited are proportional to their equivalent weights. $$ \frac{W_1}{W_2} = \frac{E_1}{E_2} $$ 9. Chemical Kinetics Rate of Reaction: Change in concentration of reactant/product per unit time. For $aA + bB \to cC + dD$: $$ \text{Rate} = -\frac{1}{a}\frac{d[A]}{dt} = -\frac{1}{b}\frac{d[B]}{dt} = \frac{1}{c}\frac{d[C]}{dt} = \frac{1}{d}\frac{d[D]}{dt} $$ Rate Law: Rate $= k[A]^x[B]^y$. Order of Reaction: Sum of powers of concentration terms ($x+y$). Can be fractional or zero. Molecularity: Number of reacting species in an elementary step. Always a whole number. Integrated Rate Laws and Half-Life ($t_{1/2}$): Zero Order: Rate $= k$. $$ [A]_t = [A]_0 - kt $$ $$ t_{1/2} = \frac{[A]_0}{2k} $$ First Order: Rate $= k[A]$. $$ \ln[A]_t = \ln[A]_0 - kt \quad \text{or} \quad [A]_t = [A]_0 e^{-kt} $$ $$ t_{1/2} = \frac{0.693}{k} $$ Second Order: Rate $= k[A]^2$. $$ \frac{1}{[A]_t} = \frac{1}{[A]_0} + kt $$ $$ t_{1/2} = \frac{1}{k[A]_0} $$ Arrhenius Equation: Effect of temperature on rate constant. $$ k = Ae^{-E_a/(RT)} $$ $$ \ln \left(\frac{k_2}{k_1}\right) = \frac{E_a}{R} \left(\frac{1}{T_1} - \frac{1}{T_2}\right) $$ $E_a$ = Activation energy, $A$ = Pre-exponential factor. Collision Theory: Reaction occurs when molecules collide with sufficient energy ($>E_a$) and correct orientation. 10. Solutions and Colligative Properties Ideal Solutions: Obeys Raoult's Law. $\Delta H_{mix}=0, \Delta V_{mix}=0$. Raoult's Law: For volatile components, partial vapor pressure of a component is proportional to its mole fraction. $$ P_A = X_A P_A^\circ $$ For non-volatile solute: $P_{solution} = X_{solvent} P_{solvent}^\circ$. Colligative Properties: Depend only on number of solute particles, not their nature. Relative Lowering of Vapor Pressure (RLVP): $$ \frac{P^\circ - P_s}{P^\circ} = X_{solute} = \frac{n_{solute}}{n_{solute} + n_{solvent}} $$ Elevation in Boiling Point ($\Delta T_b$): $$ \Delta T_b = K_b m $$ $K_b$ = molal elevation constant (ebullioscopic constant). Depression in Freezing Point ($\Delta T_f$): $$ \Delta T_f = K_f m $$ $K_f$ = molal depression constant (cryoscopic constant). Osmotic Pressure ($\Pi$): $$ \Pi = iCRT $$ $C$ = molarity, $R$ = gas constant, $T$ = temperature (K). Van't Hoff Factor ($i$): Accounts for dissociation or association of solute. $$ i = \frac{\text{Observed Colligative Property}}{\text{Normal Colligative Property}} = \frac{\text{Normal Molar Mass}}{\text{Observed Molar Mass}} $$ For dissociation: $i = 1 + (n-1)\alpha$. For association: $i = 1 + (1/n - 1)\alpha$. 11. Surface Chemistry Adsorption: Accumulation of molecular species on the surface. Physisorption: Weak van der Waals forces, multilayer, reversible, low $\Delta H_{ads}$. Chemisorption: Chemical bonds, monolayer, irreversible, high $\Delta H_{ads}$. Freundlich Adsorption Isotherm: $$ x/m = kP^{1/n} \quad (n>1) $$ $\log(x/m) = \log k + (1/n)\log P$. Catalysis: Homogeneous: Catalyst and reactants in same phase. (e.g., acid hydrolysis of ester). Heterogeneous: Catalyst and reactants in different phases. (e.g., Haber process for $NH_3$). Enzyme Catalysis: Highly specific biological catalysts. Colloids: Heterogeneous systems with particle size 1-1000 nm. Types: Sol (solid in liquid), Gel (liquid in solid), Emulsion (liquid in liquid), Foam (gas in liquid), Aerosol (solid/liquid in gas). Properties: Tyndall effect (scattering of light), Brownian movement (random motion), Electrophoresis (movement in electric field), Coagulation/Flocculation (precipitation by adding electrolyte). Hardy-Schulze Rule: Greater the valence of the flocculating ion, greater is its coagulating power. 12. Periodicity and Classification of Elements Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers. Periodic Trends: Atomic Radius: Decreases across period (due to increased $Z_{eff}$), Increases down group (new shells). Ionization Enthalpy ($IE_1$): Energy required to remove 1st electron. Increases across period, Decreases down group. Exceptions: $IE_1(N) > IE_1(O)$, $IE_1(Be) > IE_1(B)$. Electron Gain Enthalpy ($\Delta H_{eg}$): Energy released/absorbed when an electron is added. Generally becomes more negative across period, less negative down group. Exceptions: Group 17 (halogens) highly negative. Group 18 (noble gases) positive. $Cl > F$, $S > O$. Electronegativity (EN): Tendency of an atom to attract shared electrons. Increases across period, Decreases down group. (Pauling scale). F > O > N $\approx$ Cl. Metallic Character: Decreases across period, Increases down group. Oxidizing Nature: Increases across period, Decreases down group. Reducing Nature: Decreases across period, Increases down group. Diagonal Relationship: Similarities between elements diagonally placed in 2nd and 3rd periods (e.g., Li-Mg, Be-Al, B-Si). Due to similar charge/radius ratio. 13. General Principles and Processes of Isolation of Elements (Metallurgy) Concentration of Ore: Gravity Separation (Levigation): For heavy oxide/carbonate ores (e.g., SnO$_2$). Magnetic Separation: For magnetic ores (e.g., Magnetite $Fe_3O_4$, Chromite $FeCr_2O_4$). Froth Flotation: For sulfide ores (e.g., ZnS, PbS). Uses pine oil, collectors, froth stabilizers. Leaching: Chemical method using a suitable solvent. (e.g., Bauxite with NaOH for Al). Conversion of Concentrated Ore to Oxide: Calcination: Heating in absence/limited supply of air below melting point. Removes volatile impurities ($H_2O, CO_2$). (e.g., $CaCO_3 \to CaO + CO_2$). Roasting: Heating in excess air below melting point. For sulfide ores, converts to oxide. (e.g., $2ZnS + 3O_2 \to 2ZnO + 2SO_2$). Reduction of Metal Oxide to Metal: Smelting (Pyrometallurgy): Reduction with C, CO or other reducing agents at high temp. Example: $Fe_2O_3 + 3CO \to 2Fe + 3CO_2$ (Blast Furnace). Electrolytic Reduction: For highly electropositive metals (Na, Mg, Al). (e.g., Hall-Heroult process for Al). Hydrometallurgy: Reduction by more reactive metal from aqueous solution. (e.g., $Cu^{2+} + Zn \to Cu + Zn^{2+}$). Refining: Distillation: For low boiling point metals (Zn, Hg, Cd). Liquation: For low melting point metals (Sn, Pb). Electrolytic Refining: For Cu, Zn, Al, etc. Impure metal as anode, pure metal as cathode. Zone Refining: For ultra-pure semiconductors (Si, Ge). Based on principle that impurities are more soluble in molten metal. Vapour Phase Refining: For metals like Ni (Mond process), Ti, Zr (Van Arkel method). Metal converted to volatile compound, then decomposed to pure metal. Thermodynamic Principles (Ellingham Diagram): Plots $\Delta G^\circ$ vs T for formation of metal oxides. Explains feasibility of reduction. Lower line can reduce oxide whose line is above it. Slope indicates $\Delta S$. Positive slope means $\Delta S 14. Hydrogen Isotopes: Protium ($^1H$), Deuterium ($^2H$, D), Tritium ($^3H$, T, radioactive). Preparation: Laboratory: $Zn + dil. HCl \to ZnCl_2 + H_2$. Industrial: Electrolysis of water, Bosch process (steam reforming of hydrocarbons). Properties: Lightest element, flammable gas, reducing agent. Hydrides: Ionic (Saline): Group 1 & 2 metals react with H$_2$. (e.g., $NaH, CaH_2$). Covalent (Molecular): p-Block elements. (e.g., $CH_4, NH_3, H_2O$). Metallic (Interstitial): d- & f-Block elements. Non-stoichiometric. Water ($H_2O$): High boiling point, specific heat, heat of vaporization due to H-bonding. Hard Water: Contains $Ca^{2+}, Mg^{2+}$ ions. Temporary Hardness: Due to bicarbonates. Removed by boiling (Clark's method). Permanent Hardness: Due to chlorides/sulfates. Removed by washing soda, Calgon, ion-exchange. Hydrogen Peroxide ($H_2O_2$): Structure (non-planar, open book). Oxidizing and reducing agent. Oxidizing: $2KI + H_2O_2 \to 2KOH + I_2$. Reducing: $Ag_2O + H_2O_2 \to 2Ag + H_2O + O_2$. 15. s-Block Elements Alkali Metals (Group 1: Li, Na, K, Rb, Cs): Properties: Soft, silvery-white, low melting/boiling points, highly reactive, strong reducing agents. Form ionic compounds. Reducing power: $Li > Cs > Rb > K > Na$ (in aqueous solution due to high hydration enthalpy of Li$^+$). Flame Coloration: Li (crimson red), Na (golden yellow), K (lilac). Reactions: React with water ($2Na + 2H_2O \to 2NaOH + H_2$), form oxides ($Na_2O$), peroxides ($Na_2O_2$), superoxides ($KO_2$). Important Compounds: $NaOH$ (Caustic Soda): Preparation (Castner-Kellner cell). Strong base. $Na_2CO_3$ (Washing Soda): Preparation (Solvay process). $NaHCO_3$ (Baking Soda): Less soluble, precipitates in Solvay process. Alkaline Earth Metals (Group 2: Be, Mg, Ca, Sr, Ba, Ra): Properties: Harder, higher melting points than Group 1, less reactive than Group 1. Strong reducing agents. Flame Coloration: Ca (brick red), Sr (crimson), Ba (apple green). Be and Mg do not give flame test. Diagonal Relationship: Be with Al (similar charge/radius ratio, forms covalent compounds, amphoteric oxide). Important Compounds: $CaO$ (Quicklime): Used in cement. $CaCO_3$ (Limestone): Various forms (marble, chalk). Plaster of Paris ($CaSO_4 \cdot 1/2 H_2O$): From gypsum ($CaSO_4 \cdot 2H_2O$). Milk of Magnesia ($Mg(OH)_2$): Antacid. 16. p-Block Elements Group 13 (Boron Family: B, Al, Ga, In, Tl): Boron: Non-metal, forms covalent compounds. Unique behavior due to small size, high IE. Aluminium: Metal, amphoteric oxide ($Al_2O_3$). Electron Deficient Compounds: Tendency to form trihalides (e.g., $BF_3$). $BF_3$ is a Lewis acid. Diborane ($B_2H_6$): Banana bonds (3-center 2-electron bonds). B B H H H H H H Structure of Diborane: Two B atoms and two bridging H atoms form 3-center 2-electron bonds. Four terminal H atoms form normal 2-center 2-electron bonds. Group 14 (Carbon Family: C, Si, Ge, Sn, Pb): Catenation: Tendency to form long chains/rings. Max for Carbon. Allotropes of Carbon: Diamond: $sp^3$ hybridized, tetrahedral, hardest substance. Graphite: $sp^2$ hybridized, layered, good conductor, lubricant. Fullerenes ($C_{60}$): Buckminsterfullerene, cage-like structure. Oxides: $CO$ (neutral), $CO_2$ (acidic), $SiO_2$ (acidic), $SnO_2, PbO_2$ (amphoteric). Silicates: Basic structural unit $SiO_4^{4-}$. Group 15 (Nitrogen Family: N, P, As, Sb, Bi): Nitrogen: Diatomic ($N_2$), triple bond, inert. Ammonia ($NH_3$): Haber process. Lewis base. Pyramidal shape. Nitric Acid ($HNO_3$): Ostwald process. Strong oxidizing agent. Phosphorus: Allotropes (White, Red, Black). White P is most reactive, stored in water. Phosphine ($PH_3$): Highly poisonous, weaker base than $NH_3$. Oxoacids of Phosphorus: $H_3PO_2$ (hypophosphorous, reducing), $H_3PO_3$ (phosphorous, reducing), $H_3PO_4$ (phosphoric). Group 16 (Oxygen Family: O, S, Se, Te, Po): Oxygen: Diatomic ($O_2$), paramagnetic. Ozone ($O_3$): Allotrope of oxygen, strong oxidizing agent. Bent shape. Sulfur: Allotropes (Rhombic, Monoclinic). Rhombic is stable at room temp. Sulfuric Acid ($H_2SO_4$): Contact process. Strong dehydrating agent, oxidizing agent, acid. Hydrides ($H_2O, H_2S, H_2Se, H_2Te$): Thermal stability decreases down group. Acidic strength increases down group. Group 17 (Halogen Family: F, Cl, Br, I, At): Properties: Highly reactive non-metals, strong oxidizing agents. Reactivity $F_2 > Cl_2 > Br_2 > I_2$. Bond Dissociation Enthalpy: $Cl_2 > Br_2 > F_2 > I_2$ (due to lone pair repulsion in $F_2$). Oxidizing Power: $F_2 > Cl_2 > Br_2 > I_2$. Hydrogen Halides ($HX$): Acidic strength $HF Interhalogen Compounds: $XY, XY_3, XY_5, XY_7$. More reactive than constituent halogens (except $F_2$). Group 18 (Noble Gases: He, Ne, Ar, Kr, Xe, Rn): Properties: Inert, stable electron configuration. Low boiling points. Xenon Compounds: $XeF_2, XeF_4, XeF_6$. (Can be formed due to lower IE and availability of d-orbitals). Structure of $XeF_4$: Square Planar ($sp^3d^2$ with 2 lone pairs). 17. d- and f-Block Elements d-Block (Transition Elements): Properties: High melting/boiling points, high tensile strength, variable oxidation states, colored ions, paramagnetism (unpaired electrons), catalytic activity, complex formation. Electronic Configuration: $(n-1)d^{1-10} ns^{1-2}$. Exceptions: Cr ($3d^5 4s^1$), Cu ($3d^{10} 4s^1$). Oxidation States: Exhibit multiple OS (e.g., Mn from +2 to +7). Ionization Enthalpy: Generally increases from left to right. Standard Electrode Potentials: Generally negative, indicate reducing nature. $E^\circ(Cu^{2+}/Cu)$ is positive. Important Compounds: Potassium Dichromate ($K_2Cr_2O_7$): Orange. Strong oxidizing agent in acidic medium. $$ Cr_2O_7^{2-} + 14H^+ + 6e^- \to 2Cr^{3+} + 7H_2O $$ Interconvertible with chromate ($CrO_4^{2-}$ yellow) in pH change. Potassium Permanganate ($KMnO_4$): Purple. Strong oxidizing agent. $$ MnO_4^- + 8H^+ + 5e^- \to Mn^{2+} + 4H_2O \quad (\text{acidic medium}) $$ $$ MnO_4^- + 2H_2O + 3e^- \to MnO_2 + 4OH^- \quad (\text{neutral/weakly alkaline}) $$ f-Block (Inner Transition Elements): Lanthanoids (Ce to Lu): Electronic Configuration: $[Xe] 4f^{0-14} 5d^{0-1} 6s^2$. Oxidation State: Dominant +3. Some show +2, +4. Lanthanoid Contraction: Steady decrease in atomic and ionic radii from Ce to Lu. Causes similar radii for 4d and 5d elements (e.g., Zr and Hf). Actinoids (Th to Lr): Electronic Configuration: $[Rn] 5f^{0-14} 6d^{0-1} 7s^2$. Oxidation State: More variable than lanthanoids (e.g., U, Np, Pu show +3, +4, +5, +6). Radioactive. 18. Coordination Compounds Definitions: Central Metal Atom/Ion: Lewis acid. Ligand: Lewis base (electron pair donor). Monodentate, bidentate, polydentate. Coordination Number: Number of ligand donor atoms directly attached to central metal. Counter Ion: Balances charge, outside coordination sphere. Coordination Sphere: Central metal + ligands. Nomenclature (IUPAC): Ligands named first (alphabetically), then metal. Oxidation state of metal in Roman numerals. Isomerism: Structural Isomerism: Linkage: Ambidentate ligands ($NO_2^-, SCN^-$). (e.g., $[Co(NH_3)_5(NO_2)]^{2+}$ vs $[Co(NH_3)_5(ONO)]^{2+}$). Coordination: Exchange of ligands between cationic and anionic complexes. (e.g., $[Co(NH_3)_6][Cr(CN)_6]$ vs $[Cr(NH_3)_6][Co(CN)_6]$). Ionization: Exchange of ligand with counter ion. (e.g., $[Co(NH_3)_5Br]SO_4$ vs $[Co(NH_3)_5SO_4]Br$). Hydrate: Water molecule acts as ligand or solvent. (e.g., $[Cr(H_2O)_6]Cl_3$ vs $[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O$). Stereoisomerism: Geometric (cis-trans): For square planar ($MA_2B_2$, $MA_2BC$) and octahedral ($MA_4B_2$, $MA_3B_3$) complexes. Optical (Enantiomers): Non-superimposable mirror images. Chiral complexes. (e.g., $[Co(en)_3]^{3+}$). Bonding Theories: Valence Bond Theory (VBT): Explains hybridization and magnetic properties. Inner orbital complex ($d^2sp^3$): Strong field ligands, electrons pair up. Outer orbital complex ($sp^3d^2$): Weak field ligands, electrons don't pair up. Crystal Field Theory (CFT): Explains color, magnetic properties, and stability based on ligand interaction with metal d-orbitals. Octahedral field: d-orbitals split into $t_{2g}$ (lower energy, $d_{xy}, d_{yz}, d_{zx}$) and $e_g$ (higher energy, $d_{x^2-y^2}, d_{z^2}$). Energy difference $\Delta_o$. Tetrahedral field: d-orbitals split into $e$ (lower energy) and $t_2$ (higher energy). Energy difference $\Delta_t = (4/9)\Delta_o$. Spectrochemical Series: Order of ligand field strength (ability to cause d-orbital splitting). $I^- Strong field ligands cause large $\Delta$, favor low spin (pairing). Weak field ligands cause small $\Delta$, favor high spin (no pairing). 19. Organic Chemistry - Some Basic Principles and Techniques Hybridization: $sp^3$: Single bonds, tetrahedral, 109.5$^\circ$ (e.g., alkanes). $sp^2$: Double bonds, trigonal planar, 120$^\circ$ (e.g., alkenes, carbonyls). $sp$: Triple bonds, linear, 180$^\circ$ (e.g., alkynes, nitriles). Homolytic vs. Heterolytic Fission: Homolytic: Forms free radicals (e.g., $Cl-Cl \xrightarrow{hv} Cl^\cdot + Cl^\cdot$). Heterolytic: Forms carbocations or carbanions. Reaction Intermediates: Carbocations ($R_3C^+$): $sp^2$, planar, electron deficient. Stability: $3^\circ > 2^\circ > 1^\circ > CH_3^+$. Stabilized by +I, +R effects. Carbanions ($R_3C^-$): $sp^3$, pyramidal, electron rich. Stability: $CH_3^- > 1^\circ > 2^\circ > 3^\circ$. Stabilized by -I, -R effects. Free Radicals ($R_3C^\cdot$): $sp^2$ or $sp^3$, planar or pyramidal. Stability: $3^\circ > 2^\circ > 1^\circ > CH_3^\cdot$. Stabilized by +I, +R effects. Electronic Effects: Inductive Effect (I): Permanent electron displacement along a $\sigma$-bond. $+I$ (electron releasing): Alkyl groups. $-I$ (electron withdrawing): $-NO_2, -CN, -COOH, -X$. Resonance Effect (R) / Mesomeric Effect (M): Delocalization of $\pi$ electrons/lone pairs. $+R$ (electron donating): $-OH, -OR, -NH_2, -X$. $-R$ (electron withdrawing): $-NO_2, -CHO, -COOH, -CN$. Hyperconjugation: Delocalization of $\sigma$ electrons of C-H bond with adjacent $\pi$ system or empty p-orbital. Stabilizes carbocations, free radicals, alkenes. Acidic/Basic Strength: Acids: Factors stabilizing conjugate base increase acidity (e.g., -I, -R groups). Bases: Factors increasing electron density on N/O increase basicity (e.g., +I, +R groups). Purification Techniques: Distillation, Crystallization, Sublimation, Chromatography. 20. Hydrocarbons Alkanes (C-C single bonds, saturated, $C_nH_{2n+2}$): Preparation: Wurtz reaction, Reduction of alkyl halides, Hydrogenation of alkenes/alkynes. Reactions: Free Radical Halogenation: $CH_4 + Cl_2 \xrightarrow{hv} CH_3Cl + HCl$. (Chain initiation, propagation, termination). Combustion. Alkenes (C=C double bond, unsaturated, $C_nH_{2n}$): Preparation: Dehydration of alcohols, Dehydrohalogenation of alkyl halides (Saytzeff's rule), Partial hydrogenation of alkynes. Reactions (Electrophilic Addition): Hydrogenation: $R-CH=CH-R' + H_2 \xrightarrow{Ni/Pt/Pd} R-CH_2-CH_2-R'$. Halogenation: $R-CH=CH-R' + Br_2 \to R-CH(Br)-CH(Br)-R'$. (Anti addition). Hydrohalogenation: $R-CH=CH_2 + HBr \to R-CH(Br)-CH_3$ (Markovnikov's Rule: H adds to C with more H's). Anti-Markovnikov's: In presence of peroxide, $HBr$ adds opposite to Markovnikov's rule. Hydration: $R-CH=CH_2 + H_2O \xrightarrow{H^+} R-CH(OH)-CH_3$ (Markovnikov's). Ozonolysis: $R-CH=CH-R' \xrightarrow{O_3, Zn/H_2O} RCHO + R'CHO$. (Cleavage of C=C). Baeyer's Test (Cold, dil. alkaline $KMnO_4$): Forms diols (vicinal). Pink color decolorizes. Alkynes (C$\equiv$C triple bond, unsaturated, $C_nH_{2n-2}$): Preparation: Dehydrohalogenation of vicinal dihalides. Reactions: Acidic Nature of Terminal Alkynes: $R-C \equiv C-H + Na \to R-C \equiv C^-Na^+ + 1/2 H_2$. (React with Na, NaNH$_2$, ammoniacal $AgNO_3$, $Cu_2Cl_2$). Hydrogenation: Complete: $R-C \equiv C-R' + 2H_2 \xrightarrow{Ni/Pt/Pd} R-CH_2-CH_2-R'$. Partial (cis-alkene): $R-C \equiv C-R' + H_2 \xrightarrow{\text{Lindlar's catalyst}} \text{cis-alkene}$. Partial (trans-alkene): $R-C \equiv C-R' + Na/\text{liq. } NH_3 \to \text{trans-alkene (Birch Reduction)}$. Electrophilic Addition: Similar to alkenes, but two additions. Cyclic Polymerization: $3 C_2H_2 \xrightarrow{\text{Red hot iron tube}} C_6H_6$. Aromatic Hydrocarbons (Benzene): Aromaticity: Planar, cyclic, fully conjugated, $(4n+2)\pi$ electrons (Huckel's Rule). Preparation: Aromatization of alkanes, Decarboxylation of aromatic acids, Cyclic polymerization of ethyne. Reactions (Electrophilic Aromatic Substitution - EAS): Halogenation: $C_6H_6 + Cl_2 \xrightarrow{FeCl_3} C_6H_5Cl + HCl$. Nitration: $C_6H_6 + HNO_3 \xrightarrow{\text{conc. } H_2SO_4} C_6H_5NO_2 + H_2O$. Sulfonation: $C_6H_6 + H_2SO_4 \xrightarrow{\Delta} C_6H_5SO_3H + H_2O$. Friedel-Crafts Alkylation: $C_6H_6 + RCl \xrightarrow{AlCl_3} C_6H_5R + HCl$. (Rearrangement possible). Friedel-Crafts Acylation: $C_6H_6 + RCOCl \xrightarrow{AlCl_3} C_6H_5COR + HCl$. Directing Groups in EAS: Ortho-para directing, activating: $-OH, -NH_2, -OR, -CH_3$. Meta directing, deactivating: $-NO_2, -COOH, -CHO, -CN$. Halogens: Ortho-para directing, deactivating. 21. Organic Compounds Containing Halogens (Haloalkanes and Haloarenes) Preparation: From Alcohols: $ROH + HX \to RX + H_2O$. $ROH + PCl_5 \to RCl + POCl_3 + HCl$. From Hydrocarbons: Free radical halogenation (alkanes), Addition of HX/X$_2$ (alkenes), EAS (arenes). Halogen Exchange: Finkelstein reaction ($RI \leftarrow RCl + NaI$), Swarts reaction ($RF \leftarrow RBr + AgF$). Reactions of Haloalkanes: Nucleophilic Substitution ($S_N1, S_N2$): $S_N1$: Unimolecular, 2 steps, carbocation intermediate. Rate $\propto [RX]$. Order of reactivity: $3^\circ > 2^\circ > 1^\circ$. Racemization. $S_N2$: Bimolecular, 1 step (concerted), transition state. Rate $\propto [RX][Nu^-]$. Order of reactivity: $1^\circ > 2^\circ > 3^\circ$. Inversion of configuration (Walden inversion). Nucleophiles: $OH^-, CN^-, OR^-, NH_3, H_2O$, etc. Elimination (Dehydrohalogenation) (E1, E2): Forms alkenes. Saytzeff's Rule: Major product is the more substituted alkene. $E1$: 2 steps, carbocation intermediate. $3^\circ > 2^\circ > 1^\circ$. $E2$: 1 step, concerted. $3^\circ > 2^\circ > 1^\circ$. Reaction with Metals: Grignard Reagent: $RX + Mg \xrightarrow{\text{dry ether}} RMgX$. (Highly reactive, nucleophilic carbon). Wurtz Reaction: $2RX + 2Na \xrightarrow{\text{dry ether}} R-R + 2NaX$. (For symmetrical alkanes). Reactions of Haloarenes: Less reactive towards nucleophilic substitution due to resonance stabilization and $sp^2$ carbon. Nucleophilic Substitution: Only under harsh conditions (e.g., Dow's process for phenol). Activated by electron-withdrawing groups at o/p positions. Electrophilic Substitution: Halogens are ortho-para directing but deactivating. Wurtz-Fittig Reaction: $Ar-X + R-X + 2Na \to Ar-R + 2NaX$. Fittig Reaction: $2Ar-X + 2Na \to Ar-Ar + 2NaX$. 22. Organic Compounds Containing Oxygen (Alcohols, Phenols, Ethers, Aldehydes, Ketones, Carboxylic Acids) 22.1 Alcohols, Phenols, Ethers Alcohols ($R-OH$): Preparation: Hydration of alkenes, Reduction of aldehydes/ketones/carboxylic acids, Grignard reagents. Acidity: $1^\circ > 2^\circ > 3^\circ$ (due to +I effect). Less acidic than water. Reactions: Dehydration: $R-CH_2-CH_2-OH \xrightarrow{\text{conc. } H_2SO_4, \Delta} R-CH=CH_2$. (E1 mechanism, Saytzeff's rule). Oxidation: $1^\circ \text{ alcohol} \xrightarrow{PCC} \text{aldehyde} \xrightarrow{K_2Cr_2O_7} \text{carboxylic acid}$. $2^\circ \text{ alcohol} \xrightarrow{PCC} \text{ketone}$. $3^\circ \text{ alcohol}$ resists oxidation. Esterification: $R-OH + R'-COOH \xrightarrow{H^+} R'-COOR + H_2O$. Phenols ($Ar-OH$): Preparation: Dow's process (chlorobenzene), Cumene process. Acidity: More acidic than alcohols due to resonance stabilization of phenoxide ion. Electron-withdrawing groups (e.g., $-NO_2$) increase acidity at o/p positions. Reactions (EAS): -OH is o/p directing, activating. Nitration: Phenol $\xrightarrow{\text{dil. } HNO_3} \text{o- and p-nitrophenol}$. $\xrightarrow{\text{conc. } HNO_3} \text{picric acid (2,4,6-trinitrophenol)}$. Halogenation: Phenol $\xrightarrow{Br_2/CS_2} \text{o- and p-bromophenol}$. $\xrightarrow{Br_2/H_2O} \text{2,4,6-tribromophenol}$. Kolbe's Reaction: Phenol $\xrightarrow{NaOH} \text{Sodium Phenoxide} \xrightarrow{CO_2, H^+} \text{Salicylic Acid}$. Reimer-Tiemann Reaction: Phenol $\xrightarrow{CHCl_3, NaOH} \text{Salicylaldehyde}$. Reaction with Zn dust: Phenol $\xrightarrow{Zn \text{ dust}} \text{Benzene}$. OH Phenol Ethers ($R-O-R'$): Preparation: Williamson Ether Synthesis ($R-X + NaOR' \to R-O-R'$). (Best for $1^\circ$ alkyl halides). Reactions: Cleavage by strong acids (HI, HBr). $R-O-R' + HI \to RI + R'OH \xrightarrow{HI} RI + R'I$. 22.2 Aldehydes and Ketones ($R-CHO, R-CO-R'$, Carbonyl Compounds) Preparation: Oxidation of alcohols. Ozonolysis of alkenes. Rosenmund Reduction: $R-COCl \xrightarrow{H_2, Pd/BaSO_4} R-CHO$. Stephen Reaction: $R-CN \xrightarrow{SnCl_2/HCl} R-CH=NH \xrightarrow{H_3O^+} R-CHO$. Etard Reaction: Toluene $\xrightarrow{CrO_2Cl_2} \text{benzaldehyde}$. Gattermann-Koch Reaction: Benzene $\xrightarrow{CO, HCl, AlCl_3} \text{benzaldehyde}$. Reactions (Nucleophilic Addition): Carbonyl carbon is electrophilic. Addition of HCN: Forms cyanohydrins. Addition of NaHSO$_3$: Forms bisulfite addition product. Addition of Grignard Reagent: Formaldehyde $\to 1^\circ$ alcohol. Aldehyde $\to 2^\circ$ alcohol. Ketone $\to 3^\circ$ alcohol. Addition of Alcohols: Forms hemiacetals/acetals (from aldehydes), hemiketals/ketals (from ketones). Addition of Ammonia Derivatives: Forms imines, oximes, hydrazones, semicarbazones. Reduction: To Alcohols: $\xrightarrow{LiAlH_4 \text{ or } NaBH_4} R-CH_2-OH$. To Hydrocarbons: Clemmensen Reduction: $\xrightarrow{Zn-Hg/\text{conc. } HCl} R-CH_2-R'$. Wolff-Kishner Reduction: $\xrightarrow{NH_2NH_2, KOH/\text{ethylene glycol}} R-CH_2-R'$. Oxidation: Aldehydes easily oxidized to carboxylic acids. Ketones resist oxidation. Tollens' Test: Aldehydes $\xrightarrow{Ag(NH_3)_2^+OH^-} \text{silver mirror}$. Fehling's Test: Aldehydes $\xrightarrow{\text{Fehling's soln.}} \text{red ppt. of } Cu_2O$. Reactions involving $\alpha$-Hydrogen: Aldol Condensation: Aldehydes/ketones with $\alpha$-H react in presence of dilute base to form $\beta$-hydroxy aldehydes/ketones, then $\alpha,\beta$-unsaturated carbonyls. Cannizzaro Reaction: Aldehydes without $\alpha$-H react in presence of conc. base to give alcohol and carboxylic acid salt (disproportionation). (e.g., HCHO, $C_6H_5CHO$). Haloform Reaction (Iodoform Test): Compounds with $CH_3CO-$ or $CH_3CH(OH)-$ groups react with $X_2/NaOH$ to give haloform ($CHX_3$) and carboxylate salt. R C H O R C R' O Aldehyde (left) and Ketone (right) 22.3 Carboxylic Acids ($R-COOH$) Preparation: Oxidation of $1^\circ$ alcohols/aldehydes, Hydrolysis of nitriles/esters, Grignard reagents with $CO_2$. Acidity: More acidic than alcohols/phenols. Electron-withdrawing groups increase acidity. Formic acid ($HCOOH$) is strongest simple carboxylic acid. Reactions: Esterification: $R-COOH + R'-OH \xrightarrow{H^+} R-COOR' + H_2O$. Reduction: $R-COOH \xrightarrow{LiAlH_4} R-CH_2-OH$. (Strong reducing agent needed). Decarboxylation: $R-COONa \xrightarrow{NaOH/CaO, \Delta} R-H$. HVZ Reaction (Hell-Volhard-Zelinsky): $R-CH_2-COOH \xrightarrow{X_2/\text{red P}} R-CH(X)-COOH$. (Halogenation at $\alpha$-carbon). Formation of Derivatives: Acid chlorides ($RCOCl$), Anhydrides ($(RCO)_2O$), Esters ($RCOOR'$), Amides ($RCONH_2$). R C OH O Carboxylic Acid 23. Organic Compounds Containing Nitrogen (Amines, Diazonium Salts) Amines ($R-NH_2, R_2NH, R_3N$): Preparation: Reduction of Nitro Compounds: $R-NO_2 \xrightarrow{Sn/HCl \text{ or } H_2/Pd} R-NH_2$. Ammonolysis of Alkyl Halides: $RX + NH_3 \to RNH_2 + R_2NH + R_3N + R_4N^+X^-$. (Mixture of $1^\circ, 2^\circ, 3^\circ$ amines and quaternary salt). Gabriel Phthalimide Synthesis: For $1^\circ$ aliphatic amines. Phthalimide $\xrightarrow{KOH} \text{K-Phthalimide} \xrightarrow{RX} \text{N-alkylphthalimide} \xrightarrow{H_3O^+ \text{ or } NaOH} R-NH_2$. Hofmann Bromamide Degradation: $R-CONH_2 \xrightarrow{Br_2/NaOH} R-NH_2 + Na_2CO_3 + NaBr + H_2O$. (Decreases carbon by one). Reduction of Nitriles/Amides: $R-CN \xrightarrow{LiAlH_4} R-CH_2-NH_2$. $R-CONH_2 \xrightarrow{LiAlH_4} R-CH_2-NH_2$. Basicity: Amines are basic due to lone pair on N. Aqueous: $2^\circ > 1^\circ > 3^\circ > NH_3$ (for methyl amines). $2^\circ > 3^\circ > 1^\circ > NH_3$ (for ethyl amines). (Due to solvation +I effect). Gaseous: $3^\circ > 2^\circ > 1^\circ > NH_3$ (only +I effect). Aromatic amines are less basic than aliphatic amines (resonance effect). Reactions: Acylation: $R-NH_2 + CH_3COCl \to R-NHCOCH_3$. Carbylamine Reaction (Isocyanide Test): $1^\circ$ amine $\xrightarrow{CHCl_3, KOH} R-NC$ (foul smelling). Hinsberg Test: Distinguishes $1^\circ, 2^\circ, 3^\circ$ amines using benzenesulfonyl chloride. $1^\circ$: Forms product soluble in KOH. $2^\circ$: Forms product insoluble in KOH. $3^\circ$: No reaction. Diazotisation: $1^\circ$ aromatic amine $\xrightarrow{NaNO_2/HCl (0-5^\circ C)} Ar-N_2^+Cl^-$. Diazonium Salts ($Ar-N_2^+X^-$): Reactions: Sandmeyer Reaction: $Ar-N_2^+Cl^- \xrightarrow{CuCl/HCl \text{ or } CuBr/HBr \text{ or } CuCN/KCN} Ar-Cl/Ar-Br/Ar-CN$. Gattermann Reaction: $Ar-N_2^+Cl^- \xrightarrow{Cu/HCl \text{ or } Cu/HBr} Ar-Cl/Ar-Br$. Balz-Schiemann Reaction: $Ar-N_2^+Cl^- \xrightarrow{HBF_4, \Delta} Ar-F$. Reduction: $Ar-N_2^+Cl^- \xrightarrow{H_3PO_2 \text{ or } EtOH} Ar-H$. Coupling Reactions: With phenols or anilines to form azo dyes (e.g., p-hydroxyazobenzene). 24. Biomolecules Carbohydrates: Polyhydroxy aldehydes or ketones. Monosaccharides: Glucose, Fructose (non-hydrolyzable). Glucose (an aldohexose): Forms pyranose ring. $\alpha, \beta$ anomers. Fructose (a ketohexose): Forms furanose ring. Disaccharides: Sucrose (glucose + fructose, non-reducing), Maltose (glucose + glucose, reducing), Lactose (glucose + galactose, reducing). Polysaccharides: Starch (amylose + amylopectin), Cellulose, Glycogen. Proteins: Polymers of $\alpha$-amino acids linked by peptide bonds. Amino Acids: Zwitterionic form. Essential vs. Non-essential. Peptide Bond: $-CO-NH-$. Structure: Primary: Sequence of amino acids. Secondary: $\alpha$-helix, $\beta$-pleated sheet (H-bonding). Tertiary: 3D folding (disulfide, H-bonding, ionic, van der Waals). Quaternary: Arrangement of multiple polypeptide units. Denaturation: Loss of 2$^\circ$, 3$^\circ$, 4$^\circ$ structure (by heat, pH change). Primary structure remains. Vitamins: Organic compounds required in small amounts. Fat-soluble: A, D, E, K. Water-soluble: B complex, C. Nucleic Acids: DNA (Deoxyribonucleic Acid) and RNA (Ribonucleic Acid). Monomer: Nucleotide (Nitrogenous base + Pentose sugar + Phosphate group). Bases: Purines (Adenine, Guanine), Pyrimidines (Cytosine, Thymine (DNA), Uracil (RNA)). DNA: Double helix (Watson-Crick), A-T (2 H-bonds), G-C (3 H-bonds). RNA: Single stranded. 25. Polymers Classification: Source: Natural (starch, cellulose, rubber), Synthetic (nylon, polythene). Structure: Linear, Branched, Cross-linked. Molecular Forces: Elastomers, Fibers, Thermoplastics, Thermosetting plastics. Synthesis: Addition Polymerization: Unsaturated monomers add to each other without loss of small molecules (e.g., polythene from ethene, PVC from vinyl chloride). Condensation Polymerization: Monomers combine with elimination of small molecules (e.g., H$_2$O, HCl) (e.g., Nylon 6,6, Terylene). Important Polymers: Polythene: Low Density (LDPE), High Density (HDPE). PVC (Polyvinyl Chloride): From vinyl chloride. Teflon (Polytetrafluoroethene): From tetrafluoroethene. (Non-stick cookware). Nylon 6,6: Hexamethylenediamine + Adipic acid. Nylon 6: From Caprolactam. Terylene (Dacron): Ethylene glycol + Terephthalic acid. Bakelite: Phenol + Formaldehyde (thermosetting). Natural Rubber: Isoprene (cis-1,4-polyisoprene). Vulcanization: Heating with sulfur to improve properties. Biodegradable Polymers: PHBV (Poly-$\beta$-hydroxybutyrate-co-$\beta$-hydroxyvalerate). 26. Chemistry in Everyday Life Drugs: Antacids: Neutralize excess acid in stomach (e.g., $Mg(OH)_2, Al(OH)_3$). Antihistamines: Counteract histamine, anti-allergic (e.g., Cimetidine, Terfenadine). Tranquilizers: Reduce anxiety, induce sense of well-being (e.g., Equanil). Analgesics: Reduce pain (e.g., Aspirin, Paracetamol). Narcotic (morphine) vs. Non-narcotic. Antipyretics: Reduce fever (e.g., Aspirin, Paracetamol). Antiseptics: Applied to living tissues (e.g., Dettol, Bithional). Disinfectants: Applied to non-living objects (e.g., Chlorine, $SO_2$). Antibiotics: Inhibit growth or kill microorganisms (e.g., Penicillin, Chloramphenicol). Bactericidal vs. Bacteriostatic. Food Additives: Artificial Sweetening Agents: Saccharin, Aspartame, Sucralose, Alitame. Food Preservatives: Sodium benzoate, Sodium metabisulfite. Antioxidants: BHT, BHA. Cleansing Agents: Soaps: Sodium/potassium salts of long chain fatty acids. Not effective in hard water. Detergents: Synthetic, effective in hard water. Anionic, Cationic, Non-ionic. 27. Environmental Chemistry Atmospheric Pollution: Tropospheric Pollution: Gaseous Pollutants: Oxides of Sulfur ($SO_2, SO_3$), Nitrogen ($NO_x$), Carbon ($CO, CO_2$), Hydrocarbons. Particulate Pollutants: Dust, smoke, smog (photochemical smog). Acid Rain: $SO_2, NO_x$ react with water to form $H_2SO_4, HNO_3$. pH Stratospheric Pollution: Ozone Layer Depletion: By CFCs ($CF_2Cl_2 \xrightarrow{hv} \cdot Cl + \cdot CF_2Cl$, $\cdot Cl + O_3 \to ClO^\cdot + O_2$). Protects from UV radiation. Water Pollution: Sources: Pathogens, organic wastes, chemical pollutants. BOD (Biochemical Oxygen Demand): Amount of oxygen required by bacteria to decompose organic matter. High BOD indicates high pollution. Eutrophication: Nutrient enrichment of water bodies leading to algal blooms, oxygen depletion, and death of aquatic life. Soil Pollution: Pesticides, industrial wastes. Green Chemistry: Design of chemical products and processes that reduce or eliminate the use and generation of hazardous substances. (e.g., use of water/supercritical $CO_2$ as solvent instead of organic solvents).