Chemistry
रसायन विज्ञान · Class 12 · Code 3021 (CDC Che.302)
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chapters done
- General and Physical Chemistry
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The syllabus asks for
- Introduction to gravimetric analysis, volumetric analysis and equivalent weight
- Relationship between equivalent weight, atomic weight and valency
- Equivalent weight of compounds: acid, base, salt, oxidizing and reducing agents
- Concentration units: percentage, g/L, molarity, molality, normality, formality, ppm and ppb
- Primary and secondary standard substances
- Law of equivalence and normality equation
- Titration and its types: acid-base titration, redox titration (with numerical problems)
Key points
- Gravimetric analysis weighs a precipitate; volumetric analysis measures the volume of a standard solution. Volumetric work is faster, gravimetric more accurate.
- E = molar mass/n-factor, and for an element E = atomic weight/valency. The n-factor is basicity for an acid, acidity for a base, total ionic charge for a salt, and electrons transferred for a redox reagent.
- KMnO₄ has three equivalent weights: 31.6 (acidic, 5e⁻), 52.67 (neutral, 3e⁻) and 158 (strongly alkaline, 1e⁻).
- Key concentration relations: M = w × 1000/(Mw Vcm^3), N = M × n, S = N × E, M = 10pd/Mw, m = 1000M/(1000d - MMw). Molality and percentage by mass do not change with temperature; molarity and normality do.
- A primary standard (oxalic acid, Na₂CO₃, K₂Cr₂O₇) is pure, stable and of high equivalent weight, so it can be weighed directly; a secondary standard (HCl, NaOH, KMnO₄, Na₂S₂O₃) must be titrated against one.
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The syllabus asks for
- Limitation of Arrhenius concept
- Bronsted-Lowry definition; Conjugate acid-base pairs
- Relative strength of acids and bases
- Lewis definition
- Ionization of weak electrolyte (Ostwald's dilution law)
- Ionic product of water (Kw)
- Dissociation constant of acid and base (Ka and Kb); concept of pKa and pKb
- pH value: pH of strong and weak acids; pH of strong and weak bases
- Solubility and solubility product principle
- Common ion effect
- Application of solubility product principle and common ion effect in precipitation reactions
- Buffer solution and its application
- Indicators and selection of indicators in acid-base titration
- Types of salts: acidic, basic, simple, complex
- Hydrolysis of salts: strong acid + strong base; weak acid + strong base; weak base + strong acid
Key points
- Arrhenius acids/bases work only in water; Bronsted-Lowry defines acid = proton donor, base = proton acceptor, giving conjugate pairs that differ by one H⁺; Lewis defines acid = electron-pair acceptor, base = electron-pair donor.
- For a conjugate pair Ka × Kb = Kw and pKa + pKb = 14; the stronger the acid, the weaker its conjugate base.
- Ostwald's dilution law: K = Cα^2/(1-α) ≈ Cα^2, so α = √K/C = √KV. It fails for strong electrolytes.
- Kw = [H^+][OH^-] = 1.0×10^-14 at 25 °C; pH + pOH = 14. Kw rises with temperature, so neutral pH falls below 7 on heating.
- pH master formulas: strong acid pH = -log C; weak acid pH = (1)/(2)(pKa - log C); strong base pH = 14 + log C; weak base pOH = (1)/(2)(pKb - log C).
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The syllabus asks for
- Introduction
- Rate of reactions: average and instantaneous
- Rate law and its expressions
- Rate constant, its unit and significance
- Order and molecularity
- Integrated rate equation for zero and first order reaction
- Half-life of zero and first order reactions
- Collision theory; activation energy and activated complex
- Factors affecting rate: concentration, temperature (Arrhenius equation), catalyst (energy profile diagram)
- Catalysis and types: homogeneous, heterogeneous and enzyme catalysis
Key points
- Kinetics deals with rates, the factors affecting them and mechanism. Average rate is the slope of a chord, instantaneous rate the slope of the tangent, on a concentration-time graph; the unit is mol L⁻¹ s⁻¹.
- For aA + bB → cC + dD the rate of reaction is -(1)/(a)(d[A])/(dt) = (1)/(c)(d[C])/(dt): always divide by the stoichiometric coefficient.
- The rate law r = k[A]^x[B]^y is experimental; its powers need not match the coefficients in the balanced equation.
- k is the specific reaction rate, has units (mol L^-1)^1-ns^-1, is independent of concentration and increases with temperature.
- Order is experimental and may be zero or fractional; molecularity is theoretical and is a small whole number. They are equal only for elementary reactions. Reactions like ester hydrolysis are pseudo-first-order because the solvent is in excess.
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The syllabus asks for
- Introduction; energy in chemical reactions
- Internal energy
- First law of thermodynamics
- Enthalpy and enthalpy changes (endothermic and exothermic)
- Enthalpy of reaction, solution, formation, combustion
- Laws of thermochemistry: Laplace law and Hess's law
- Entropy and spontaneity
- Second law of thermodynamics
- Gibbs free energy and prediction of spontaneity
- Relationship between dG and equilibrium constant
Key points
- A system exchanges matter and/or energy with its surroundings. U, H, S, G, p, V, T are state functions; q and w are not.
- First law: Δ U = q + w with w = -pextΔ V. At constant volume qv = Δ U; at constant pressure qp = Δ H.
- H = U + pV, so Δ H = Δ U + Δ ngRT for gas reactions. Δ H < 0 exothermic, Δ H > 0 endothermic.
- Δ Hr^° = ΣΔ Hf^°(products) - ΣΔ Hf^°(reactants), and Δ Hf^° of an element in its standard state is zero.
- Lavoisier–Laplace: reversing a reaction reverses the sign of Δ H. Hess's law: Δ H is route-independent — the basis of the Born–Haber cycle, which gives lattice enthalpy (−787 kJ mol⁻¹ for NaCl).
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The syllabus asks for
- Electrode potential and standard electrode potential
- Types of electrodes: standard hydrogen electrode and calomel electrode
- Electrochemical series and its applications
- Voltaic cell: Zn-Cu cell, Ag-Cu cell
- Cell potential and standard cell potential
- Relationship between cell potential and free energy
- Commercial batteries and fuel cells (hydrogen/oxygen)
Key points
- Electrode potential arises from the electrical double layer at a metal–solution boundary; only differences can be measured, so all E^° values are quoted against the SHE (1 M H⁺, H₂ at 1 bar, 298 K, defined as 0.000 V). The calomel electrode (+0.242 V saturated) is the practical secondary reference.
- The electrochemical series orders standard reduction potentials. More negative = better reducing agent; more positive = better oxidising agent. A metal displaces any metal below it, and metals above hydrogen liberate H₂ from dilute acids.
- In a galvanic cell, oxidation occurs at the negative anode (written on the left) and reduction at the positive cathode. Daniell cell: Zn | Zn²⁺(1 M) || Cu²⁺(1 M) | Cu, E^°cell = 1.10 V. Ag–Cu cell: Cu | Cu²⁺ || Ag⁺ | Ag, E^°cell = 0.46 V.
- E^°cell = E^°cathode - E^°anode; positive means spontaneous. E^° is intensive and is never multiplied by a stoichiometric coefficient.
- Nernst equation: Ecell = E^°cell - (0.0591)/(n)log Q at 298 K.
- Inorganic Chemistry
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The syllabus asks for
- Characteristics of transition metals
- Oxidation states
- Complex ions and metal complexes
- Shapes of complex ions
- d-orbitals in complex ions (crystal field theory, octahedral complex)
- Reasons for the colour of transition metal compounds
- Catalytic properties of transition metals
Key points
- A transition element has a partially filled d subshell in the atom or in a common ion. Zn (3d¹⁰4s², Zn²⁺ = 3d¹⁰) is therefore not typical. Cr and Cu have 4s¹ configurations because half-filled and full d subshells are stable.
- Characteristic properties: hardness, high melting point and density, variable oxidation states, coloured and paramagnetic ions, complex formation, catalysis, alloys and interstitial compounds.
- Variable oxidation states arise because 3d and 4s are close in energy. The maximum state rises to +7 at Mn, then falls; the +2 state dominates on the right.
- A complex ion is a central metal ion bonded to ligands by coordinate bonds. EAN = Z - (oxidation state) + 2(CN).
- Shapes: CN 2 linear (sp); CN 4 tetrahedral (sp³) or square planar (dsp²); CN 6 octahedral (d²sp³).
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The syllabus asks for
- Copper: occurrence and extraction from copper pyrite; properties with air, acids, aqueous ammonia, metal ions; uses; chemistry of blue vitriol; red oxide and black oxide (formula and uses)
- Zinc: occurrence and extraction from zinc blende; properties with air, acid, alkali, displacement; uses; chemistry of white vitriol
- Mercury: occurrence and extraction from cinnabar; properties; chemistry of calomel and corrosive sublimate
- Iron: occurrence and extraction; properties and uses; manufacture of steel by Basic Oxygen Method and Open Hearth Process; corrosion of iron and its prevention
- Silver: occurrence and extraction by cyanide process; preparation and uses of silver chloride and silver nitrate
Key points
- Copper comes from copper pyrite, CuFeS₂: froth flotation → roasting → smelting with silica (FeSiO₃ slag, Cu₂S + FeS matte) → bessemerisation, where 2Cu₂O + Cu₂S → 6Cu + SO₂ gives blister copper by auto-reduction → poling and electrolytic refining (anode mud holds Ag, Au, Pt).
- Copper is unattacked by dilute HCl or H₂SO₄ but dissolves in HNO₃ and in hot conc. H₂SO₄; with aqueous ammonia and air it gives deep blue Cu(NH₃)₄₂, Schweitzer's reagent. Blue vitriol is [Cu(H₂O)₄]SO₄·H₂O, white when anhydrous. Cu₂O is the red oxide, CuO the black.
- Zinc comes from zinc blende, ZnS: flotation → roasting (2ZnS + 3O₂ → 2ZnO + 2SO₂) → reduction with coke (ZnO + C → Zn + CO), the zinc distilling off as vapour. Zinc is amphoteric: it gives H₂ with both dilute acid and hot alkali. White vitriol is ZnSO₄·7H₂O.
- Mercury comes from cinnabar, HgS, by roasting alone: HgS + O₂ → Hg + SO₂, because HgO decomposes on heating. Excess mercury gives mercurous (Hg₂²⁺) salts, excess acid or halogen gives mercuric (Hg²⁺) salts. Calomel Hg₂Cl₂ is white, insoluble, harmless and blackened by ammonia; corrosive sublimate HgCl₂ is soluble, covalent and poisonous,
- Iron comes from haematite Fe₂O₃ in the blast furnace, the key step being FeO + CO → Fe + CO₂, with limestone as flux giving CaSiO₃ slag; the product is pig iron with ~4% C. Conc. HNO₃ makes iron passive.
- Organic Chemistry
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The syllabus asks for
- Introduction; nomenclature, isomerism and classification of monohaloalkanes
- Preparation from alkanes, alkenes and alcohols
- Physical properties
- Chemical properties; substitution reactions SN1 and SN2 (basic concept)
- Formation of alcohol, nitrile, amine, ether, thioether, carbylamines, nitrite and nitroalkane
- Elimination reaction (dehydrohalogenation, Saytzeff's rule); Reduction; Wurtz reaction
- Preparation of trichloromethane from ethanol and propanone
- Chemical properties of trichloromethane: oxidation, reduction, action on silver powder, conc. nitric acid, propanone, aqueous alkali
Key points
- A monohaloalkane is CₙH₂ₙ₊₁X. It is 1°, 2° or 3° according to the number of carbons on the halogen-bearing carbon, and shows chain, position and (when that carbon has four different groups) optical isomerism. C₄H₉Br has four structural isomers.
- Preparation: free radical halogenation of alkanes (gives mixtures); Markovnikov addition of HX to alkenes, reversed for HBr by peroxides; and — best — alcohols with conc. HX/ZnCl₂, PCl₃, PCl₅ or SOCl₂ (Darzen's process). Finkelstein (NaI/acetone) and Swarts (AgF) exchange halogens.
- Boiling point rises RF < RCl < RBr < RI and with chain length, and falls with branching. Dipole moments run CH₃Cl > CH₃F > CH₃Br > CH₃I.
- Reactivity in nucleophilic substitution is R–I > R–Br > R–Cl > R–F, because the C–X bond enthalpy falls in that order.
- S<sub>N</sub>2: one step, backside attack, rate = k[RX][Nu^-], transition state with five groups on carbon, inversion of configuration, favoured by CH₃X > 1° > 2° > 3° and strong nucleophiles.
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The syllabus asks for
- Introduction; nomenclature and isomerism
- Preparation of chlorobenzene from benzene and benzene diazonium chloride
- Physical properties
- Low reactivity vs haloalkanes in nucleophilic substitution
- Reduction of chlorobenzene; electrophilic substitution reactions
- Action with Na: Fittig and Wurtz-Fittig reaction; action with chloral
- Uses
Key points
- A haloarene has the halogen on an sp² ring carbon (C₆H₅Cl); an aralkyl halide has it on the side chain (C₆H₅CH₂Cl) and behaves like a haloalkane.
- C₆H₄Cl₂ has three positional isomers; the para isomer has much the highest melting point because of symmetrical crystal packing.
- Preparation: C₆H₆ + Cl₂ --(anhyd. AlCl₃, dark)--> C₆H₅Cl + HCl; industrially 2C₆H₆ + 2HCl + O₂ --(CuCl₂, 500 K)--> 2C₆H₅Cl + 2H₂O (Raschig).
- From the diazonium salt: Sandmeyer (CuCl/HCl) and Gattermann (Cu/HCl), both giving C₆H₅Cl + N₂↑; KI alone gives iodobenzene.
- Haloarenes are colourless, sweet-smelling, steam-volatile liquids, denser than and insoluble in water; boiling point rises Cl → Br → I.
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The syllabus asks for
- Introduction; nomenclature, isomerism and classification of monohydric alcohol
- Distinction of 1, 2, 3 alcohols by Victor Meyer's Method
- Preparation from haloalkane, primary amines and esters
- Industrial preparation: oxo process, hydroboration-oxidation of ethene, fermentation of sugar
- Common terms: absolute alcohol, power alcohol, denatured alcohol, rectified spirit, alcoholic beverage
- Physical properties
- Chemical properties: HX/PX3/PCl5/SOCl2; reactive metals; dehydration; oxidation of 1/2/3 alcohols; catalytic dehydrogenation; esterification; test of ethanol
Key points
- A monohydric alcohol is CₙH₂ₙ₊₁OH; it is 1°, 2° or 3° according to the number of carbon atoms (1, 2 or 3) attached to the carbinol carbon.
- Victor Meyer: HI/red P → AgNO₂ → HNO₂/NaOH gives red (1°, nitrolic acid), blue (2°, pseudonitrole), colourless (3°). Lucas: turbidity at once (3°), in 5 min (2°), none cold (1°).
- Preparations: R–X + aq. KOH; R–NH₂ + HNO₂ → R–OH + N₂; hydrolysis of esters. Industrial: oxo process (alkene + CO + H₂ → aldehyde → 1° alcohol, one extra carbon), hydroboration–oxidation (anti-Markovnikov), fermentation (C₆H₁₂O₆ --zymase--> 2C₂H₅OH + 2CO₂).
- Rectified spirit = 95.6 % azeotrope; absolute alcohol = 100 %, made with CaO; denatured alcohol contains methanol and pyridine; power alcohol is petrol + 20–25 % ethanol.
- High boiling point and water solubility come from intermolecular hydrogen bonding; both fall as the alkyl chain grows, and boiling point falls with branching.
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The syllabus asks for
- Introduction and nomenclature
- Preparation from chlorobenzene, diazonium salt, benzene sulphonic acid
- Physical properties
- Acidic nature: comparison with alcohol and water
- Action with NH3, Zn, Na, benzene diazonium chloride, phthalic anhydride
- Acylation, Kolbe's reaction, Reimer-Tiemann reaction
- Electrophilic substitution: nitration, sulphonation, bromination, Friedel-Crafts alkylation
- Test of phenol (FeCl3 test, aqueous bromine test, Libermann test); uses
Key points
- A phenol has –OH bonded directly to the ring; benzyl alcohol has it on a side chain and behaves as a 1° alcohol.
- Preparations: chlorobenzene + NaOH at 300 °C/150 atm (Dow); warming C₆H₅N₂Cl with water; fusion of C₆H₅SO₃Na with NaOH. Every route ends in sodium phenoxide and must be acidified.
- Phenol is a colourless solid, m.p. 41 °C, sparingly soluble in water, corrosive, and turns pink in air.
- Acidity: Ka = 1.0×10^-10, pKa = 10.0 — stronger than water (15.7) and ethanol (15.9) but weaker than carbonic acid (6.35). Hence it dissolves in NaOH but gives no CO₂ with NaHCO₃. The cause is resonance stabilisation of the phenoxide ion over C-2, C-4 and C-6.
- With NH₃ over ZnCl₂ at 300 °C → aniline; with Zn dust → benzene; with Na → sodium phenoxide + H₂; with C₆H₅N₂Cl in cold alkali → p-hydroxyazobenzene (orange dye); with phthalic anhydride → phenolphthalein.
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The syllabus asks for
- Introduction; nomenclature, classification and isomerism
- Preparation of aliphatic and aromatic ethers by Williamson's synthesis
- Physical properties
- Chemical properties of ethoxyethane: action with HI, conc. HCl, conc. H2SO4, air and Cl2
- Uses
Key points
- An ether is R–O–R′, formula CₙH₂ₙ₊₂O — a functional isomer of an alcohol. Simple ethers have identical groups, mixed ethers different ones.
- IUPAC names are alkoxyalkanes, the smaller group becoming the alkoxy substituent: C₂H₅OC₂H₅ is ethoxyethane, C₆H₅OCH₃ is methoxybenzene (anisole).
- Ethers show metamerism (different sharing of carbons about the oxygen), chain isomerism and functional isomerism with alcohols. C₄H₁₀O has 3 ethers and 4 alcohols.
- Williamson's synthesis: R–ONa + R′–X → R–O–R′ + NaX, an SN₂ reaction. The halide must be methyl or primary; a 3° halide gives an alkene and a haloarene gives nothing. For anisole use C₆H₅ONa + CH₃I.
- No O–H means no self-hydrogen-bonding, so ethoxyethane boils at 34.6 °C, about 83 °C below butan-1-ol and close to pentane. It is still water-soluble, because oxygen can accept a hydrogen bond from water.
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The syllabus asks for
- Aliphatic aldehydes and ketones: introduction, nomenclature, isomerism
- Preparation: dehydrogenation and oxidation of alcohol, ozonolysis of alkenes, acid chloride, gem dihaloalkane, catalytic hydration of alkynes
- Physical properties; structure and nature of carbonyl group
- Distinction by 2,4-DNP, Tollen's reagent, Fehling's solution
- Addition of H2, HCN, NaHSO3
- Action with ammonia derivatives: NH2OH, NH2-NH2, phenyl hydrazine, semicarbazide
- Aldol condensation; Cannizzaro's reaction
- Clemmensen's reduction; Wolf-Kishner reduction
- Action with PCl5 and LiAlH4; action of methanal with ammonia and phenol
- Formalin and its uses
- Aromatic: preparation of benzaldehyde from toluene and acetophenone from benzene
- Properties of benzaldehyde: Perkin condensation, Benzoin condensation, Cannizzaro's reaction, electrophilic substitution
Key points
- Aldehydes are R—CHO, ketones R—CO—R′; both are CₙH₂ₙO and are functional isomers of each other. Ketones also show position isomerism and metamerism.
- The carbonyl carbon is sp², planar, 120°, C=O = 122 pm, strongly polarised Cδ+=Oδ−. This makes the carbon electrophilic, the oxygen basic and the α-hydrogen acidic.
- Preparations: oxidation or Cu-dehydrogenation of alcohols (1° → aldehyde, 2° → ketone), reductive ozonolysis of alkenes, Rosenmund reduction of acid chlorides, hydrolysis of gem-dihalides, HgSO₄/H₂SO₄ hydration of alkynes (ethyne → ethanal, all others → ketone).
- Both give 2,4-DNP; only aldehydes give Tollens' silver mirror, Fehling's Cu₂O and Schiff's pink. Benzaldehyde gives Tollens but not Fehling. Iodoform tests for CH₃CO—, not for the aldehyde group.
- Nucleophilic addition: HCN → cyanohydrin, NaHSO₃ → bisulphite adduct, H₂/LiAlH₄ → alcohol. Reactivity HCHO > RCHO > RCOR′ > ArCHO, from +I and steric effects.
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The syllabus asks for
- Introduction, nomenclature, isomerism of aliphatic and aromatic carboxylic acids
- Preparation of monocarboxylic acids from aldehydes, nitriles, dicarboxylic acid, sodium alkoxide, trihaloalkanes
- Preparation of benzoic acid from alkyl benzene
- Physical properties
- Chemical properties: alkalies, metal oxides, metal carbonates, metal bicarbonates, PCl3, LiAlH4, dehydration
- Hell-Volhard-Zelinsky reaction
- Electrophilic substitution of benzoic acid: bromination, nitration, sulphonation
- Effect of constituents on acidic strength; abnormal behaviour of methanoic acid
- Derivatives (acid halides, amides, esters, anhydrides): preparation from carboxylic acid
- Comparative physical and chemical properties: hydrolysis, ammonolysis, amines, alcoholysis, reduction
- Claisen condensation; Hofmann bromamide reaction; amphoteric nature of amide; relative reactivity of acid derivatives
Key points
- Carboxylic acids are R—COOH, CₙH₂ₙO₂; the carboxyl carbon is always C-1. They are functional isomers of esters (C₃H₆O₂ = propanoic acid, methyl ethanoate, ethyl methanoate).
- Preparations: oxidation of a 1° alcohol or aldehyde, hydrolysis of a nitrile (+1 C), decarboxylation of a diacid (−1 C), carbonylation of a sodium alkoxide, hydrolysis of a trihaloalkane; benzoic acid from any alkylbenzene with a benzylic H by hot alkaline KMnO₄.
- They form hydrogen-bonded dimers, so they boil above alcohols of the same molar mass (ethanoic acid 118 °C against propan-1-ol 97 °C).
- Acidity: RCOOH > H₂CO₃ > C₆H₅OH > H₂O > ROH, because RCOO⁻ is resonance stabilised with both C—O bonds 127 pm. Effervescence with NaHCO₃ distinguishes an acid from a phenol.
- −I substituents strengthen the acid (pKa: CH₃COOH 4.76, ClCH₂COOH 2.87, Cl₂CHCOOH 1.29, Cl₃CCOOH 0.66); +I groups weaken it; the effect falls off sharply with distance.
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The syllabus asks for
- Nitroalkanes: introduction, nomenclature, isomerism
- Preparation from haloalkane and alkane; physical properties; reduction
- Nitrobenzene: preparation from benzene; physical properties
- Reduction in different media
- Electrophilic substitution: nitration, sulphonation, bromination
- Uses of nitro-compounds
Key points
- Nitroalkanes are R–NO₂ (CₙH₂ₙ₊₁NO₂); the nitro group is a resonance hybrid with two identical N–O bonds. Alkyl nitrites R–O–N=O are their functional isomers.
- Nitroalkanes show chain, position, functional and nitro–aci tautomeric isomerism. Only 1° and 2° (α-H present) are acidic and dissolve in NaOH.
- Preparation: R–X + AgNO₂ → R–NO₂ (with NaNO₂ you get the nitrite instead); CH₄ + HNO₃ at 400–450 °C → CH₃NO₂ + H₂O.
- Reduction: 6[H] gives R–NH₂; 4[H] with Zn/NH₄Cl stops at R–NHOH.
- Nitrobenzene: C₆H₆ + HNO₃ (conc. H₂SO₄, 50–60 °C) → C₆H₅NO₂ + H₂O, electrophile NO₂⁺. Pale yellow oil, b.p. 211 °C, bitter-almond smell, poisonous.
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The syllabus asks for
- Aliphatic amines: introduction, nomenclature, classification, isomerism
- Separation of 1, 2, 3 amines by Hoffmann's method
- Preparation of primary amines from haloalkane, nitriles, nitroalkanes and amides
- Physical properties; basicity and comparative basic nature of 1, 2, 3 amines
- Reaction of primary amines with chloroform, conc. HCl, R-X, RCOX and nitrous acid
- Test of 1, 2, 3 amines (nitrous acid test)
- Aniline: preparation from nitrobenzene and phenol; physical properties
- Basicity of aniline compared with aliphatic amines and ammonia
- Alkylation, acylation, diazotization, carbylamine and coupling reaction
- Electrophilic substitution: nitration, sulphonation, bromination; uses
Key points
- Amines are 1° (RNH₂), 2° (R₂NH) or 3° (R₃N) by the number of carbons on nitrogen; IUPAC names them alkanamines, with N- prefixes for groups on nitrogen. C₃H₉N has four isomers covering all three classes.
- Hoffmann's method separates a mixture with diethyl oxalate: 1° → solid oxamide, 2° → liquid oxamic ester, 3° → no reaction; alkaline hydrolysis returns each amine.
- Primary amines: C₂H₅Br + 2NH₃ → C₂H₅NH₂ + NH₄Br; CH₃CN + 4[H] → CH₃CH₂NH₂ (one more C); CH₃NO₂ + 6[H] → CH₃NH₂ + 2H₂O; CH₃CONH₂ + Br₂ + 4NaOH → CH₃NH₂ + Na₂CO₃ + 2NaBr + 2H₂O (one fewer C).
- Boiling point 1° > 2° > 3° at equal mass (hydrogen bonding); amines boil below alcohols of similar mass.
- Kb = [RNH3^+][OH^-]/[RNH2], and smaller pKb means stronger base. In water: (CH₃)₂NH > CH₃NH₂ > (CH₃)₃N > NH₃; in the gas phase 3° > 2° > 1° > NH₃.
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The syllabus asks for
- Introduction, general formula and examples of organolithium, organocopper and organocadmium compounds
- Nature of metal-carbon bond
- Grignard reagent: preparation using haloalkane and haloarene
- Reaction with water, aldehydes and ketones (preparation of 1, 2, 3 alcohols), carbon dioxide, HCN, RCN, ester and acid chloride
Key points
- An organometallic compound has a direct C–M bond; C₂H₅ONa and CH₃COONa do not qualify.
- Classes: R–Li (organolithium), R–Mg–X (Grignard), R₂CuLi (Gilman cuprate), R₂Cd (organocadmium). Reactivity RLi > RMgX > R₂Cd; R₂Cd converts RCOCl into a ketone and stops there.
- Carbon (2.55) is more electronegative than every one of these metals, so the bond is Cδ⁻–Mδ⁺. Ionic character: C–Li ≈ 46 %, C–Mg ≈ 32 %, C–Cd ≈ 17 %. More ionic means more carbanion-like and more reactive.
- Preparation: R–X + Mg in dry ether → RMgX; reactivity R–I > R–Br > R–Cl. The ether coordinates to Mg. Chlorobenzene needs THF at reflux, not ether.
- Any active hydrogen kills the reagent: CH₃MgI + H₂O → CH₄↑ + Mg(OH)I (the basis of Zerewitinoff's method).
- Applied Chemistry
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The syllabus asks for
- Polymers: addition and condensation polymers; elastomers and fibres; natural and synthetic polymers
- Synthetic polymers: polythene, PVC, Teflon, polystyrene, nylon and bakelite
- Dyes: introduction; types of dyes on the basis of structure and method of application
- Drugs: characteristics; natural and synthetic drugs; classification of common drugs; habit-forming drugs and drug addiction
- Pesticides: introduction to insecticides, herbicides and fungicides
Key points
- A polymer is many monomers joined; n is the degree of polymerisation and M(polymer) = n × M(monomer) for an addition polymer.
- Addition polymerisation needs an unsaturated monomer and loses nothing; condensation polymerisation needs bifunctional monomers and expels H₂O, HCl or NH₃ at each link.
- Elastomers stretch and recover (rubber), fibres are strong and thread-like (nylon), thermoplastics remould on heating (polythene), thermosets do not (bakelite).
- Key polymers: polythene (ethene), PVC (chloroethene), Teflon (tetrafluoroethene), polystyrene (styrene) — all addition; nylon-6,6 (hexamethylenediamine + adipic acid) and bakelite (phenol + methanal) — both condensation.
- A dye needs a chromophore (–N=N–, –NO₂, >C=O, quinonoid) for colour and an auxochrome (–OH, –NH₂, –SO₃H) to deepen it and fix it to the fibre.
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The syllabus asks for
- Introduction
- Raw materials for cement production
- Main steps in cement production: crushing and grinding, strong heating, final grinding
- Types of cement: OPC and PPC
- Portland cement process with flow-sheet diagram
- Cement Industry in Nepal
Key points
- Cement is a hydraulic binder of calcium silicates and aluminates; Aspdin patented Portland cement in 1824.
- Composition of OPC: CaO 60–67 %, SiO₂ 17–25 %, Al₂O₃ 3–8 %, Fe₂O₃ 0.5–6 %. Hydraulic modulus ≈ 2; silica modulus 2.5–4. Excess lime makes cement unsound.
- Raw materials: calcareous (limestone), argillaceous (clay), correctives (sand, bauxite, iron ore) and 2–5 % gypsum added after burning.
- Three stages: crushing/grinding (wet or dry process) → burning in a rotary kiln → final grinding with gypsum.
- Kiln zones: drying (to 400 °C), calcination 900–1000 °C (CaCO₃ → CaO + CO₂↑), clinkering 1400–1500 °C (formation of C₃S, C₂S, C₃A, C₄AF), then rapid cooling.
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The syllabus asks for
- Introduction
- Raw materials and their sources
- Stages in production of paper
- Flow-sheet diagram for paper production
- Quality of paper
Key points
- Paper is a mat of cellulose fibres, (C₆H₁₀O₅)ₙ, bonded by hydrogen bonds; pulping is the removal of lignin, which is brown, stiff and light-sensitive.
- Wood is roughly 45 % cellulose, 25 % hemicellulose and 25 % lignin; non-wood fibres such as babiyo grass and straw contain much less lignin.
- Raw materials: softwood (long fibre, strength), hardwood (short fibre, smoothness), non-wood fibre, rags and lokta bark, and recycled waste paper.
- Mechanical pulping gives 90–95 % yield but weak, yellowing pulp; chemical pulping gives 45–55 % yield of strong, stable fibre.
- Kraft (sulphate) process: cook chips with white liquor (NaOH + Na₂S) at 170 °C for 2–4 h; the sulphite process uses Ca(HSO₃)₂ with free SO₂.
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The syllabus asks for
- Natural and artificial radioactivity
- Units of radioactivity
- Nuclear reactions
- Nuclear fission and fusion reactions
- Nuclear power and nuclear weapons
- Industrial uses of radioactivity
- Medical uses of radioactivity
- Radiocarbon dating
- Harmful effects of nuclear radiations
Key points
- Radioactivity is the spontaneous disintegration of an unstable nucleus; discovered by Becquerel (1896), it is unaffected by temperature, pressure or chemical combination.
- α = ⁴₂He (high ionising, low penetrating), β = ⁰₋₁e, γ = electromagnetic (low ionising, high penetrating). α emission: A − 4, Z − 2; β emission: A same, Z + 1.
- Artificial radioactivity (Joliot-Curie, 1934): ²⁷₁₃Al + ⁴₂He → ³⁰₁₅P + ¹₀n, and ³⁰₁₅P → ³⁰₁₄Si + ⁰₊₁e.
- Units: activity 1 Bq = 1 dps, 1 Ci = 3.7 × 10¹⁰ Bq; absorbed dose 1 Gy = 1 J kg⁻¹ = 100 rad; equivalent dose 1 Sv = 100 rem.
- In every nuclear equation the mass numbers and the atomic numbers must balance; the mass defect appears as energy, with 1 u = 931.5 MeV.
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The paper
3hr, FM 75 written (+25 internal)
| Section | Questions | Marks each | Total |
|---|---|---|---|
| Group A MCQ | 11 | 1 | 11 |
| Group B short | 8 | 5 | 40 |
| Group C long | 3 | 8 | 24 |
Where the hours go
Teaching hours are the board's guide to how much of the paper each part carries.
- Organic Chemistry55 h
- General and Physical Chemistry40 h
- Inorganic Chemistry20 h
- Applied Chemistry13 h