Chemistry HSC Quick Reference

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1. Solid State Crystalline Solids: Regular, periodic arrangement of particles; sharp melting points; anisotropic (except cubic systems). Amorphous Solids: Random particle arrangement; melt gradually; isotropic; supercooled liquids. Isomorphism: Two or more substances with the same crystal structure and atomic ratio. Polymorphism: Single substance existing in two or more crystalline forms under different conditions. Allotropy is polymorphism in elements. Types of Crystalline Solids: Ionic Solids: Ions held by electrostatic forces; hard, brittle; high melting point; non-conductors in solid state, but good conductors when molten or dissolved. Covalent Network Solids: Atoms linked by continuous covalent bonds forming giant molecules; very hard; high melting/boiling points; poor conductors (except graphite). Molecular Solids: Molecules held by weak intermolecular forces (dipole-dipole, London, H-bonds); soft; low melting points; poor conductors. Metallic Solids: Metal ions in a "sea of mobile electrons" (metallic bonds); malleable, ductile; good electrical/thermal conductors. Crystal Lattice: Geometrical arrangement of points in 3D periodic array. Basis: Constituent particles attached to lattice points. Unit Cell: Smallest repeating structural unit of a crystalline solid. Coordination Number: Number of nearest neighbors touching a given particle. Packing Efficiency: Percentage of total space occupied by particles in a unit cell. Simple Cubic (sc): 52.4% Body-Centered Cubic (bcc): 68% Face-Centered Cubic (fcc)/Hexagonal Close-Packed (hcp): 74% Crystal Defects (Imperfections): Irregularities in particle arrangement. Point Defects: Irregularities at lattice points. Stoichiometric Defects: Stoichiometry remains unchanged. Vacancy Defect: Missing particle from lattice site. Self-Interstitial Defect: Particle occupies interstitial site. Schottky Defect: Equal numbers of cations and anions missing (density decreases). Frenkel Defect: Ion displaced from lattice site to interstitial site (density unchanged). Impurity Defects: Foreign atoms present. Substitutional: Impurity at host lattice site. Interstitial: Impurity at interstitial site. Nonstoichiometric Defects: Stoichiometry changes. Metal Deficiency: Missing metal ions, balanced by higher oxidation state of other metal ions. Metal Excess: Extra metal atoms/ions in interstitial sites or anion vacancies (F-centers). Band Theory: Explains electrical conductivity in solids. Conduction Band: Highest energy band, partially occupied or vacant; electrons are mobile. Valence Band: Lower energy band, electrons tightly bound. Band Gap: Energy difference between valence and conduction bands. Conductors: Overlapping or partially filled bands (small/zero band gap). Insulators: Large band gap. Semiconductors: Small band gap. Extrinsic Semiconductors (Doping): n-type: Doped with Group 15 elements (excess electrons). p-type: Doped with Group 13 elements (electron vacancies/holes). Magnetic Properties: Diamagnetic: All electrons paired; weakly repelled by magnetic field. Paramagnetic: Unpaired electrons; weakly attracted by magnetic field (temporary). Ferromagnetic: Large number of unpaired electrons; strongly attracted (permanent magnetism possible). 2. Solutions Solution: Homogeneous mixture of two or more pure substances. Solute (minor component), Solvent (major component). Saturated Solution: Contains maximum amount of solute at a given temperature. Dynamic equilibrium between dissolution and crystallization. Supersaturated Solution: Contains greater than equilibrium amount of solute; unstable. Solubility: Amount of solute per unit volume of saturated solution at a specific temperature (mol L$^{-1}$). Factors Affecting Solubility: Nature of Solute and Solvent: "Like dissolves like." Polar solutes in polar solvents, nonpolar in nonpolar. Temperature: Endothermic dissolution: Solubility increases with temperature (Le-Chatelier's principle). Exothermic dissolution: Solubility decreases with temperature. Gases in liquids: Solubility usually decreases with increasing temperature (condensation is exothermic). Pressure (for gases in liquids): Solubility increases with increasing pressure. Henry's Law: Solubility of a gas in a liquid is directly proportional to the partial pressure of the gas over the solution. $S = K_H P$. Raoult's Law (Volatile components): Partial vapor pressure of a volatile component in a solution is equal to the vapor pressure of the pure component multiplied by its mole fraction in the solution. $P_1 = x_1 P_1^0$. Ideal Solutions: Obey Raoult's law over entire concentration range. $\Delta_{mix}H = 0$ (no heat change on mixing). $\Delta_{mix}V = 0$ (no volume change on mixing). Solvent-solute, solute-solute, solvent-solvent interactions are comparable. Nonideal Solutions: Do not obey Raoult's law. Positive Deviations: Vapour pressure higher than ideal; weaker solute-solvent interactions. Negative Deviations: Vapour pressure lower than ideal; stronger solute-solvent interactions. Colligative Properties: Properties of solutions that depend on the number of solute particles, not their nature. Vapour pressure lowering ($\Delta P$). Boiling point elevation ($\Delta T_b$). Freezing point depression ($\Delta T_f$). Osmotic pressure ($\pi$). Vapour Pressure Lowering: When a nonvolatile solute is added to a solvent, its vapor pressure decreases. Relative lowering of vapor pressure is equal to the mole fraction of solute: $\frac{\Delta P}{P_1^0} = x_2$. Boiling Point Elevation: $\Delta T_b = T_b - T_b^0$. For dilute solutions, $\Delta T_b = K_b m$, where $K_b$ is the ebullioscopic constant. Freezing Point Depression: $\Delta T_f = T_f^0 - T_f$. For dilute solutions, $\Delta T_f = K_f m$, where $K_f$ is the cryoscopic constant. Osmosis: Net spontaneous flow of solvent molecules through a semipermeable membrane from a region of higher solvent concentration to lower solvent concentration. Osmotic Pressure ($\pi$): Hydrostatic pressure that stops osmosis. For dilute solutions, $\pi = MRT$. Isotonic Solutions: Solutions with the same osmotic pressure. Hypertonic Solution: Higher osmotic pressure. Hypotonic Solution: Lower osmotic pressure. Reverse Osmosis: Solvent flow from solution to pure solvent by applying pressure greater than osmotic pressure. van't Hoff Factor ($i$): Ratio of observed colligative property to theoretical colligative property for electrolytes. $i = \frac{\text{actual moles of particles in solution}}{\text{moles of formula units dissolved}}$. For electrolytes, colligative properties are modified by $i$. Degree of Dissociation ($\alpha$): $\alpha = \frac{i-1}{n-1}$, where $n$ is the number of ions produced per formula unit. 3. Ionic Equilibria Ionic Equilibrium: Equilibrium between ions and unionized molecules in solution. Electrolytes: Substances that produce ions when dissolved in water. Strong Electrolytes: Ionize completely or almost completely (strong acids, strong bases, most salts). Weak Electrolytes: Dissociate to a small extent (weak acids, weak bases). Degree of Dissociation ($\alpha$): Fraction of total moles of electrolyte that dissociates into ions. Arrhenius Theory: Acid: Substance giving H$^+$ ions in aqueous solution. Base: Substance giving OH$^-$ ions in aqueous solution. Brønsted-Lowry Theory: Acid: Proton donor. Base: Proton acceptor. Conjugate Acid-Base Pair: Acid and base differing by a proton. Lewis Theory: Acid: Electron pair acceptor. Base: Electron pair donor. Amphoteric Nature of Water: Water can act as both an acid and a base. Dissociation Constant: Weak Acid ($K_a$): $K_a = \frac{[H^+][A^-]}{[HA]}$ Weak Base ($K_b$): $K_b = \frac{[B^+][OH^-]}{[BOH]}$ Ostwald's Dilution Law: For weak electrolytes, $\alpha = \sqrt{\frac{K_a}{c}}$ (for acids) or $\alpha = \sqrt{\frac{K_b}{c}}$ (for bases). Degree of dissociation increases with dilution. Autoionization of Water: Water dissociates to H$_3$O$^+$ and OH$^-$. Ionic Product of Water ($K_w$): $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at 298 K. pH Scale: pH = -log$_{10}$[H$^+$]. pOH = -log$_{10}$[OH$^-$]. pH + pOH = 14. Acidic: pH 7. Hydrolysis of Salts: Reaction of salt ions with water to produce acidity or alkalinity. Strong acid + Strong base salt: Neutral solution (no hydrolysis). Strong acid + Weak base salt: Acidic solution (cation hydrolysis). Weak acid + Strong base salt: Basic solution (anion hydrolysis). Weak acid + Weak base salt: pH depends on relative strengths of acid/base formed. Buffer Solutions: Solutions that resist drastic changes in pH upon addition of small amounts of strong acid or base. Acidic Buffer: Weak acid + its salt with a strong base (e.g., CH$_3$COOH/CH$_3$COONa). pH = p$K_a$ + log$_{10}\frac{[\text{salt}]}{[\text{acid}]}$. Basic Buffer: Weak base + its salt with a strong acid (e.g., NH$_4$OH/NH$_4$Cl). pOH = p$K_b$ + log$_{10}\frac{[\text{salt}]}{[\text{base}]}$. Solubility Product ($K_{sp}$): For a sparingly soluble salt, product of equilibrium concentrations of its ions, each raised to the power of its stoichiometric coefficient. $IP = K_{sp}$: Saturated solution, equilibrium. $IP > K_{sp}$: Supersaturated, precipitation occurs. $IP Common Ion Effect: Ionization of a weak electrolyte is suppressed in the presence of a strong electrolyte containing a common ion. It also decreases the solubility of sparingly soluble salts. 4. Chemical Thermodynamics Thermodynamics: Study of energy changes in physical and chemical transformations. System: Part of the universe under thermodynamic investigation. Surroundings: All other parts of the universe outside the system. Types of Systems: Open System: Exchanges both energy and matter with surroundings. Closed System: Exchanges energy but not matter with surroundings. Isolated System: Exchanges neither energy nor matter with surroundings. Properties of System: Extensive Property: Depends on amount of matter (e.g., mass, volume, internal energy). Intensive Property: Independent of amount of matter (e.g., pressure, temperature, density). State Functions: Properties that depend only on the state of the system, not on the path taken to reach that state (e.g., P, V, T, U, H, S, G). Path Functions: Properties that depend on the path taken (e.g., work (W), heat (Q)). Thermodynamic Equilibrium: State where system properties do not vary with time. Processes: Transition from one equilibrium state to another. Isothermal Process: Temperature (T) is constant ($\Delta T = 0$, $\Delta U = 0$). Isobaric Process: Pressure (P) is constant ($\Delta P = 0$). Isochoric Process: Volume (V) is constant ($\Delta V = 0$). Adiabatic Process: No heat exchange (Q = 0). Reversible Process: Occurs infinitely slowly through infinite steps; driving and opposing forces differ infinitesimally; can be reversed by infinitesimal change. Work (W): Energy exchange due to displacement against a force. In thermodynamics, primarily PV work. $W = -P_{ext} \Delta V$. Work done BY system: W Work done ON system: W > 0. Free Expansion (in vacuum): $P_{ext} = 0$, so $W = 0$. Heat (Q): Energy exchange due to temperature difference. Heat absorbed BY system: Q > 0. Heat released BY system: Q First Law of Thermodynamics: Energy is conserved. $\Delta U = Q + W$. Isothermal: $Q = -W$. Adiabatic: $\Delta U = W$. Isochoric: $\Delta U = Q_V$. Isobaric: $Q_p = \Delta U + P\Delta V$. Enthalpy (H): $H = U + PV$. Change in enthalpy: $\Delta H = \Delta U + \Delta(PV)$. At constant pressure, $\Delta H = \Delta U + P\Delta V = Q_p$. Relationship between $\Delta H$ and $\Delta U$ for Reactions: $\Delta H = \Delta U + \Delta n_g RT$, where $\Delta n_g$ is change in moles of gaseous species. Work Done in Chemical Reaction: $W = -\Delta n_g RT$. Enthalpies of Phase Transition: Fusion ($\Delta_{fus}H$): Solid to liquid. Vaporization ($\Delta_{vap}H$): Liquid to gas. Sublimation ($\Delta_{sub}H$): Solid to gas. $\Delta_{sub}H = \Delta_{fus}H + \Delta_{vap}H$. Enthalpies for Atomic/Molecular Change: Ionization ($\Delta_{ion}H$): Removal of electron from gaseous atom. Electron Gain ($\Delta_{eg}H$): Addition of electron to gaseous atom. Atomization ($\Delta_{atom}H$): Dissociation of gaseous substance into atoms. Solution ($\Delta_{soln}H$): Dissolution of one mole of substance in solvent. $\Delta_{soln}H = \Delta_LH + \Delta_{hyd}H$. Thermochemistry: Study of enthalpy changes in chemical reactions. Enthalpy of Reaction ($\Delta_rH$): $\Delta_rH = \sum H_{products} - \sum H_{reactants}$. Exothermic Reaction: $\Delta_rH Endothermic Reaction: $\Delta_rH > 0$ (heat absorbed). Standard Enthalpy of Reaction ($\Delta_rH^0$): Enthalpy change under standard conditions (1 bar, 298 K, 1 M for solutions). Standard Enthalpy of Formation ($\Delta_fH^0$): Enthalpy change when one mole of compound is formed from its elements in their standard states ($\Delta_fH^0$ for element = 0). Standard Enthalpy of Combustion ($\Delta_cH^0$): Enthalpy change when one mole of substance is completely oxidized. Bond Enthalpy: Energy required to break a particular covalent bond in one mole of gaseous molecule. $\Delta_rH^0 = \sum \Delta H^0_{reactants} - \sum \Delta H^0_{products}$. Hess's Law of Constant Heat Summation: Overall enthalpy change for a reaction is the sum of enthalpy changes of individual steps. Enthalpy is a state function. Spontaneous Process: Occurs naturally without external influence. Proceeds in one direction till equilibrium. Can be rapid or slow. Exothermicity is not the sole criterion for spontaneity. Entropy (S): Measure of molecular disorder or randomness. $\Delta S = \frac{Q_{rev}}{T}$. Second Law of Thermodynamics: Total entropy of a system and its surroundings increases in a spontaneous process. $\Delta S_{total} = \Delta S_{sys} + \Delta S_{surr} > 0$. $\Delta S_{total} > 0$: Spontaneous. $\Delta S_{total} $\Delta S_{total} = 0$: At equilibrium. Gibbs Energy (G): $G = H - TS$. Change at constant T and P: $\Delta G = \Delta H - T\Delta S$. Gibbs Energy and Spontaneity: $\Delta G = -T\Delta S_{total}$. $\Delta G $\Delta G > 0$: Nonspontaneous. $\Delta G = 0$: At equilibrium. Temperature of Equilibrium: $T = \frac{\Delta H}{\Delta S}$ (when $\Delta G = 0$). Gibbs Function and Equilibrium Constant: $\Delta G = \Delta G^0 + RT \ln Q$. At equilibrium, $\Delta G = 0$ and $Q = K$. So, $\Delta G^0 = -RT \ln K = -2.303 RT \log_{10} K$. 5. Electrochemistry Electrochemistry: Study of interconversion of chemical and electrical energy. Electric Conduction: Metallic Conduction: Flow of electrons through metals (electronic conductors). Electrolytic Conduction: Movement of ions through molten electrolytes or electrolyte solutions (ionic conductors). Accompanied by chemical change. Electrical Conductance (G): Reciprocal of resistance (R). $G = \frac{1}{R}$. Unit: Siemens (S) or $\Omega^{-1}$. Resistivity ($\rho$): Resistance of a conductor of unit length and unit cross-sectional area. $R = \rho \frac{l}{a}$. Unit: $\Omega$ m. Conductivity ($\kappa$): Reciprocal of resistivity. $k = \frac{1}{\rho} = G \frac{l}{a}$. Electrical conductance of unit volume of solution. Unit: S m$^{-1}$ or $\Omega^{-1}$ cm$^{-1}$. Molar Conductivity ($\Lambda$): Electrolytic conductivity divided by molar concentration. $\Lambda = \frac{k}{c}$. Unit: S m$^2$ mol$^{-1}$ or $\Omega^{-1}$ cm$^2$ mol$^{-1}$. Increases with dilution for strong electrolytes. Increases rapidly with dilution for weak electrolytes. Kohlrausch's Law of Independent Migration of Ions: At infinite dilution, each ion contributes independently to the total molar conductivity of an electrolyte. $\Lambda^0 = n_+\lambda_+^0 + n_-\lambda_-^0$. Degree of Dissociation ($\alpha$) of Weak Electrolytes: $\alpha = \frac{\Lambda_c}{\Lambda_0}$. Electrochemical Cell: Device where redox reactions generate electricity (galvanic) or electricity drives non-spontaneous reactions (electrolytic). Electrodes: Surfaces where oxidation/reduction occur. Cathode: Reduction occurs (gains electrons). Anode: Oxidation occurs (loses electrons). Electrolytic Cell: External electricity drives non-spontaneous reaction (electrolysis). Electrical energy to chemical energy. Anode is positive (+), cathode is negative (-). Galvanic (Voltaic) Cell: Spontaneous chemical reaction produces electricity. Chemical energy to electrical energy. Anode is negative (-), cathode is positive (+). Salt Bridge: Connects two half-cells in a galvanic cell, maintains electrical neutrality, completes circuit. Faraday's Laws of Electrolysis: Relate quantity of electricity to mass of substance produced. Quantity of electricity (Q) = Current (I) $\times$ Time (t). 1 Faraday (F) = 96500 C = charge of one mole of electrons. Electrode Potential: Potential difference established at the interface of a metal electrode and its ion solution due to half-reaction. Standard Electrode Potential ($E^0$): Electrode potential under standard conditions (1 M solution, 1 atm gas pressure, 25 $^0$C). Standard reduction potential (IUPAC convention). $E^0_{cell} = E^0_{cathode} - E^0_{anode}$. Nernst Equation: Relates cell potential to concentrations. $E_{cell} = E^0_{cell} - \frac{RT}{nF} \ln Q$ or $E_{cell} = E^0_{cell} - \frac{0.0592}{n} \log_{10} Q$ at 25 $^0$C. Gibbs Energy of Cell Reactions: $\Delta G = -nFE_{cell}$. Under standard conditions, $\Delta G^0 = -nFE^0_{cell}$. Standard Cell Potential and Equilibrium Constant (K): $E^0_{cell} = \frac{RT}{nF} \ln K = \frac{0.0592}{n} \log_{10} K$ at 25 $^0$C. Reference Electrodes: Electrodes with arbitrarily zero or exactly known potentials (e.g., Standard Hydrogen Electrode, SHE). Types of Voltaic Cells: Primary Cells: Non-rechargeable (e.g., dry cell). Secondary Cells: Rechargeable (e.g., lead storage battery, nickel-cadmium cell). Fuel Cells: Reactants continuously supplied to electrodes to produce electricity (e.g., H$_2$-O$_2$ fuel cell). High efficiency, non-polluting. Electrochemical Series: List of standard reduction potentials, arranged in decreasing order. Higher $E^0$: Stronger oxidizing agent. Lower $E^0$: Stronger reducing agent. Spontaneous redox reaction: $E^0_{cell} > 0$. 6. Chemical Kinetics Chemical Kinetics: Study of reaction rates and factors affecting them. Rate of Reaction: How rapidly reactants are consumed or products formed. Average Rate: $\frac{\Delta C}{\Delta t}$. Instantaneous Rate: $\frac{dC}{dt}$. For $aA + bB \rightarrow cC + dD$, Rate $= -\frac{1}{a}\frac{d[A]}{dt} = -\frac{1}{b}\frac{d[B]}{dt} = \frac{1}{c}\frac{d[C]}{dt} = \frac{1}{d}\frac{d[D]}{dt}$. Units: mol dm$^{-3}$ s$^{-1}$. Rate Law: Experimentally determined expression relating reaction rate to reactant concentrations. Rate $= k[A]^x[B]^y$. Rate Constant (k): Proportionality constant in rate law; independent of concentration, depends on temperature. Order of Reaction: Sum of powers of concentration terms in the rate law ($x+y$). Experimentally determined; can be integer, fraction, or zero. Molecularity of Reaction: Number of reactant molecules taking part in an elementary reaction. Always an integer. Elementary Reaction: Occurs in a single step. Complex Reaction: Occurs in a series of elementary steps. Rate Determining Step: The slowest step in a complex reaction, which determines the overall rate. Integrated Rate Laws: First Order Reaction: $A \rightarrow P$. $\ln \frac{[A]_t}{[A]_0} = -kt$ or $k = \frac{2.303}{t} \log_{10} \frac{[A]_0}{[A]_t}$. Units of k: s$^{-1}$. Half-life ($t_{1/2}$): Time for concentration to halve. $t_{1/2} = \frac{0.693}{k}$. Independent of initial concentration. Zero Order Reaction: $A \rightarrow P$. Rate is independent of reactant concentration. $kt = [A]_0 - [A]_t$. Units of k: mol dm$^{-3}$ s$^{-1}$. Half-life ($t_{1/2}$): $t_{1/2} = \frac{[A]_0}{2k}$. Proportional to initial concentration. Pseudo-First Order Reactions: Reactions that appear to be first order but are actually higher order due to one reactant being in large excess (e.g., hydrolysis of ester in excess water). Collision Theory: For a reaction to occur, reactant molecules must collide with: Sufficient Kinetic Energy: Equal to or greater than activation energy ($E_a$). Proper Orientation: Reacting groups must align correctly. Activation Energy ($E_a$): Minimum kinetic energy required for colliding molecules to react. Potential Energy Barrier: Energy barrier between reactants and products. Arrhenius Equation: Describes temperature dependence of rate constant. $k = A e^{-E_a/RT}$. A: Pre-exponential factor (frequency factor). $\log_{10} k = -\frac{E_a}{2.303R} \frac{1}{T} + \log_{10} A$. For two temperatures: $\log_{10} \frac{k_2}{k_1} = \frac{E_a}{2.303R} \left(\frac{T_2 - T_1}{T_1 T_2}\right)$. Effect of Catalyst: Increases reaction rate without being consumed. Provides an alternative reaction pathway with lower activation energy. 7. Elements of Groups 16, 17 and 18 7.1 Group 16 (Oxygen Family / Chalcogens) Elements: O, S, Se, Te, Po. Electronic Configuration: ns$^2$np$^4$. Occurrence: Oxygen (most abundant), Sulfur (sulfates, sulfides). Atomic Properties: Atomic/Ionic Radii: Increase down the group. Ionization Enthalpy: High, decreases down the group. Lower than Group 15 due to np$^4$ configuration. Electronegativity: Decreases down the group. Oxygen is highly electronegative. Electron Gain Enthalpy: Becomes less negative down the group. Oxygen has less negative than sulfur due to small size. Physical Properties: Oxygen (gas), others (solids). O, S (nonmetals); Se, Te (metalloids); Po (metal). Melting/boiling points increase down the group. All exhibit allotropy (e.g., O$_2$, O$_3$; Rhombic and Monoclinic S). Anomalous Behavior of Oxygen: Small size, high electronegativity, absence of d-orbitals. Diatomic (O$_2$) vs. polyatomic. Paramagnetic O$_2$. Oxidation states: -2, -1, +2 (in OF$_2$). Hydride (H$_2$O) is liquid. Max covalency 2 (rarely 4). Chemical Properties: Oxidation States: -2, +2, +4, +6. Oxygen common -2 (except OF$_2$, peroxides). Stability of +6 decreases, +4 increases down the group. Reactivity towards H: Form H$_2$E. Acidic character increases H$_2$O Reactivity towards O: Form EO$_2$ and EO$_3$. Acidic oxides. Reactivity towards Halogens: Form EX$_6$, EX$_4$, EX$_2$. Stability: fluorides > chlorides > bromides > iodides. Reactivity towards Metals: Form corresponding oxides/sulfides. Oxoacids of Sulfur: H$_2$SO$_3$, H$_2$SO$_4$, H$_2$S$_2$O$_7$, etc. Dioxygen (O$_2$): Preparation: Heating KClO$_3$, decomposition of metal oxides, H$_2$O$_2$; electrolysis of water; fractional distillation of liquid air. Properties: Colorless, odorless gas; sparingly soluble in water; paramagnetic. Reactions: With metals, nonmetals, some compounds (e.g., 2SO$_2$ + O$_2 \rightarrow$ 2SO$_3$). Ozone (O$_3$): Allotrope of oxygen. Preparation: Silent electric discharge through O$_2$. Properties: Pale blue gas; characteristic smell; diamagnetic. Reactions: Powerful oxidizing agent (O$_3 \rightarrow$ O$_2$ + O); bleaching agent; reduces peroxides. Ozone Depletion: Caused by NO, CFCs. Sulfur Dioxide (SO$_2$): Preparation: Burning S, Na$_2$SO$_3$ + H$_2$SO$_4$, roasting sulfides. Properties: Colorless gas, pungent smell; highly soluble in water (sulfurous acid); reducing agent (in presence of moisture). Structure: Angular, resonance hybrid. Sulfuric Acid (H$_2$SO$_4$): Preparation (Contact Process): S $\rightarrow$ SO$_2$; SO$_2 \rightarrow$ SO$_3$ (V$_2$O$_5$ catalyst); SO$_3$ + H$_2$SO$_4 \rightarrow$ H$_2$S$_2$O$_7$ (oleum); H$_2$S$_2$O$_7$ + H$_2$O $\rightarrow$ H$_2$SO$_4$. Properties: Strong acid, oxidizing agent (hot, conc.), dehydrating agent. 7.2 Group 17 (Halogen Family) Elements: F, Cl, Br, I, At. Electronic Configuration: ns$^2$np$^5$. Occurrence: Very reactive, found as compounds (fluorides, chlorides, bromides, iodides). Atomic Properties: Atomic/Ionic Radii: Increase down the group. Smallest in their periods. Ionization Enthalpy: High, decreases down the group. Electronegativity: Very high, decreases down the group. Fluorine is most electronegative. Electron Gain Enthalpy: Becomes less negative down the group. Fluorine has less negative than chlorine due to small size. Physical Properties: F$_2$, Cl$_2$ (gases); Br$_2$ (liquid); I$_2$ (solid). Colored (F$_2$ yellow, Cl$_2$ greenish-yellow, Br$_2$ red, I$_2$ violet). Bond dissociation enthalpy: Cl-Cl > Br-Br > F-F > I-I (F-F is low due to lone pair repulsion). Anomalous Behavior of Fluorine: Small size, high electronegativity, absence of d-orbitals, low F-F bond enthalpy. Higher IE, EN, electrode potential. Lower ionic/covalent radii, MP, BP, electron gain enthalpy. Most reactions exothermic. Forms only one oxoacid (HOF). HF is liquid (H-bonding). Chemical Properties: Oxidation States: -1 (all). Cl, Br, I also +1, +3, +5, +7 (due to empty d-orbitals). F only -1. Reactivity towards H: Form HX (hydrogen halides). Acidic strength: HF HCl > HBr > HI. Reactivity towards O: Form many unstable oxides. F forms OF$_2$, O$_2$F$_2$. Cl forms Cl$_2$O, ClO$_2$, Cl$_2$O$_6$, Cl$_2$O$_7$. Higher oxides are more stable. Reactivity towards Halogens: Form interhalogen compounds (XX', XX$_3$', XX$_5$', XX$_7'$). Reactivity towards Metals: Form metal halides. Ionic character: MF > MCl > MBr > MI. Oxoacids of Halogens: HXO, HXO$_2$, HXO$_3$, HXO$_4$. Acid strength increases with oxidation state of halogen. Chlorine (Cl$_2$): Preparation: Oxidation of HCl (MnO$_2$, KMnO$_4$); Deacon's process (HCl + O$_2$/CuCl$_2$); electrolysis of brine. Properties: Greenish-yellow gas, pungent odor, poisonous; dissolves in water (chlorine water). Reactions: With metals, nonmetals; high affinity for H; with NH$_3$ (excess NH$_3 \rightarrow$ N$_2$, excess Cl$_2 \rightarrow$ NCl$_3$); with alkalis (cold/dilute $\rightarrow$ hypochlorite, hot/conc. $\rightarrow$ chlorate); bleaching agent (requires moisture). Hydrogen Chloride (HCl): Preparation: NaCl + H$_2$SO$_4$. Properties: Colorless, pungent gas; highly soluble in water (hydrochloric acid, strong acid). Reactions: With NH$_3 \rightarrow$ NH$_4$Cl; with noble metals in aqua regia (conc. HNO$_3$ + conc. HCl). Interhalogen Compounds: Compounds formed by combination of different halogens (e.g., ClF, BrF$_3$, IF$_7$). Halogen with larger size and more electropositive is central atom (X). 7.3 Group 18 (Noble Gases) Elements: He, Ne, Ar, Kr, Xe, Rn. Electronic Configuration: ns$^2$np$^6$ (stable, completely filled valence shell). Occurrence: All (except Rn) in atmosphere. He, Ne in radioactive minerals. Atomic Properties: Atomic Radii: Increase down the group. Ionization Enthalpy: Very high, decreases down the group. Electron Gain Enthalpy: Large positive values (no tendency to accept electrons). Physical Properties: Monoatomic gases. Sparingly soluble in water. Very low melting/boiling points (due to weak van der Waals forces). He lowest BP (4.2 K). Chemical Properties: Chemically inert (stable electronic configuration). Heavier noble gases (Kr, Xe, Rn) can form compounds due to lower ionization energy. Xenon Compounds: Direct reaction with F$_2$ $\rightarrow$ XeF$_2$, XeF$_4$, XeF$_6$. Also form oxides (XeO$_3$) and oxyfluorides (XeOF$_2$, XeO$_2$F$_2$, XeOF$_4$). 8. Transition and Inner Transition Elements 8.1 Transition Elements (d-block) Definition: Elements with incomplete d-subshell in atomic or ionic state. Groups 3-12. Electronic Configuration: (n-1)d$^{1-10}$ ns$^{1-2}$. Exceptions: Cr ([Ar]3d$^5$4s$^1$), Cu ([Ar]3d$^{10}$4s$^1$). Zn, Cd, Hg: Not considered transition elements as they have completely filled d-orbitals in ground and common oxidation states. Oxidation States: Exhibit variable oxidation states (due to small energy difference between (n-1)d and ns orbitals). Maximum oxidation states near the middle of the series (e.g., Mn +2 to +7). Lower oxidation states: Ionic compounds. Higher oxidation states: Covalent compounds. Physical Properties: All are metals: Hard, lustrous, malleable, ductile. High melting/boiling points (strong metallic bonding involving d-electrons). Good conductors of heat and electricity. Form alloys. Atomic and Ionic Radii: Decrease gradually across a period (increasing effective nuclear charge, poor shielding by d-electrons). Ionic radii decrease with increasing oxidation state. Ionization Enthalpy: Intermediate between s-block and p-block. Increases across a period. Third series has higher IE than first two (due to poor shielding by 4f electrons - Lanthanoid contraction). Magnetic Properties: Paramagnetic: Due to unpaired electrons. Magnetic moment ($\mu$) = $\sqrt{n(n+2)}$ BM (spin-only formula). Diamagnetic: All electrons paired. Ferromagnetic: Fe, Co, Ni. Strong attraction, can be permanently magnetized. Colour: Compounds are often colored. Caused by d-d transitions (absorption of visible light). Ions with no unpaired d-electrons (d$^0$ or d$^{10}$) are colorless. Color depends on ligands and geometry (Crystal Field Theory). Catalytic Properties: Many transition metals and their compounds are good catalysts. Ability to participate in redox steps by changing oxidation states. Provide surface for reactions (heterogeneous catalysis). Formation of Interstitial Compounds: Small atoms (H, C, N) trapped in interstitial spaces of crystal lattice. Hard, high melting points, similar chemical properties to parent metal. Formation of Alloys: Metals with similar radii form alloys (e.g., steel, brass). Compounds of Mn and Cr: KMnO$_4$ (Potassium Permanganate): Strong oxidizing agent in acidic, neutral, or alkaline media. K$_2$Cr$_2$O$_7$ (Potassium Dichromate): Strong oxidizing agent in acidic medium. 8.2 Inner Transition Elements (f-block) Definition: Elements where f-orbitals are filled. Placed separately at bottom of periodic table. Lanthanoids (4f series): Ce to Lu. Electronic configuration: [Xe]4f$^{0-14}$5d$^{0-1}$6s$^2$. Actinoids (5f series): Th to Lr. Electronic configuration: [Rn]5f$^{0-14}$6d$^{0-2}$7s$^2$. Properties of Lanthanoids: Soft, silvery-white metals. Moderate densities. Highly reactive (react with water). Common oxidation state: +3. Some show +2 (Eu$^{2+}$, Yb$^{2+}$) or +4 (Ce$^{4+}$). Many +3 ions are colored (due to f-f transitions). Lanthanoid Contraction: Gradual decrease in atomic and ionic radii across the series (due to poor shielding by 4f electrons, increasing nuclear charge). Strongly paramagnetic. Properties of Actinoids: Highly reactive, radioactive. Common oxidation state: +3. Exhibit a wider range of oxidation states (+2 to +7) than lanthanoids. Actinoid Contraction: Decrease in atomic and ionic radii across the series (poor shielding by 5f electrons). Less reactive than lanthanoids. Similarities between Lanthanoids and Actinoids: Both show +3 oxidation state. f-orbitals are filled gradually. Ionic radii decrease with atomic number (contractions). Low electronegativity, highly reactive. Differences between Lanthanoids and Actinoids: Lanthanoids max OS +4, Actinoids up to +7. Actinoids form complexes more easily. Most lanthanoids non-radioactive (except Pm), all actinoids radioactive. Actinoids form oxocations (e.g., UO$_2^{2+}$). Most lanthanoids colorless, actinoids colored. Postactinoid Elements: Elements with atomic number > 103. Man-made, radioactive, very short half-lives. 9. Coordination Compounds Coordination Compound: Central metal atom/ion surrounded by ligands linked by coordinate bonds. Ligand: Ion or molecule that donates electron pair to central metal ion (Lewis base). Central Metal Ion: Electron pair acceptor (Lewis acid). Types of Ligands: Monodentate: Single donor atom (e.g., Cl$^-$, NH$_3$). Polydentate: Two or more donor atoms (e.g., ethylenediamine (en), oxalate (C$_2$O$_4^{2-}$), EDTA (hexadentate)). Ambidentate: Two donor atoms, but uses only one at a time (e.g., NO$_2^-$, SCN$^-$). Coordination Sphere: Central metal ion + ligands, enclosed in square brackets. Complex Ion: Coordination sphere carrying a net charge. Counter Ions: Ionizable groups outside the coordination sphere. Oxidation State (O.S.) of Metal: Charge carried by the metal ion. Coordination Number (C.N.): Number of ligand donor atoms directly attached to the central metal ion. Double Salt: Dissociates completely into simple ions in water (e.g., Mohr's salt). Coordination Complex: Dissociates to give at least one complex ion in water (e.g., K$_4$[Fe(CN)$_6$]). Werner's Theory: Metal has two types of valencies: primary (ionizable) and secondary (non-ionizable). Primary valencies (O.S.) satisfied by anions. Secondary valencies (C.N.) satisfied by ligands (anions or neutral molecules), fixed spatial arrangement. Classification of Complexes: Based on Ligands: Homoleptic: Only one type of ligand (e.g., [Co(NH$_3$)$_6$]$^{3+}$). Heteroleptic: More than one type of ligand (e.g., [Co(NH$_3$)$_4$Cl$_2$]$^+$). Based on Charge: Cationic: Positively charged complex ion. Anionic: Negatively charged complex ion. Neutral: No net charge on complex (e.g., [Ni(CO)$_4$]). IUPAC Nomenclature: Ligands named first, then metal. Anionic ligands end in '-o' (e.g., chloro, cyano). Neutral ligands often have special names (e.g., ammine, aqua). Prefixes (di-, tri-, etc.) for multiple identical ligands; (bis-, tris-, etc.) if ligand name has numerical prefix. Metal oxidation state in Roman numerals in parentheses. Metal name ends in '-ate' for anionic complexes. Effective Atomic Number (EAN) Rule: Metal ion accepts electron pairs until it attains the electronic configuration of the next noble gas. EAN = Z - X + Y. Isomerism: Different compounds with the same molecular formula. Stereoisomers: Same connectivity, different spatial arrangement. Geometric (cis/trans): Ligands in adjacent (cis) or opposite (trans) positions. Common in square planar (MA$_2$B$_2$, MA$_2$BC) and octahedral (MA$_4$B$_2$, M(AA)$_2$B$_2$) complexes. Not in tetrahedral. Optical (enantiomers): Non-superimposable mirror images. Chiral complexes. Rotate plane-polarized light. Structural (Constitutional) Isomers: Different linkages among atoms. Linkage: Ambidentate ligand coordinates via different donor atoms (e.g., -NO$_2$ vs -ONO). Ionization: Exchange of ligands between coordination and ionization spheres (e.g., [Co(NH$_3$)$_5$SO$_4$]Br vs [Co(NH$_3$)$_5$Br]SO$_4$). Coordination: Interchange of ligands between cationic and anionic complex ions. Solvate (Hydrate): Water as ligand or solvent of crystallization. Stability of Coordination Compounds: Measured by stability constant (K). Greater K means greater stability. Factors: Charge-to-size ratio of metal ion, nature (basicity) of ligand. Theories of Bonding: Valence Bond Theory (VBT): Based on hybridization. Metal provides vacant orbitals, hybridization occurs, ligands donate electrons. Predicts geometry and magnetic properties. Limitations: Doesn't explain high/low spin, color, or accurate structures. Crystal Field Theory (CFT): Ligands are point charges, electrostatic interaction. Degeneracy of d-orbitals is destroyed (split) in ligand field. Octahedral Splitting: 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$_{z^2}$) sets. Energy difference is $\Delta_o$. Ligand Strength: Strong field ligands (C, N, P donors) cause large $\Delta_o$, favoring low spin (pairing). Weak field ligands (halogens, O, S donors) cause small $\Delta_o$, favoring high spin. Tetrahedral Splitting: d-orbitals split into e$_g$ (lower energy) and t$_{2g}$ (higher energy). $\Delta_t$ is smaller than $\Delta_o$, so pairing is not favored (always high spin). Color: Explained by d-d transitions (electrons jump between split d-orbitals by absorbing visible light). Applications: Biology (chlorophyll, hemoglobin), medicine (cisplatin, EDTA), water hardness estimation, electroplating. 10. Halogen Derivatives Definition: Hydrocarbons where H atoms are replaced by halogen atoms. Classification: Based on Hydrocarbon Skeleton: Haloalkanes, Haloalkenes, Haloalkynes, Haloarenes. Based on Number of Halogen Atoms: Mono-, Di-, Tri-, Polyhalogen compounds. Monohalogen Compounds: Alkyl Halides (Haloalkanes): Halogen attached to sp$^3$ carbon of saturated skeleton (1$^0$, 2$^0$, 3$^0$). Allylic Halides: Halogen attached to sp$^3$ carbon next to a C=C double bond. Benzylic Halides: Halogen attached to sp$^3$ carbon next to an aromatic ring. Vinylic Halides: Halogen attached to sp$^2$ carbon of a C=C double bond. Haloalkynes: Halogen attached to sp carbon of a C$\equiv$C triple bond. Aryl Halides (Haloarenes): Halogen directly attached to sp$^2$ carbon of an aromatic ring. Nomenclature: Common names (alkyl halide), IUPAC (haloalkane/haloarene). Methods of Preparation of Alkyl Halides: From Alcohols: R-OH + HX $\rightarrow$ R-X + H$_2$O (reactivity: 3$^0$ > 2$^0$ > 1$^0$; HI > HBr > HCl). R-OH + PX$_3 \rightarrow$ R-X. R-OH + SOCl$_2 \rightarrow$ R-Cl (preferred for R-Cl). From Hydrocarbons: Addition of HX to alkenes (Markownikoff's rule; anti-Markownikoff with peroxides for HBr). Addition of X$_2$ to alkenes (vicinal dihalides). Allylic halogenation (high temperature). Halogen Exchange: Finkelstein Reaction: R-Cl/Br + NaI $\rightarrow$ R-I + NaCl/NaBr (in acetone). Swartz Reaction: R-Cl/Br + AgF/Hg$_2$F$_2$/AsF$_3$/SbF$_3 \rightarrow$ R-F + AgCl/Br. Methods of Preparation of Aryl Halides: Electrophilic Substitution: Benzene/derivatives + X$_2$ (FeX$_3$ catalyst). Sandmeyer's Reaction: Diazonium salt $\rightarrow$ Aryl halide (CuX/HX). Physical Properties: Nature of Intermolecular Forces: C-X bond is polar ($\delta^+$C-$\delta^-$X). Moderately polar compounds. Boiling Point: Higher than corresponding alkanes (polarity, higher molecular mass). Increases with increasing atomic mass of halogen (RI > RBr > RCl > RF). Increases with increasing carbon number. Decreases with branching. Solubility: Insoluble in water (cannot form H-bonds). Soluble in non-polar organic solvents. Melting Point of Dihalobenzenes: Para isomer has higher MP due to symmetrical packing. Optical Isomerism: Chiral Carbon: Carbon atom bonded to four different groups. Chiral Molecule: Non-superimposable on its mirror image (exhibits handedness/chirality). Plane Polarized Light: Light oscillating in a single plane. Optical Activity: Property of a substance to rotate the plane of plane polarized light. Dextrorotatory (d or +): Rotates right (clockwise). Laevorotatory (l or -): Rotates left (anticlockwise). Enantiomers: Non-superimposable mirror images with equal and opposite optical rotation. Identical physical properties (except optical rotation), identical chemical properties (except with chiral reagents). Racemic Mixture: Equimolar mixture of enantiomers; optically inactive. Representation: Fischer projection formula, Wedge formula. Chemical Properties of Alkyl Halides: Laboratory Test: Hydrolysis to X$^-$, then Ag$^+$ + X$^- \rightarrow$ AgX precipitate. Nucleophilic Substitution Reactions (S$_N$): Halogen replaced by nucleophile. C-X bond breaks (heterolysis), new C-Nu bond forms. Reactivity: R-I > R-Br > R-Cl; 3$^0$ > 2$^0$ > 1$^0$. S$_N$2 Mechanism (Substitution Nucleophilic Bimolecular): Single step, concerted mechanism. Rate depends on [Alkyl Halide] and [Nucleophile]. Backside attack of nucleophile. Pentacoordinate transition state. Inversion of configuration (Walden inversion). Favored by 1$^0$ halides, strong nucleophiles, aprotic solvents. S$_N$1 Mechanism (Substitution Nucleophilic Unimolecular): Two-step mechanism. Step 1 (slow): Heterolysis of C-X bond to form planar carbocation intermediate. Step 2 (fast): Nucleophile attacks carbocation. Rate depends on [Alkyl Halide] only. Racemization (product is nearly racemic). Favored by 3$^0$ halides, weak nucleophiles, polar protic solvents. Elimination Reaction (Dehydrohalogenation): Alkyl halide + alc. KOH $\rightarrow$ alkene. $\beta$-elimination (H from $\beta$-carbon, X from $\alpha$-carbon). Saytzeff's Rule: Preferred product is the more substituted alkene. Reaction with Active Metals: Grignard Reagent: R-X + Mg (dry ether) $\rightarrow$ R-Mg-X. (Organometallic compound, nucleophilic carbon). Wurtz Reaction: 2R-X + 2Na (dry ether) $\rightarrow$ R-R + 2NaX. Reactions of Haloarenes: Reaction with Metals: Wurtz-Fittig Reaction: Aryl-X + R-X + 2Na (dry ether) $\rightarrow$ Aryl-R + 2NaX. Fittig Reaction: 2Aryl-X + 2Na (dry ether) $\rightarrow$ Aryl-Aryl + 2NaX. Nucleophilic Substitution (S$_N$): Low reactivity due to: Resonance effect (C-X bond has partial double bond character, stronger and shorter). sp$^2$ hybridization of carbon (stronger bond). Unstable phenyl carbocation. Steric hindrance from aromatic ring. Activated by electron-withdrawing groups at ortho/para positions. Electrophilic Substitution (S$_E$): Halogen is ortho/para directing but ring deactivating (due to -I effect). Halogenation, Nitration, Sulfonation, Friedel-Crafts alkylation/acylation. Uses and Environmental Effects of Polyhalogen Compounds: Dichloromethane (CH$_2$Cl$_2$): Solvent, propellant. Toxic. Chloroform (CHCl$_3$): Solvent, refrigerant R-22 precursor. CNS depressant, forms poisonous phosgene (COCl$_2$) in air/light. Carbon Tetrachloride (CCl$_4$): Solvent, source of Cl. Toxic to liver, carcinogenic, greenhouse gas. Iodoform (CHI$_3$): Antiseptic. Toxic. Freons (CFCs): Refrigerants, propellants. Cause ozone depletion. DDT (Dichlorodiphenyltrichloroethane): Insecticide. Persistent organic pollutant, bioaccumulates, toxic to fish, banned in many countries. 11. Alcohols, Phenols and Ethers 11.1 Alcohols Definition: Organic compounds with -OH group attached to saturated carbon. Classification: Mono/Di/Tri/Polyhydric: Based on number of -OH groups. Based on Hybridization of C-OH bond: sp$^3$C-OH: -OH attached to sp$^3$ carbon. 1$^0$, 2$^0$, 3$^0$ alcohols. Allylic alcohols (sp$^3$ C-OH next to C=C). Benzylic alcohols (sp$^3$ C-OH next to aromatic ring). sp$^2$C-OH: -OH attached to sp$^2$ carbon (vinylic alcohols). Nomenclature: Common: Alkyl alcohol (e.g., methyl alcohol). Carbinol: Derivatives of methyl alcohol (carbinol). IUPAC: Alkanols (replace 'e' of alkane with 'ol'). Preparation: From Alkyl Halides: Hydrolysis with aq. NaOH/KOH or moist Ag$_2$O. From Alkenes: Acid-catalyzed hydration (Markownikoff's addition). Hydroboration-oxidation (anti-Markownikoff's addition). From Carbonyl Compounds: Reduction of aldehydes (1$^0$ alcohols) and ketones (2$^0$ alcohols) with H$_2$/Ni or LiAlH$_4$. Reduction of carboxylic acids (1$^0$ alcohols) with LiAlH$_4$ (expensive) or catalytic hydrogenation of esters. From Grignard Reagents: R-Mg-X + aldehyde/ketone $\rightarrow$ adduct $\rightarrow$ alcohol on hydrolysis. HCHO $\rightarrow$ 1$^0$ alcohol. Aldehyde $\rightarrow$ 2$^0$ alcohol. Ketone $\rightarrow$ 3$^0$ alcohol. Epoxide $\rightarrow$ 1$^0$ alcohol. Physical Properties: Intermolecular Forces: Strong H-bonding due to polar -OH group. Physical State: Lower alcohols are toxic liquids, characteristic odor. Boiling Points: Increase with molecular mass. Higher than alkanes/ethers due to H-bonding. Decreases with branching. Solubility: Lower alcohols soluble in water (H-bonding with water). Decreases with increasing hydrocarbon part. Chemical Properties: Laboratory Tests: Litmus Test: Aqueous alcohols are neutral. Sodium Test: 2R-OH + 2Na $\rightarrow$ 2R-ONa + H$_2 \uparrow$ (detects -OH group). Lucas Test (for 1$^0$, 2$^0$, 3$^0$ alcohols): Conc. HCl + ZnCl$_2$. Turbidity indicates alkyl chloride formation. 3$^0$ (instant) > 2$^0$ (slow) > 1$^0$ (on heating). Reactions due to O-H bond cleavage (Acidic Character): Alcohols are very weak acids (weaker than water). Alkyl groups destabilize alkoxide ion (+I effect). Esterification: R-OH + R'-COOH $\rightarrow$ R'-COOR + H$_2$O (acid-catalyzed). Also with acid halides/anhydrides. Reactions due to C-O bond cleavage: With Hydrogen Halides: R-OH + HX $\rightarrow$ R-X + H$_2$O (reactivity: HI > HBr > HCl; 3$^0$ > 2$^0$ > 1$^0$). With Phosphorous Halides: R-OH + PX$_3 \rightarrow$ R-X. R-OH + PCl$_5 \rightarrow$ R-Cl. Dehydration to Alkenes: With conc. H$_2$SO$_4$/H$_3$PO$_4$/Al$_2$O$_3$. Follows Saytzeff's rule (more substituted alkene). Oxidation: 1$^0$ alcohols: R-CHO (PCC) or R-COOH (strong oxidants like KMnO$_4$). 2$^0$ alcohols: Ketones (CrO$_3$). 3$^0$ alcohols: Resistant; strong oxidants/high temp. cause C-C bond cleavage to give mixture of carboxylic acids. Dehydrogenation (with Cu/573K): 1$^0 \rightarrow$ aldehyde, 2$^0 \rightarrow$ ketone, 3$^0 \rightarrow$ alkene. 11.2 Phenols Definition: -OH group directly attached to aromatic ring. Nomenclature: Phenol (common and IUPAC). Substituted phenols use o-/m-/p- or locants. Preparation: From Chlorobenzene (Dow Process): Chlorobenzene + NaOH (high T/P) $\rightarrow$ Sodium phenoxide $\rightarrow$ Phenol (acidification). From Cumene: Cumene + O$_2$ $\rightarrow$ Cumene hydroperoxide $\rightarrow$ Phenol + Acetone (acid). From Benzene Sulfonic Acid: Benzene sulfonic acid $\rightarrow$ Sodium benzene sulfonate $\rightarrow$ Sodium phenoxide (fusion with NaOH) $\rightarrow$ Phenol (acidification). From Aniline: Aniline $\rightarrow$ Benzene diazonium chloride (HNO$_2$/0$^0$C) $\rightarrow$ Phenol (hydrolysis). Physical Properties: Intermolecular Forces: Strong H-bonding. Physical State: Colorless, toxic, low melting solid, carbolic odor. Boiling Points: Higher than corresponding alcohols. Solubility: Appreciable solubility in water (H-bonding). Chemical Properties: Laboratory Tests: Litmus Test: Aqueous phenols turn blue litmus red (weakly acidic). FeCl$_3$ Test: Phenols + neutral FeCl$_3 \rightarrow$ deep (purple/violet/green) coloration of ferric phenoxide. Acidic Character: Weaker acids than carboxylic acids, stronger than alcohols. React with NaOH but not NaHCO$_3$. Phenoxide ion is resonance stabilized. Esterification: Ar-OH + R'-COOH $\rightarrow$ Ar-O-COR' + H$_2$O. Also with acid halides/anhydrides. Electrophilic Substitution (o-/p- directing, ring activating): Halogenation: With Br$_2$/H$_2$O $\rightarrow$ 2,4,6-tribromophenol. With Br$_2$ in nonpolar solvent $\rightarrow$ o-/p-bromophenol. Nitration: Dilute HNO$_3 \rightarrow$ o-/p-nitrophenol. Conc. HNO$_3 \rightarrow$ 2,4,6-trinitrophenol (picric acid). Sulfonation: Conc. H$_2$SO$_4 \rightarrow$ o- (room T) or p- (373K) phenolsulfonic acid. Reimer-Tiemann Reaction: Phenol + CHCl$_3$ + aq. NaOH $\rightarrow$ Salicylaldehyde (followed by acid hydrolysis). Kolbe Reaction: Sodium phenoxide + CO$_2$ (high T/P) $\rightarrow$ Sodium salicylate $\rightarrow$ Salicylic acid (acid hydrolysis). Oxidation: Phenol + CrO$_3$/Na$_2$Cr$_2$O$_7$/H$_2$SO$_4 \rightarrow$ p-benzoquinone. Oxidizes slowly in air. Catalytic Hydrogenation: Phenol + H$_2$ (Ni/433K) $\rightarrow$ Cyclohexanol. Reduction: Phenol + Zn dust $\rightarrow$ Benzene. 11.3 Ethers Definition: Organic oxides (R-O-R', R-O-Ar, Ar-O-Ar'). Classification: Symmetrical (Simple): R and R' are identical. Unsymmetrical (Mixed): R and R' are different. Nomenclature: Common: Alkyl alkyl ether (e.g., diethyl ether). IUPAC: Alkoxyalkane (smaller alkyl group as alkoxy substituent). Preparation: Dehydration of Alcohols: 2R-OH $\rightarrow$ R-O-R + H$_2$O (conc. H$_2$SO$_4$/413K). Favored for 1$^0$ alcohols. Symmetrical ethers. Williamson Synthesis: R-X + NaOR' $\rightarrow$ R-O-R' + NaX. (S$_N$2 mechanism). Requires 1$^0$ alkyl halide. Physical Properties: Physical State: Lower ethers are gases, others are colorless liquids with pleasant odor. Highly volatile and inflammable. Boiling Points: Increase with molecular mass. Similar to alkanes of comparable mass (no H-bonding). Polarity: Small net dipole moment (C-O-C bond angle $\neq$ 180$^0$). Solubility: Miscible with water (can form H-bonds with water). Chemical Properties: Laboratory Test: Dissolve in cold conc. H$_2$SO$_4$ (form oxonium salts). Reactions involving Alkyl Group: Peroxide Formation: Ethers + atmospheric O$_2 \rightarrow$ peroxides (hazardous). Reactions involving C-O bond: Hydrolysis: With hot dilute H$_2$SO$_4 \rightarrow$ alcohols/phenols. With PCl$_5$: R-O-R' + PCl$_5 \rightarrow$ R-Cl + R'-Cl + POCl$_3$. With Hot Concentrated HX: R-O-R + HX $\rightarrow$ R-X + R-OH. Order of reactivity: HI > HBr > HCl. For t-butyl ether, SN1 mechanism gives t-alkyl halide. Aryl alkyl ethers give phenol + alkyl halide. Electrophilic Substitution (in aromatic ethers): Alkoxy group is o-/p- directing and ring activating (+R effect). Halogenation, Friedel-Crafts alkylation/acylation, Nitration. 12. Aldehydes, Ketones and Carboxylic Acids 12.1 Aldehydes and Ketones Definition: Organic compounds containing a carbonyl group (>C=O). Aldehydes: Carbonyl carbon bonded to at least one H atom (-CHO, formyl group). R-CHO. Ketones: Carbonyl carbon bonded to two alkyl/aryl groups (>C=O, ketonic carbonyl group). R-CO-R'. Classification: Aldehydes: Aliphatic: -CHO group attached to sp$^3$ carbon (or H for formaldehyde). Aromatic: -CHO group attached to aromatic ring. Ketones: Aliphatic: >C=O attached to two alkyl groups. Simple (Symmetrical): Alkyl groups are identical. Mixed (Unsymmetrical): Alkyl groups are different. Aromatic: >C=O attached to two aryl groups or one aryl and one alkyl group. Nomenclature: Aldehydes: Trivial: Derived from carboxylic acids (replace '-ic acid' with '-aldehyde'). Positions indicated by $\alpha, \beta, \gamma$. IUPAC: Alkanal (replace 'e' of alkane with 'al'). -CHO carbon is C-1. For cyclic: -carbaldehyde. Ketones: Trivial: Alkyl alkyl ketone (alphabetical order). IUPAC: Alkanone (replace 'e' of alkane with 'one'). Numbered from end closer to carbonyl. For cyclic: -one. Preparation: General Methods (for both): Oxidation of Alcohols: 1$^0$ alcohol $\rightarrow$ aldehyde; 2$^0$ alcohol $\rightarrow$ ketone. Dehydrogenation of Alcohols: Vapors over hot Cu (industrial). Ozonolysis of Alkenes: Alkene + O$_3 \rightarrow$ ozonide $\rightarrow$ aldehyde/ketone (Zn/H$_2$O). Hydration of Alkynes: Alkyne + H$_2$O (H$_2$SO$_4$/HgSO$_4$) $\rightarrow$ aldehyde/ketone. Other Methods: From Acyl Chlorides: Aldehydes (Rosenmund Reduction): R-COCl + H$_2$ (Pd-BaSO$_4$) $\rightarrow$ R-CHO. Ketones: R-COCl + R$_2$Cd $\rightarrow$ R-CO-R (dialkyl cadmium from Grignard). Aromatic ketones by Friedel-Crafts acylation. From Nitriles: Aldehydes (Stephen Reaction): R-C$\equiv$N + SnCl$_2$/HCl $\rightarrow$ imine hydrochloride $\rightarrow$ R-CHO (H$_3$O$^+$). Also by DIBAl-H. Ketones: R-C$\equiv$N + R'-Mg-X $\rightarrow$ imine $\rightarrow$ R-CO-R' (H$_3$O$^+$). From Aromatic Hydrocarbons: Aldehydes: Etard reaction (methyl arene + CrO$_2$Cl$_2 \rightarrow$ chromium complex $\rightarrow$ benzaldehyde); Oxidation with CrO$_3$/acetic anhydride; Side-chain chlorination of toluene; Gatterman-Koch reaction (benzene + CO/HCl/AlCl$_3$). Ketones: Friedel-Crafts acylation. Aldehydes from Esters: R-COOR' + DIBAl-H $\rightarrow$ R-CHO. Physical Properties: Intermolecular Forces: Dipole-dipole attraction due to polar C=O. Physical State: Formaldehyde (gas), acetaldehyde (volatile liquid), higher aldehydes (pleasant odor), ketones (pleasant odor). Boiling Points: Increase with molecular mass. Higher than alkanes/ethers, lower than alcohols/carboxylic acids (no H-bonding). Solubility: Lower aldehydes/ketones soluble in water (H-bonding with water). Decreases with increasing hydrocarbon part. Polarity and Reactivity of Carbonyl Group: C=O is polar ($\delta^+$C-$\delta^-$O). Carbonyl carbon is electrophilic, susceptible to nucleophilic attack. Reactivity: Aldehydes > Ketones. Electronic Effects: Ketones have two +I alkyl groups decreasing electrophilicity. Steric Effects: Ketones have two bulky alkyl groups causing steric hindrance. Chemical Properties: Laboratory Tests: For Aldehydes: Schiff Test: Pink/red color. Tollens' Test (Silver Mirror Test): Aldehyde $\rightarrow$ carboxylate, Ag$^+ \rightarrow$ Ag. Fehling Test: Aldehyde $\rightarrow$ carboxylate, Cu$^{2+} \rightarrow$ Cu$_2$O (red ppt). For Ketones: Sodium Nitroprusside Test: Red color with NaOH. Reactions with Nucleophiles (Nucleophilic Addition): Addition of HCN: Aldehyde/ketone + HCN $\rightarrow$ cyanohydrin. Addition of NaHSO$_3$: Aldehyde/ketone + NaHSO$_3 \rightarrow$ bisulfite adduct (crystalline). Addition of Alcohols: Aldehyde + 1 alcohol $\rightarrow$ hemiacetal $\rightarrow$ acetal (stable). Ketone + 1 alcohol $\rightarrow$ hemiketal $\rightarrow$ ketal. Ketone + diol $\rightarrow$ cyclic ketal. Addition of Grignard Reagents: $\rightarrow$ Alcohols. Nucleophilic Addition-Elimination with Ammonia Derivatives (NH$_2$-Z): $\rightarrow$ Imines (C=N bonds). Forms crystalline derivatives useful for characterization. Examples: oximes (NH$_2$OH), hydrazones (NH$_2$NH$_2$), phenylhydrazones (NH$_2$NHC$_6$H$_5$), semicarbazones (NH$_2$NHCONH$_2$). Haloform Reaction (Iodoform Test): Given by acetaldehyde, methyl ketones (CH$_3$-CO-R), and alcohols with CH$_3$-CH(OH)- group. CH$_3$-CO-R + I$_2$/NaOH $\rightarrow$ CHI$_3$ (yellow ppt) + R-COONa. Aldol Condensation: Aldehydes/ketones with at least one $\alpha$-H atom, in presence of dilute alkali. Aldehyde $\rightarrow \beta$-hydroxy aldehyde (aldol) $\rightarrow$ $\alpha,\beta$-unsaturated aldehyde (on warming). Ketone $\rightarrow \beta$-hydroxy ketone (ketol) $\rightarrow$ $\alpha,\beta$-unsaturated ketone. Cross Aldol Condensation: Between two different aldehydes/ketones. Cannizzaro Reaction: Aldehydes with no $\alpha$-H atom, on heating with conc. alkali. Self-oxidation and reduction (disproportionation). One molecule $\rightarrow$ alcohol, another $\rightarrow$ carboxylic acid salt. Oxidation and Reduction Reactions: Oxidation: Aldehydes $\rightarrow$ carboxylic acids (dilute HNO$_3$, KMnO$_4$, K$_2$Cr$_2$O$_7$). Ketones $\rightarrow$ mixture of carboxylic acids (strong oxidants, C-C bond cleavage). Reduction of Carbonyl Group to Methylene (-CH$_2$-): Clemmensen Reduction: Aldehyde/ketone + Zn-Hg/conc. HCl $\rightarrow$ alkane. Wolf-Kishner Reduction: Aldehyde/ketone + Hydrazine $\rightarrow$ hydrazone $\rightarrow$ alkane (KOH/ethylene glycol). Electrophilic Substitution (Aromatic Aldehydes/Ketones): Carbonyl group is electron-withdrawing and meta-directing. Nitration, Sulfonation, Halogenation give meta-products. 12.2 Carboxylic Acids Definition: Organic compounds containing a carboxyl group (-COOH). R-COOH. Classification: Aliphatic: -COOH group bonded to alkyl group (or H for formic acid). Mono/Di/Tri carboxylic acids. Aromatic: -COOH group bonded directly to aromatic ring. Nomenclature: Trivial: Often from Latin names of source (e.g., formic, acetic). Positions indicated by $\alpha, \beta, \gamma$. IUPAC: Alkanoic acid (replace 'e' of alkane with 'oic acid'). -COOH carbon is C-1. For cyclic: -carboxylic acid. Preparation: From Nitriles and Amides: Nitrile $\rightarrow$ amide $\rightarrow$ carboxylic acid (acid hydrolysis). From Acyl Chlorides and Anhydrides: Hydrolysis with water. From Esters: Acid hydrolysis or alkaline hydrolysis (saponification). From Alkyl Benzene: Oxidation of alkyl benzene (alkaline KMnO$_4$/acidic KMnO$_4$/chromic acid). Alkyl chain is oxidized to -COOH. From Alkenes: Oxidation with KMnO$_4$/dil. H$_2$SO$_4$. From Grignard Reagent: R-Mg-X + CO$_2$ (dry ice) $\rightarrow$ R-COOMgX $\rightarrow$ R-COOH (acid hydrolysis). Physical Properties: Physical State: Lower aliphatic acids (liquids, irritating odor), higher homologues (waxy solids, odorless). Boiling Points: Higher than alkanes, ethers, alcohols, aldehydes, ketones of comparable mass. Due to dimer formation via two H-bonds in liquid and gas phase. Solubility: Lower acids soluble in water (H-bonding with water). Decreases with molecular mass. Insoluble acids soluble in less polar organic solvents. Chemical Properties: Acidic Character: -COOH group imparts acidity. Ionize to R-COO$^-$ + H$^+$. Carboxylate ion is resonance stabilized by two equivalent structures. Stronger acids than alcohols and phenols. Acidity increased by electron-withdrawing groups (-I effect, -R effect) and decreased by electron-donating groups (+I effect, +R effect). Laboratory Tests: Litmus Test: Aqueous solution turns blue litmus red. Sodium Bicarbonate Test: Brisk effervescence of CO$_2$ with NaHCO$_3$. (Distinguishes from phenols). Ester Test: Carboxylic acid + alcohol (H$^+$/heat) $\rightarrow$ ester (fruity smell). Formation of Acyl Chloride: R-COOH + PCl$_3$/PCl$_5$/SOCl$_2 \rightarrow$ R-COCl (SOCl$_2$ preferred). Reaction with Ammonia: R-COOH + NH$_3 \rightarrow$ R-COONH$_4$ (ammonium carboxylate) $\rightarrow$ R-CONH$_2$ (amide) (on heating). Formation of Acid Anhydride: 2R-COOH (P$_2$O$_5$/heat) $\rightarrow$ (R-CO)$_2$O. Or R-COONa + R-COCl. Decarboxylation: Sodium salt of carboxylic acid + soda lime (heat) $\rightarrow$ alkane (one C less). Reduction: R-COOH + LiAlH$_4$ (or diborane) $\rightarrow$ R-CH$_2$OH (1$^0$ alcohol). 13. Amines 13.1 Amines Definition: Nitrogen-containing organic compounds with basic character. Derivatives of ammonia. Classification: Primary (1$^0$): One H of NH$_3$ replaced by alkyl/aryl group (R-NH$_2$). Secondary (2$^0$): Two H of NH$_3$ replaced (R$_2$NH). Tertiary (3$^0$): Three H of NH$_3$ replaced (R$_3$N). Simple (Symmetrical): All alkyl/aryl groups are identical. Mixed (Unsymmetrical): Alkyl/aryl groups are different. Aliphatic: Alkyl groups attached to N. Aromatic: Aryl groups attached to N. Nomenclature: Common: Alkyl amine (e.g., methylamine), or derivatives of aniline. IUPAC: Alkanamine (replace 'e' of alkane with 'amine'); N-substituted derivatives for 2$^0$/3$^0$ amines. Preparation: Ammonolysis of Alkyl Halides: R-X + NH$_3$ (excess alc.) $\rightarrow$ R-NH$_2$. (Can lead to 2$^0$, 3$^0$ amines, and quaternary ammonium salts if NH$_3$ is not in excess). Reactivity: R-I > R-Br > R-Cl. Reduction of Nitrocompounds: R-NO$_2$ or Ar-NO$_2$ + Sn/HCl or H$_2$/Ni or LiAlH$_4 \rightarrow$ R-NH$_2$ or Ar-NH$_2$. Reduction of Alkyl Cyanides (Alkanenitriles): R-C$\equiv$N + Na/C$_2$H$_5$OH (Mendius reduction) or LiAlH$_4 \rightarrow$ R-CH$_2$-NH$_2$. (Step-up reaction). Reduction of Amides: R-CONH$_2$ + LiAlH$_4$ or Na/C$_2$H$_5$OH $\rightarrow$ R-CH$_2$-NH$_2$. Gabriel Phthalimide Synthesis: For 1$^0$ amines only. Phthalimide $\rightarrow$ potassium salt $\rightarrow$ N-alkyl phthalimide (with R-X) $\rightarrow$ 1$^0$ amine (alkaline hydrolysis). Aryl amines cannot be prepared. Hofmann Degradation (Hofmann Bromamide Degradation): R-CONH$_2$ + Br$_2$ + 4KOH $\rightarrow$ R-NH$_2$ + 2KBr + K$_2$CO$_3$ + 2H$_2$O. (Step-down reaction, 1 C less). Physical Properties: Intermolecular Forces: N-H bond is polar. 1$^0$ and 2$^0$ amines have H-bonding (strongest in 1$^0$). 3$^0$ amines have dipole-dipole forces. Boiling Points: 1$^0$ > 2$^0$ > 3$^0$ (due to H-bonding). Lower than alcohols/carboxylic acids of comparable mass. Physical State: Lower aliphatic amines (gases, fishy odor), middle (liquids), higher (solids). Arylamines (colorless liquids, get colored on oxidation). Solubility: Lower aliphatic amines soluble in water (H-bonding with water). Decreases with increasing alkyl group size. Aromatic/higher aliphatic amines are insoluble. Basicity of Amines: Due to lone pair of electrons on N atom (Lewis base). Kb value (larger for stronger base), pKb value (smaller for stronger base). Aliphatic Amines: Basic strength (aq. medium): R$_2$NH > R-NH$_2$ > R$_3$N > NH$_3$. (Order influenced by +I effect and solvation of conjugate acid). Arylamines: Weaker bases than NH$_3$ and aliphatic amines. Lone pair on N is delocalized by resonance with aromatic ring, making it less available for protonation. Anilinium ion is less resonance stabilized than aniline. Chemical Properties: Laboratory Tests: Basic Test: Water soluble amines turn red litmus blue. Insoluble amines dissolve in HCl. Diazotization/Orange Dye Test: For aromatic 1$^0$ amines. Ar-NH$_2$ + NaNO$_2$/HCl (0$^0$C) $\rightarrow$ Ar-N$_2^+$Cl$^-$ (diazonium salt). Then + $\beta$-naphthol/NaOH $\rightarrow$ orange dye. Alkylation (Hofmann's Exhaustive Alkylation): R-NH$_2$ + R'X (excess) $\rightarrow$ R$_4$N$^+$X$^-$ (quaternary ammonium halide). Hofmann Elimination: Quaternary ammonium hydroxide (from R$_4$N$^+$X$^-$ + moist Ag$_2$O) $\rightarrow$ alkene (least substituted) + 3$^0$ amine (on strong heating). Acylation: 1$^0$ and 2$^0$ amines + acyl chloride/acetic anhydride $\rightarrow$ amide (N-substituted). (Nucleophilic substitution, carried out in presence of base like pyridine). Carbylamine Reaction (Isocyanide Test): 1$^0$ amines + CHCl$_3$ + 3KOH (heat) $\rightarrow$ R-NC (foul-smelling alkyl isocyanide). (Test for 1$^0$ amines). Reaction with Nitrous Acid (HNO$_2$): Aliphatic 1$^0$ amines: $\rightarrow$ unstable diazonium salt $\rightarrow$ alcohol + N$_2 \uparrow$. Aromatic 1$^0$ amines: $\rightarrow$ stable diazonium salt (Ar-N$_2^+$Cl$^-$) at 0-5$^0$C. Reactions of Arene Diazonium Salts: Displacement of Diazo Group (-N$_2^+$): Replaced by -Cl, -Br, -CN (Sandmeyer/Gatterman), -I (KI), -F (HBF$_4$/heat), -H (H$_3$PO$_2$/H$_2$O or CH$_3$CH$_2$OH), -OH (H$_2$O/heat). Retention of Diazo Group (Coupling Reactions): With phenols/aromatic amines $\rightarrow$ azo compounds (Ar-N=N-Ar') (brightly colored dyes). Electrophilic aromatic substitution. Reaction with Arenesulfonyl Chloride (Hinsberg's Test): 1$^0$ amine + C$_6$H$_5$SO$_2$Cl $\rightarrow$ N-alkylbenzenesulfonamide (soluble in alkali). 2$^0$ amine + C$_6$H$_5$SO$_2$Cl $\rightarrow$ N,N-dialkylbenzenesulfonamide (insoluble in alkali). 3$^0$ amines do not react. Electrophilic Aromatic Substitution (in aromatic amines): Amino group is o-/p- directing and powerful ring activating. Bromination: Aniline + Br$_2$/H$_2$O $\rightarrow$ 2,4,6-tribromoaniline. (To get monobromo, protect -NH$_2$ by acetylation). Nitration: Direct nitration gives mixture of o-/m-/p-nitroanilines (due to protonation of -NH$_2$ to -NH$_3^+$ in acidic medium). (To get p-nitroaniline, protect -NH$_2$ by acetylation). Sulfonation: Aniline + conc. H$_2$SO$_4 \rightarrow$ sulfanilic acid (exists as zwitter ion). 14. Biomolecules 14.1 Carbohydrates Definition: Polyhydroxy aldehydes or ketones, or compounds yielding them on hydrolysis. Also called saccharides. Molecular formula often C$_x$(H$_2$O)$_y$. Classification: Monosaccharides: Do not hydrolyze further (e.g., glucose, fructose, ribose). Oligosaccharides: Yield 2-10 monosaccharide units on hydrolysis. Disaccharides: 2 units (e.g., sucrose, maltose, lactose). Trisaccharides: 3 units (e.g., raffinose). Polysaccharides: Yield large number of monosaccharide units (e.g., starch, glycogen, cellulose). Nomenclature of Monosaccharides: Aldose: Contains an aldehyde group. Ketose: Contains a ketone group. Prefixes (tri-, tetr-, pent-, hex-): Indicate number of carbons. E.g., glucose is aldohexose, fructose is ketohexose. Glucose (C$_6$H$_12$O$_6$): Aldohexose. Preparation: From sucrose or starch (acid-catalyzed hydrolysis). Structure and Properties: Molecular formula C$_6$H$_12$O$_6$. Six carbons in a straight chain (n-hexane on HI reduction). Contains one carbonyl group (forms oxime with NH$_2$OH, cyanohydrin with HCN). Carbonyl group is aldehyde (oxidized to gluconic acid by Br$_2$ water). Contains five -OH groups (forms pentaacetate). Contains one primary alcoholic -CH$_2$OH group (oxidized to saccharic acid by HNO$_3$). Optical Isomerism: Contains four chiral carbons (C-2, C-3, C-4, C-5). D/L Configuration: Relative configuration based on glyceraldehyde. D-glucose has -OH on right at lowest chiral carbon. Ring Structure: Exists as two cyclic hemiacetal structures ( $\alpha$- and $\beta$-anomers) in equilibrium with open chain form. Formed by reaction of -CHO and C-5 -OH. Six-membered ring (pyranose structure). Haworth formula used. Anomeric carbon (C-1) is a new chiral center. Reducing Nature: Hemiacetal group is potential aldehyde, gives positive Tollens' and Fehling tests. Fructose (C$_6$H$_12$O$_6$): Ketohexose. Laevorotatory. Reducing sugar. Exists as fructopyranose (major) and fructofuranose (in combined state). Ring structure is hemiketal. Disaccharides: Two monosaccharide units linked by glycosidic (-O-) linkage (formed by removal of H$_2$O). Sucrose: D-glucose + D-fructose, linked by $\alpha$-1,2-glycosidic linkage. Non-reducing sugar. Hydrolysis gives invert sugar. Maltose: Two D-glucose units, linked by $\alpha$-1,4-glycosidic linkage. Reducing sugar. Lactose: D-galactose + D-glucose, linked by $\beta$-1,4-glycosidic linkage. Reducing sugar. Polysaccharides: Large number of monosaccharide units linked by glycosidic linkages. Starch: Polymer of $\alpha$-D-glucose. Two components: Amylose (15-20%): Unbranched, $\alpha$-1,4-glycosidic linkages. Water soluble. Amylopectin (80-85%): Branched, $\alpha$-1,4- and $\alpha$-1,6-glycosidic linkages. Water insoluble. Cellulose: Straight chain polymer of $\beta$-glucose units linked by $\beta$-1,4-glycosidic bonds. Very strong, difficult to hydrolyze. Not digestible by humans. Glycogen: Animal starch. Similar to amylopectin but more highly branched. 14.2 Proteins Definition: Polyamides, high molecular weight polymers of $\alpha$-amino acids. $\alpha$-Amino Acids: Carboxylic acids with an amino (-NH$_2$) group bonded to the $\alpha$-carbon. $\alpha$-carbon is chiral (except glycine). L-configuration in proteins. Classification (based on R group): Acidic (R has -COOH), Basic (R has -NH$_2$), Neutral (other R groups). High melting, water soluble crystalline solids due to zwitter ion structure (dipolar ion formed by proton transfer from -COOH to -NH$_2$). Peptide Bond: Amide linkage (-CO-NH-) connecting $\alpha$-amino acid residues. Peptides: Formed by linking amino acids. Dipeptide (2 units), tripeptide (3 units), polypeptide (>10 units). Proteins: Polypeptides with >100 amino acid residues. Have N-terminal (free amino group) and C-terminal (free carboxyl group). Types of Proteins: Globular: Spherical shape, soluble in water (e.g., insulin, enzymes). Fibrous: Elongated, rod-like shape, insoluble in water (e.g., keratin, myosin). Structure of Proteins (Four Levels): Primary: Sequence of $\alpha$-amino acid residues linked by peptide bonds. Secondary: 3D arrangement of localized regions of polypeptide chain. Stabilized by H-bonding. $\alpha$-Helix: Right-handed spiral. H-bonds parallel to axis. $\beta$-Pleated Sheet: Polypeptide strands side-by-side, held by H-bonding. Tertiary: Overall 3D shape of entire polypeptide chain. Stabilized by H-bonding, dipole-dipole, electrostatic, London forces, and disulfide bonds. Quaternary: Arrangement of multiple polypeptide subunits in a protein complex (e.g., hemoglobin). Denaturation of Proteins: Disruption of secondary, tertiary, or quaternary structure (by heat, acid, base, agitation) without breaking peptide bonds. Causes loss of biological activity. Enzymes: Biological catalysts (proteins). Highly specific. Work by lock-and-key mechanism (substrate fits active site). Lower activation energy. 14.3 Nucleic Acids Definition: Store and transfer genetic information. Polymers of nucleotides. Types: RNA (Ribonucleic Acid): Cytoplasm. DNA (Deoxyribonucleic Acid): Nuclei. Nucleotides (Monomers): Consist of three components: Monosaccharide (Sugar): RNA: D-ribose. DNA: 2-deoxy-D-ribose. Nitrogen-Containing Base: Pyrimidines (single ring): Cytosine (C), Uracil (U, in RNA), Thymine (T, in DNA). Purines (two rings): Adenine (A), Guanine (G). Phosphate Group. Nucleoside: Sugar + Base (linked via anomeric carbon of sugar and N of base). Polynucleotide Structure: Nucleotides linked by phosphodiester bonds (3'-OH of one nucleotide to 5'-phosphate of another). 5' end (free phosphate), 3' end (free -OH). Primary structure: Sequence of bases. DNA Double Helix (Watson and Crick Model): Two polynucleotide strands wound into right-handed double helix. Strands run in opposite directions (antiparallel). Sugar-phosphate backbone on outside, bases on inside (perpendicular to axis). Stabilized by H-bonding between complementary base pairs: A-T (2 H-bonds), C-G (3 H-bonds). RNA is single-stranded. 15. Introduction to Polymer Chemistry 15.1 Polymers Definition: High molecular mass macromolecules formed from repeating units (monomers). 'Poly' = many, 'mer' = unit. Monomer: Small molecules that link to form polymers. Polymerization: Process of converting monomers into polymers. Classification: Based on Source/Origin: Natural: From plants (cotton, natural rubber) or animals (wool, silk). Synthetic: Man-made (nylon, terylene). Semisynthetic: Derived from natural polymers (cellulose acetate rayon). Based on Structure: Linear (Straight Chain): Monomers joined in a linear arrangement (PVC, HDP). Branched Chain: Monomers with side chains (LDP). Cross-linked (Network): Linear chains connected by cross-links (bakelite, melamine). Based on Mode of Polymerization: Addition (Chain Growth): Monomers add without loss of small molecules. Repeating unit has same elemental composition as monomer. Often by free radical mechanism (initiation, propagation, termination). E.g., polyethylene. Condensation (Step Growth): Polyfunctional monomers react with elimination of small molecules (H$_2$O, HCl, etc.). Repeating unit has different elemental composition than monomers. E.g., terylene. Ring Opening: Cyclic compounds polymerize by ring opening. Monomer elemental composition same as repeating unit. E.g., $\epsilon$-caprolactam $\rightarrow$ Nylon 6. Based on Intermolecular Forces: Elastomers: Weak van der Waals forces, elastic (vulcanized rubber, buna-S). Fibres: Strong intermolecular forces (H-bonding, dipole-dipole), high tensile strength, crystalline (nylon, terylene). Thermoplastic Polymers: Moderately strong forces, soften on heating, harden on cooling (polyethene, polystyrene). Thermosetting Polymers: Rigid, extensive cross-linking on heating, infusible once hardened (bakelite, melamine-formaldehyde). Based on Type of Monomers: Homopolymers: One type of repeating unit (polyethylene, Nylon 6,6). Copolymers: Two or more different types of repeating units (Buna-S, Buna-N). Based on Biodegradability: Biodegradable: Degraded by microbes (e.g., PHBV, Nylon 2-Nylon 6). Non-biodegradable: Not degraded by microbes (most synthetic plastics). Some Important Polymers: Natural Rubber (Polyisoprene): Monomer: Isoprene (2-methyl-1,3-butadiene). Cis-configuration. Elastic. Vulcanization: Heating rubber with sulfur to introduce cross-links, improving properties (tensile strength, elasticity). Polythene: Low Density Polyethylene (LDP): High P, moderate T, O$_2$/peroxide initiator. Branched chain, low density. High Density Polyethylene (HDP): Low P, moderate T, Ziegler-Natta catalyst. Linear, high density. Teflon (Polytetrafluoroethylene): Monomer: Tetrafluoroethylene (CF$_2$=CF$_2$). Tough, chemically inert, heat resistant (due to C-F bond). Polyacrylonitrile (Orlon/Acrilan): Monomer: Acrylonitrile. Resembles wool. Polyamide Polymers (Nylons): Contain -CO-NH- groups. Nylon 6,6: Monomers: Adipic acid + Hexamethylenediamine. Linear condensation polymer. Nylon 6: Monomer: $\epsilon$-caprolactam (ring opening polymerization). Polyesters: Contain ester linkage. Terylene (Dacron): Monomers: Ethylene glycol + Terephthalic acid. Condensation polymerization. Phenol-Formaldehyde Polymers: Bakelite: Phenol + Formaldehyde (acid/base catalyst). Thermosetting, cross-linked. Melamine-Formaldehyde Polymer: Melamine + Formaldehyde. Thermosetting, cross-linked. Buna-S (SBR): Copolymer of Butadiene + Styrene. Elastomer. Neoprene: Addition polymer of Chloroprene (2-chloro-1,3-butadiene). Synthetic rubber. Viscose Rayon: Semisynthetic fibre (regenerated cellulose). Molecular Mass and Degree of Polymerization (DP): Polymer molecular mass is an average. DP is number of monomer units. Mechanical properties depend on DP. Biodegradable Polymers: Degraded by microbes. PHBV: Copolymer of $\beta$-hydroxybutyric acid + $\beta$-hydroxyvaleric acid. Aliphatic polyester. Nylon 2-Nylon 6: Copolymer of Glycine + $\epsilon$-aminocaproic acid. Polyamide. 16. Green Chemistry and Nanochemistry 16.1 Green Chemistry Definition: Use of chemistry for pollution prevention by environmentally conscious design of chemical products and processes that reduce or eliminate hazardous substances. Sustainable Development: Development that meets present needs without compromising future generations' ability to meet their own needs. Principles of Green Chemistry (12 Principles): Prevention: Prevent waste rather than clean it up. Atom Economy: Maximize incorporation of starting material atoms into product. Less Hazardous Chemical Syntheses: Design syntheses to use and generate substances with little or no toxicity. Designing Safer Chemicals: Design chemical products that are fully effective yet have reduced toxicity. Safer Solvents and Auxiliaries: Avoid or minimize use of hazardous solvents/auxiliaries; use safer alternatives (e.g., water, supercritical CO$_2$). Design for Energy Efficiency: Minimize energy requirements; conduct reactions at ambient T and P (e.g., using catalysts). Use of Renewable Feedstocks: Use renewable raw materials instead of depleting non-renewable ones. Reduce Derivatives: Minimize unnecessary derivatization (e.g., protecting groups) to reduce waste. Catalysis: Use catalytic reagents (as selective as possible) over stoichiometric reagents. Design for Degradation: Design products that degrade into innocuous substances after use. Real-time Analysis for Pollution Prevention: Develop analytical methods for in-process monitoring to prevent hazardous substance formation. Inherently Safer Chemistry for Accident Prevention: Choose substances and forms that minimize potential for accidents. Role of Green Chemistry: Improves quality of life, reduces pollution, saves resources, helps control greenhouse effect. 16.2 Nanochemistry Nanoscience: Study of phenomena and manipulation of materials at atomic, molecular, and macromolecular scales (1-100 nm), where properties differ significantly from bulk. Nanotechnology: Design, characterization, production, and application of structures, devices, and systems by controlling shape and size at nanometer scale. Nanometer Scale: 1-100 nm (1 nm = 10$^{-9}$ m). Nanomaterial: Material with structural components having at least one dimension in the nanometer scale. Zero-Dimensional: All three dimensions in nanoscale (nanoparticles, quantum dots). One-Dimensional: Two dimensions in nanoscale (nanowires, nanorods). Two-Dimensional: One dimension in nanoscale (thin films). Nanochemistry: Combination of chemistry and nanoscience. Deals with designing and synthesizing nanoscale materials with control over size, shape, structure, and composition. Characteristic Features of Nanoparticles: Properties differ from bulk materials due to quantum effects and high surface area. Color: Changes with size (e.g., gold nanoparticles are red). Surface Area: High surface-to-volume ratio, leading to increased reactivity. Catalytic Activity: Enhanced due to large surface area (e.g., Pd, Pt, Au nanoparticles). Thermal Properties (Melting Point): Decreases with decreasing size. Mechanical Properties: Increased hardness/strength. Electrical Conductivity: Can change (e.g., carbon nanotubes can be conductors or semiconductors). Synthesis of Nanomaterials: Bottom-Up Approach: Molecular components self-assemble into complex structures (atom by atom, molecule by molecule). E.g., colloidal dispersion. Top-Down Approach: Bulk material is broken down into nanoscale pieces. Wet Chemical Synthesis (Sol-Gel Process): Based on inorganic polymerization. Steps: hydrolysis, polycondensation, drying, thermal decomposition. Analysis/Characterization of Nanomaterials: Techniques like UV-Vis spectroscopy, XRD, SEM, TEM, FTIR are used to study size, crystal structure, morphology. History of Nanotechnology: Nanomaterials used for centuries (e.g., ruby red glass with gold nanoparticles). Term 'nanotechnology' coined by Nario Taniguchi (1974). Invention of STM (1980) led to discovery of fullerenes and carbon nanotubes. Applications of Nanomaterials: Stronger/lighter materials, scratchproof eyeglasses, sunscreens (ZnO, TiO$_2$), electronics (MRAM), water purification (AgNPs), self-cleaning materials (Lotus effect). Advantages: Revolution in electronics, economical solar power, efficient energy storage, smart drugs, cancer treatment. Disadvantages (Nanopollution): Potential risks to health (lung damage from inhaled nanoparticles), environmental pollution.