Chemistry

13.1 Elements and Periodicity

13.1.1 Describe the demarcation of the periodic table into s, p, d, and f-blocks.

• The periodic table is divided into four blocks based on electron configuration:
  • s-Block: Includes Group 1 (alkali metals) and Group 2 (alkaline earth metals). The valence electrons occupy the s-orbital.
  • p-Block: Includes Groups 13 to 18, where valence electrons enter the p-orbital.
  • d-Block: Consists of transition metals (Groups 3-12) where valence electrons are in the d-orbital.
  • f-Block: Comprises lanthanides and actinides, where electrons are filling the f-orbital.

13.1.2 Determine the group, period, and block of given elements by using electronic configuration.

• Group: Determined by the number of valence electrons in the outermost shell.
• Period: Given by the principal quantum number (n) of the highest occupied energy level.
• Block: Identified by the type of orbital where the last electron enters:
  • If the last electron enters an s-orbital, the element belongs to the s-block.
  • If the last electron enters a p-orbital, it belongs to the p-block.
  • If the last electron enters a d-orbital, it belongs to the d-block.
  • If the last electron enters an f-orbital, it belongs to the f-block.

13.1.3 Explain the periodicity of physical properties of elements within groups and periods in the periodic table.

• Atomic Radius: Decreases across a period (due to increasing nuclear charge) and increases down a group (due to additional electron shells).
• Ionisation Energy: Increases across a period (stronger nuclear attraction) and decreases down a group (shielding effect reduces nuclear attraction).
• Electronegativity: Increases across a period (due to stronger nuclear charge) and decreases down a group (due to larger atomic radius).
• Electron Affinity: Becomes more negative across a period (greater nuclear attraction for additional electrons) and less negative down a group.
• Electrical Conductivity: High in metals due to free-moving electrons, but low in nonmetals.
• Melting and Boiling Points:
  • Metals: Decrease down a group as metallic bonds weaken.
  • Nonmetals: Increase down a group due to stronger Van der Waals forces.

13.2 Period 3 (Na to Ar)

13.2.1 List the elements in Period 3.

• Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulphur (S), Chlorine (Cl), Argon (Ar).

13.2.2 Describe the reaction of Period 3 elements with water, oxygen, and chlorine.

• Reaction with Water:
  • Sodium (Na): Reacts violently to form NaOH and H₂ gas.
  • Magnesium (Mg): Reacts slowly with cold water but rapidly with steam.
  • Aluminium (Al), Silicon (Si), Phosphorus (P), Sulphur (S), Chlorine (Cl), and Argon (Ar): Do not react with water.
• Reaction with Oxygen:
  • Na, Mg, and Al form basic oxides.
  • Si, P, S form acidic oxides.
  • Cl and Ar do not react with oxygen.
• Reaction with Chlorine:
  • Na, Mg, and Al form ionic chlorides.
  • Si, P form covalent chlorides.

13.2.3 Describe the reaction of oxides and chlorides of Period 3 elements with water.

• Oxides with Water:
  • Basic oxides (Na₂O, MgO) dissolve to form alkaline solutions.
  • Amphoteric oxide (Al₂O₃) does not dissolve easily.
  • Acidic oxides (SO₂, P₄O₁₀) form acids in water.
• Chlorides with Water:
  • Ionic chlorides (NaCl, MgCl₂) dissolve without reaction.
  • Covalent chlorides (SiCl₄, PCl₃) hydrolyze, forming acidic solutions.

13.2.4 Describe physical properties and acid-base characteristics of oxides and chlorides of Period 3 elements.

• Bonding:
  • Metals form ionic compounds.
  • Nonmetals form covalent compounds.
• Conductivity:
  • Ionic oxides and chlorides conduct electricity when molten.
  • Covalent compounds do not conduct electricity.
• Solubility:
  • Ionic compounds dissolve in water.
  • Covalent chlorides hydrolyze in water.
• Acid-Base Properties:
  • Basic oxides (Na₂O, MgO): React with acids.
  • Acidic oxides (SO₂, P₄O₁₀): React with bases.
  • Amphoteric oxides (Al₂O₃): React with both acids and bases.

13.3 Group 1 (Alkali Metals)

13.3.1 Describe oxidation states and trends in physical properties in Group 1 elements.

• Oxidation State: Always +1 (loses one electron easily).
• Atomic Radius: Increases down the group due to additional electron shells.
• Ionisation Energy: Decreases down the group as valence electrons are farther from the nucleus.
• Electronegativity: Decreases down the group as atomic size increases.
• Melting and Boiling Points: Decrease down the group due to weaker metallic bonding.

13.3.2 Describe the chemical reaction of Group 1 elements with water, oxygen, and chlorine.

• Reaction with Water: Produces metal hydroxide and hydrogen gas.
• Reaction with Oxygen: Forms oxides, peroxides, or superoxides depending on the metal.
• Reaction with Chlorine: Forms ionic chlorides (e.g., NaCl, KCl).

13.3.3 Discuss the trends in solubility of hydroxides, sulphates, and carbonates of Group 1 elements.

• Hydroxides: Solubility increases down the group (LiOH < NaOH < KOH).
• Sulphates: All Group 1 sulphates are soluble in water.
• Carbonates: Solubility increases down the group.

13.3.4 Discuss the trends in thermal stability of nitrates and carbonates of Group 1 elements.

• Nitrates: Decompose on heating, forming nitrites and oxygen.
• Carbonates: Do not decompose easily, except for lithium carbonate (Li₂CO₃ → Li₂O + CO₂).

13.4 Group 2 (Alkaline Earth Metals)

13.4.1 Describe oxidation states and trends in physical properties in Group 2 elements.

• Oxidation State: Group 2 elements always exhibit a +2 oxidation state (lose two valence electrons).
• Atomic Radius: Increases down the group due to additional electron shells.
• Ionisation Energy: Decreases down the group as the outer electrons are further from the nucleus.
• Electronegativity: Decreases down the group.
• Melting and Boiling Points: Generally decrease down the group due to weaker metallic bonding.

13.4.2 Describe the chemical reaction of Group 2 elements with water, oxygen, and nitrogen.

• Reaction with Water:
  • Reactivity increases down the group.
  • Magnesium reacts slowly with cold water but rapidly with steam.
  • Calcium, strontium, and barium react vigorously, forming hydroxides and hydrogen gas.
• Reaction with Oxygen: Forms basic oxides (e.g., MgO, CaO).
• Reaction with Nitrogen: Forms nitrides (e.g., Mg₃N₂, Ca₃N₂) when heated in nitrogen gas.

13.4.3 Compare the trends in solubility of hydroxides, sulphates, and carbonates of Group 2 with Group 1 elements.

• Hydroxides: Increase in solubility down the group (Mg(OH)₂ is sparingly soluble, Ba(OH)₂ is highly soluble).
• Sulphates: Decrease in solubility down the group (BaSO₄ is insoluble).
• Carbonates: Become less soluble down the group.
• Comparison with Group 1: Group 1 hydroxides and carbonates are more soluble than Group 2 compounds.

13.4.4 Discuss the trends in thermal stability of nitrates and carbonates of Group 2 elements.

• Nitrates: Decompose upon heating, forming oxides, nitrogen dioxide (NO₂), and oxygen.
• Carbonates: Decompose upon heating to form metal oxides and carbon dioxide.
• Thermal Stability: Increases down the group due to stronger lattice energy.

13.4.5 Differentiate beryllium from other members of its group.

• Beryllium:
  • Has a higher ionisation energy and small ionic radius.
  • Forms covalent compounds instead of ionic ones (e.g., BeCl₂ is covalent, unlike MgCl₂ which is ionic).
  • Does not react with water at room temperature.
  • Forms complex compounds similar to aluminium (diagonal relationship).

13.5 Group 4 (Carbon Group)

13.5.1 Describe variation in oxidation states and trends in physical properties of Group 4 elements.

• Oxidation States:
  • +4 oxidation state is common.
  • Heavier elements (Sn, Pb) also exhibit a +2 oxidation state due to the inert pair effect.
• Trends in Physical Properties:
  • Metallic Character: Increases down the group (C is nonmetal, Sn and Pb are metals).
  • Melting and Boiling Points: Decrease down the group (C has the highest melting point due to strong covalent bonding).
  • Ionisation Energy: Decreases down the group.

13.5.2 Describe the reaction of water with chlorides of carbon, silicon, and lead.

• Carbon Chloride (CCl₄): Does not react with water.
• Silicon Chloride (SiCl₄): Reacts with water to form silicic acid.
• Lead(II) Chloride (PbCl₂): Sparingly soluble in cold water but dissolves in hot water.

13.5.3 Compare the structure and stability of chlorides of carbon, silicon, and lead.

• CCl₄: Covalent, stable, non-hydrolyzing.
• SiCl₄: Covalent, hydrolyzes rapidly.
• PbCl₂: Ionic, moderately stable, sparingly soluble in water.

13.5.4 Describe the molecular structure of CO₂ and SiO₂.

• CO₂: Linear molecule, gas at room temperature, weak intermolecular forces.
• SiO₂: Giant covalent structure, high melting point, insoluble in water.

13.5.5 Discuss the acid-base characteristics of oxides of Group 4 elements.

• CO₂: Acidic oxide, dissolves in water to form carbonic acid (H₂CO₃).
• SiO₂: Weakly acidic, reacts with strong bases.
• PbO: Amphoteric, reacts with both acids and bases.

13.6 Group 7 (Halogens)

13.6.1 Discuss oxidation states and trends in physical properties of Group 7 elements.

• Oxidation States: -1 (most common), but can also exhibit +1, +3, +5, +7.
• Physical Properties:
  • Atomic Radius: Increases down the group.
  • Electronegativity: Decreases down the group (F is most electronegative).
  • Electron Affinity: Decreases down the group (F has the highest electron affinity).
  • Bond Energy: Decreases from Cl₂ to I₂ (F-F bond is weak due to repulsion of lone pairs).
  • Melting and Boiling Points: Increase down the group (due to stronger Van der Waals forces).

13.6.2 Discuss bond enthalpies and acidic strength of hydrogen halides.

• Bond Enthalpies: Decrease from HF to HI due to weaker bonding.
• Acidic Strength: Increases from HF to HI (HF < HCl < HBr < HI), as bond strength decreases.

13.6.3 Compare the strength of halide ions as reducing agents.

• Reducing Power: Increases down the group (I⁻ > Br⁻ > Cl⁻ > F⁻).
• Explanation: Larger ions lose electrons more easily, making them better reducing agents.

13.6.4 Explain the significance of halogens in daily life.

• Fluorine (F₂): Used in toothpaste (fluoride) to prevent tooth decay.
• Chlorine (Cl₂): Used for water purification and making disinfectants (bleach).
• Bromine (Br₂): Used in flame retardants and pharmaceuticals.
• Iodine (I₂): Essential for thyroid function (prevents goitre), used as an antiseptic.

14.1 General Features of Transition Elements

14.1.1 Describe the general features of transition elements (i.e., color, variable oxidation states, use as a catalyst).

• Transition elements are found in the d-block of the periodic table and include Groups 3-12.
• Color: Many transition metal compounds are colored due to the d-d electronic transitions.
• Variable Oxidation States: Transition metals exhibit multiple oxidation states due to the availability of partially filled d-orbitals (e.g., Fe²⁺ and Fe³⁺).
• Catalytic Properties: Transition metals act as catalysts because of their ability to change oxidation states and provide surfaces for reaction (e.g., Fe in Haber Process, V₂O₅ in Contact Process).

14.2 Electronic Structure

14.2.1 Explain the anomalous behavior of chromium and copper with respect to electronic configuration.

• Chromium (Cr) and Copper (Cu) show deviations from the expected Aufbau principle due to enhanced stability of half-filled (d⁵) and fully filled (d¹⁰) orbitals.
• Chromium (Cr): Expected configuration: [Ar] 4s² 3d⁴, but actual configuration: [Ar] 4s¹ 3d⁵.
• Copper (Cu): Expected configuration: [Ar] 4s² 3d⁹, but actual configuration: [Ar] 4s¹ 3d¹⁰.
• This happens because half-filled and fully filled orbitals are more stable due to symmetrical distribution and exchange energy.

14.2.2 Determine the electronic configuration of elements and ions of d-block.

• Electronic configuration is determined by following the Aufbau principle:
  • Sc: [Ar] 4s² 3d¹
  • Ti: [Ar] 4s² 3d²
  • Fe²⁺: [Ar] 3d⁶ (4s electrons are lost first)
  • Cu²⁺: [Ar] 3d⁹ (4s electrons are lost first)

14.3 Chemistry of Some Specific Transition Elements

14.3.1 Describe redox reactions and uses of vanadium, chromium, copper, manganese, and iron as catalysts.

• Vanadium (V₂O₅): Used as a catalyst in the Contact Process for sulfuric acid production.
• Chromium (Cr): Used in chrome plating and corrosion-resistant alloys.
• Copper (Cu): Essential in electrical wiring, coinage, and catalysis.
• Manganese (MnO₂): Used in batteries and the oxidation of alkenes.
• Iron (Fe): Used in the Haber Process for ammonia synthesis.

14.3.2 Describe properties of alloys with reference to the metals that compose them.

• Alloys are mixtures of metals that enhance properties such as strength, corrosion resistance, and conductivity.
• Examples:
  • Brass (Cu + Zn): Used in musical instruments.
  • Bronze (Cu + Sn): Used in statues and medals.
  • Steel (Fe + C): Stronger than pure iron, used in construction.

14.3.3 Describe the reaction of K₂Cr₂O₇ with FeSO₄ and H₂S.

• Reaction with FeSO₄:
  • K₂Cr₂O₇ + FeSO₄ + H₂SO₄ → Cr³⁺ + Fe³⁺ + SO₄²⁻ + H₂O
  • Dichromate ion oxidizes Fe²⁺ to Fe³⁺, turning from orange to green.
• Reaction with H₂S:
  • K₂Cr₂O₇ + H₂S + H₂SO₄ → Cr³⁺ + S + H₂O
  • H₂S is oxidized to sulfur (S), and dichromate turns green.

14.3.4 Describe the reaction of KMnO₄ with FeSO₄ and H₂S.

• Reaction with FeSO₄:
  • KMnO₄ + FeSO₄ + H₂SO₄ → Mn²⁺ + Fe³⁺ + SO₄²⁻ + H₂O
  • Permanganate ion oxidizes Fe²⁺ to Fe³⁺, turning from purple to colorless.
• Reaction with H₂S:
  • KMnO₄ + H₂S → Mn²⁺ + S + H₂O
  • Permanganate oxidizes H₂S to sulfur.

14.4 Coordination Compounds

14.4.1 Define the terms:

• Ligands: Ions or molecules that donate electron pairs to a metal center (e.g., NH₃, Cl⁻, H₂O).
• Coordination Number: The number of ligand donor atoms bonded to the metal.
• Coordination Sphere: The metal ion and its surrounding ligands enclosed in brackets.
• Chelates: Complexes where ligands form ring structures with metal ions (e.g., EDTA).

14.4.2 Describe different types of ligands.

• Monodentate Ligands: Donate one pair of electrons (e.g., NH₃, Cl⁻).
• Bidentate Ligands: Donate two pairs of electrons (e.g., ethylenediamine).
• Polydentate Ligands: Donate multiple electron pairs (e.g., EDTA⁴⁻).

14.4.3 Explain shapes, color, and nomenclature of coordination compounds.

• Shapes: Determined by coordination number:
  • 4-coordinate: Tetrahedral (e.g., [NiCl₄]²⁻) or square planar (e.g., [Pt(NH₃)₂Cl₂]).
  • 6-coordinate: Octahedral (e.g., [Fe(CN)₆]³⁻).
• Color: Due to d-d transitions, affected by ligand field strength.
• Nomenclature:
  • Cation named first, then anion (e.g., Potassium hexacyanoferrate(III) → K₃[Fe(CN)₆]).
  • Ligands named in alphabetical order (e.g., Tetraamminecopper(II) sulfate → [Cu(NH₃)₄]SO₄).

14.4.4 Relate the coordination number of ions to the crystal structure of the compound of which they are a part.

• Coordination number determines crystal structure:
  • CN = 4: Tetrahedral (e.g., Zn²⁺ complexes) or square planar (e.g., Pt²⁺ complexes).
  • CN = 6: Octahedral (e.g., Fe³⁺ complexes in [Fe(CN)₆]³⁻).

15.1 Coal as a Source of Organic Compounds

• 15.1.1 Explain the destructive distillation of coal.
  • Destructive distillation of coal is the process of heating coal in the absence of air, breaking it down into various useful products.
  • Process:
    1. Coal is heated to about 1000°C in an airtight chamber.
    2. It decomposes to yield gases, liquids, and a solid residue.
  • Products:
    • Coke: A porous, carbon-rich material used in steel production and as a fuel.
    • Coal Tar: A thick black liquid containing benzene, toluene, naphthalene, anthracene, and other organic compounds used in dyes, explosives, and medicines.
    • Ammonia Liquor: Contains ammonium salts, used in fertilizers.
    • Coal Gas: A fuel composed of methane, hydrogen, and carbon monoxide.
• 15.1.2 Explain coal as a source of both aliphatic and aromatic hydrocarbons.
  • Coal consists of complex organic molecules that decompose into simpler hydrocarbons when processed.
  • Aliphatic Hydrocarbons: Found in coal gas, including methane, ethane, propane, and butane.
  • Aromatic Hydrocarbons: Derived from coal tar, including benzene, toluene, xylene, and naphthalene.
  • Industrial Uses: These hydrocarbons serve as raw materials for making solvents, plastics, explosives, synthetic fibers, and pharmaceuticals.

15.2 Classification of Organic Compounds

• 15.2.1 Classify organic compounds on the basis of their structure.
  • Aliphatic Compounds: Straight or branched-chain hydrocarbons (alkanes, alkenes, alkynes).
  • Aromatic Compounds: Compounds containing benzene rings (e.g., benzene, phenol, aniline).
  • Heterocyclic Compounds: Cyclic compounds containing oxygen, nitrogen, or sulfur in the ring (e.g., furan, pyridine).
  • Saturated vs. Unsaturated:
    • Saturated: Contain only single bonds (alkanes).
    • Unsaturated: Contain one or more double or triple bonds (alkenes, alkynes).
• 15.2.2 Identify a molecule’s functional group.
  • Hydrocarbons: Alkanes (-C-C-), Alkenes (-C=C-), Alkynes (-C≡C-), Arenes (Benzene rings).
  • Oxygen-containing Groups:
    • Alcohols (-OH) → Methanol, Ethanol.
    • Ethers (-O-) → Diethyl ether.
    • Carboxylic Acids (-COOH) → Acetic acid.
    • Esters (-COO-) → Ethyl acetate.
    • Aldehydes (-CHO) → Formaldehyde, Benzaldehyde.
    • Ketones (-CO-) → Acetone.
  • Nitrogen-containing Groups:
    • Amines (-NH₂, -NHR, -NR₂) → Methylamine.
    • Amides (-CONH₂) → Acetamide.
    • Nitriles (-C≡N) → Ethanenitrile.
    • Nitro Compounds (-NO₂) → Nitrobenzene.
  • Halides: Fluorides (-F), Chlorides (-Cl), Bromides (-Br), Iodides (-I).
  • Sulfur-containing Groups:
    • Thiols (-SH) → Ethanethiol.
    • Sulphides (-S-) → Dimethyl sulfide.

15.3 Isomerism

• 15.3.1 Exemplify isomerism, stereo-isomerism, and structural isomerism.
  • Structural Isomerism: Different connectivity of atoms:
    • Chain Isomerism: Different arrangements of carbon chains (e.g., butane vs. isobutane).
    • Position Isomerism: Functional group at different positions (e.g., 1-propanol vs. 2-propanol).
    • Functional Isomerism: Same molecular formula, different functional groups (e.g., ethanol vs. dimethyl ether).
  • Stereo-isomerism: Same connectivity, different spatial arrangement:
    • Geometrical Isomerism (Cis-Trans): Occurs in alkenes and cyclic compounds due to restricted rotation (e.g., cis-but-2-ene vs. trans-but-2-ene).
    • Optical Isomerism: Compounds with chiral centers that rotate plane-polarized light.
• 15.3.2 Define chiral center.
  • A chiral center is a carbon atom bonded to four different groups, leading to non-superimposable mirror images (enantiomers).
  • Example: 2-butanol has a chiral center at C2.
• 15.3.3 Explain optical isomerism as a result of chiral center.
  • Optical isomers (enantiomers) have identical physical properties but rotate light in opposite directions.
  • Dextrorotatory (+) enantiomer rotates light clockwise, while Levorotatory (-) enantiomer rotates light counterclockwise.
• 15.3.4 Determine chiral centers in the structural formula of a molecule.
  • Identify carbon atoms bonded to four different groups.
  • Count the number of such carbons to determine possible optical isomers.
  • Example: Lactic acid (CH₃-CHOH-COOH) has one chiral center at C2, leading to two enantiomers.
• 15.3.5 Explain isomerism in alkyl halides, amines, alcohols, phenols, aldehydes, ketones, carboxylic acids, and esters.
  • Alkyl Halides: Exhibit chain, position, and functional isomerism (e.g., 1-chloropropane vs. 2-chloropropane).
  • Amines: Exhibit chain and position isomerism (e.g., ethylamine vs. propylamine).
  • Alcohols & Phenols: Exhibit positional and functional isomerism (e.g., ethanol vs. dimethyl ether).
  • Aldehydes & Ketones: Show functional isomerism (e.g., propanal and acetone).
  • Carboxylic Acids & Esters: Exhibit functional and positional isomerism (e.g., acetic acid vs. methyl formate).

16.1 Nomenclature, Shape of Molecules, and Resonance

16.1.1 Describe the nomenclature and shapes of molecules (i.e., alkane, alkene, cycloalkane, alkynes, benzenes, and substituted benzene) based on sigma and pi carbon-carbon bonds.

• Nomenclature:
  • Alkanes: Named using the prefix indicating the number of carbon atoms and the suffix -ane.
  • Alkenes: Contain double bonds, named with the -ene suffix.
  • Alkynes: Contain triple bonds, named with the -yne suffix.
  • Cycloalkanes: Saturated hydrocarbons in a ring structure, named with the cyclo- prefix.
  • Arenes (Aromatic Compounds): Include benzene and its derivatives.
• Shapes of Molecules:
  • Alkanes: Tetrahedral geometry (sp³ hybridization).
  • Alkenes: Trigonal planar geometry (sp² hybridization).
  • Alkynes: Linear geometry (sp hybridization).
  • Benzene and Aromatic Compounds: Planar structure with delocalized π-electrons.

16.1.2 Explain the phenomenon of resonance and stability of benzene.

• Resonance in Benzene:
  • Benzene has a delocalized π-electron system, leading to resonance stability.
  • It is represented as two equivalent resonance structures.
  • The bond lengths in benzene are equal (intermediate between single and double bonds).
• Stability of Benzene:
  • Resonance energy makes benzene less reactive than typical alkenes.
  • Prefers electrophilic substitution rather than addition reactions.

16.2 Types of Organic Reactions

16.2.1 Define different types of organic reactions.

• Substitution Reaction: A functional group in a molecule is replaced by another.
  • Example: Halogenation of alkanes (CH₄ + Cl₂ → CH₃Cl + HCl).
• Elimination Reaction: The removal of atoms or groups to form a double or triple bond.
  • Example: Dehydration of alcohols (C₂H₅OH → C₂H₄ + H₂O).
• Addition Reaction: Two reactants combine to form one product.
  • Example: Hydrogenation of alkenes (C₂H₄ + H₂ → C₂H₆).
• Radical Reaction: Involves free radicals as reactive intermediates.
  • Example: Free radical halogenation of methane.

16.3 Alkanes

16.3.1 Explain the unreactive nature of alkanes towards polar reagents.

• Reasons for Low Reactivity:
  • Alkanes are saturated hydrocarbons with only sigma bonds.
  • No partial charges, making them non-polar and unattractive to polar reagents.
  • High C-H and C-C bond energies, requiring a large amount of energy to break.

16.4 Alkenes

16.4.1 Describe the preparation of ethene (using chemical equations) from:

• Dehydration of Alcohol:
  • Ethanol is heated with concentrated sulfuric acid (H₂SO₄) to remove water.
  • Equation: C₂H₅OH → C₂H₄ + H₂O
• Dehydrohalogenation of Alkyl Halide:
  • Alkyl halides react with a strong base (KOH in ethanol) to eliminate HX.
  • Equation: C₂H₅Br + KOH → C₂H₄ + KBr + H₂O

16.4.2 Describe the reactions of ethene.

• Hydrogenation: Ethene reacts with H₂ in the presence of Ni/Pt catalyst to form ethane.
• Hydration: Ethene reacts with H₂O in the presence of H₂SO₄ to form ethanol.
• Hydrohalogenation: Ethene reacts with HX (e.g., HCl, HBr) to form alkyl halides.
• Halogenation: Ethene reacts with Br₂ or Cl₂ to form dihalogen compounds.
• Halohydration: Ethene reacts with X₂ and H₂O to form halohydrin.
• Epoxidation: Ethene reacts with peracids (e.g., peracetic acid) to form epoxides.
• Ozonolysis: Ethene reacts with O₃ followed by Zn/H₂O to form aldehydes or ketones.
• Polymerisation: Ethene undergoes polymerization to form polyethylene.

16.5 Alkynes

16.5.1 Compare the reactivities of alkynes, alkenes, and alkanes.

• Alkanes: Least reactive due to single bonds (sigma bonds) and absence of electron-rich centers.
• Alkenes: More reactive than alkanes because of pi bonds, which are electron-rich and susceptible to electrophilic attack.
• Alkynes: Most reactive among them due to two pi bonds, making them highly unsaturated and prone to addition reactions.

16.5.2 Describe the preparation of alkynes using elimination reaction.

• Dehydrohalogenation of vicinal dihalides:
  • CH₃CHBrCHBrCH₃ + 2NaNH₂ → CH₃C≡CCH₃ + 2NaBr + 2NH₃
• Dehalogenation of tetrahalides:
  • CHBr₂CHBr₂ + 2Zn → CH≡CH + 2ZnBr₂

16.5.3 Explain the acidic strength of alkynes with reference to its reaction with metals.

• Alkynes have acidic hydrogen due to sp-hybridized carbon, which holds electrons closer, making hydrogen more electronegative and easily removed.
• React with metals to form acetylides:
  • HC≡CH + Na → HC≡C⁻Na⁺ + ½H₂
  • HC≡CH + AgNO₃ → HC≡CAg + HNO₃

16.5.4 Explain the chemistry of alkynes by hydrogenation, hydrohalogenation, hydration, bromination, and ozonolysis.

• Hydrogenation: Converts alkynes to alkenes (Lindlar catalyst) or alkanes (Pd, Pt, or Ni catalyst).
  • CH≡CH + H₂ → CH₂=CH₂ → CH₃CH₃
• Hydrohalogenation: Addition of HX follows Markovnikov’s rule.
  • CH≡CH + HBr → CH₂=CHBr → CH₃CHBr₂
• Hydration (Keto-enol tautomerism): Forms enol first, then rearranges to ketone.
  • CH≡CH + H₂O → CH₂=CHOH → CH₃CHO
• Bromination: Forms dibromoalkenes.
  • CH≡CH + Br₂ → CHBr=CHBr
• Ozonolysis: Cleaves alkyne into carboxylic acids.
  • CH≡CH + O₃ → 2HCOOH

16.5.5 Discuss the combustion reactions of alkanes, alkenes, and alkynes with reference to energy production.

• Alkanes: Burn to produce CO₂ and H₂O with high energy output.
• Alkenes: Also undergo complete combustion but give slightly less energy.
• Alkynes: Higher oxygen demand, may produce soot due to incomplete combustion.
  • CH≡CH + 2.5O₂ → 2CO₂ + H₂O + Energy

16.6 Benzene and Substituted Benzene

16.6.1 Compare the reactivity of benzene with alkene and alkane.

• Alkanes: Undergo substitution reactions due to strong sigma bonds.
• Alkenes: Undergo addition reactions due to highly reactive pi bonds.
• Benzene: More stable due to resonance, prefers electrophilic substitution rather than addition.

16.6.2 Describe the mechanism of electrophilic substitution reaction of benzene.

• Step 1: Generation of electrophile (E⁺).
• Step 2: Electrophile attacks benzene, forming an arenium ion.
• Step 3: Loss of proton, restoring aromaticity.
  • Example: Nitration of benzene.
    • C₆H₆ + HNO₃ → C₆H₅NO₂ + H₂O (in presence of H₂SO₄)

16.6.3 Explain orientation in benzene with reference to resonating structures, i.e., effect of ortho, meta, and para directing groups in electrophilic substitution reactions.

• Ortho/Para Directors: Electron-donating groups (-OH, -CH₃, -NH₂) activate the ring.
• Meta Directors: Electron-withdrawing groups (-NO₂, -CN, -COOH) deactivate the ring.
• Example: In nitration of toluene, NO₂ attaches to ortho and para positions.

16.6.4 Discuss the chemistry of benzene and methyl benzene by nitration, sulphonation, halogenation, Friedel-Crafts alkylation, and acylation.

• Nitration: C₆H₆ + HNO₃ → C₆H₅NO₂ + H₂O (in presence of H₂SO₄).
• Sulphonation: C₆H₆ + SO₃ → C₆H₅SO₃H (in presence of H₂SO₄).
• Halogenation: C₆H₆ + Cl₂ → C₆H₅Cl + HCl (in presence of FeCl₃).
• Friedel-Crafts Alkylation: C₆H₆ + RCl → C₆H₅R + HCl (in presence of AlCl₃).
• Friedel-Crafts Acylation: C₆H₆ + RCOCl → C₆H₅COR + HCl (in presence of AlCl₃).

17.1 Alkyl Halides

17.1.1 Apply IUPAC and trivial systems for naming alkyl halides.

• Nomenclature Rules:
  • Identify the longest carbon chain containing the halogen.
  • Number the chain so that the halogen gets the lowest possible number.
  • Use prefixes fluoro-, chloro-, bromo-, iodo- to indicate halogens.
  • Example: CH₃CH₂Br → Bromoethane.

17.1.2 Discuss physical properties and reactivity of different alkyl halides on the basis of bond energy.

• Physical Properties:
  • Boiling Point: Increases with molecular weight and polarity.
  • Solubility: Insoluble in water but soluble in organic solvents.
• Reactivity:
  • C-X Bond Strength:
    • C-F > C-Cl > C-Br > C-I (Decreasing bond strength, increasing reactivity).
  • Polar Nature: More polar bonds are more reactive in nucleophilic substitution.

17.1.3 Draw the structure of different alkyl halides using their formulae.

• Example Structures:
  • CH₃Cl → Methyl chloride
  • CH₃CH₂Br → Ethyl bromide
  • CH₃CH₂CH₂I → Propyl iodide

17.1.4 Describe the preparation of alkyl halides by the reaction of alcohol with HX, SOCl₂, PCl₃, and PCl₅.

• Reaction with Hydrogen Halides (HX):
  • CH₃CH₂OH + HBr → CH₃CH₂Br + H₂O
• Reaction with Thionyl Chloride (SOCl₂):
  • CH₃CH₂OH + SOCl₂ → CH₃CH₂Cl + SO₂ + HCl
• Reaction with Phosphorus Halides:
  • 3CH₃OH + PCl₃ → 3CH₃Cl + H₃PO₃
  • CH₃OH + PCl₅ → CH₃Cl + POCl₃ + HCl

17.2 Nucleophilic Substitution Reactions

17.2.1 Describe the mechanism of nucleophilic substitution (SN1 and SN2) reactions.

• SN1 (Unimolecular Nucleophilic Substitution):
  • Occurs in two steps:
    1. Formation of carbocation (slow step).
    2. Nucleophilic attack (fast step).
  • Example: CH₃C⁺H₂ + H₂O → CH₃CH₂OH
• SN2 (Bimolecular Nucleophilic Substitution):
  • Occurs in one step:
    • Nucleophile attacks while the leaving group departs simultaneously.
  • Example: CH₃CH₂Br + OH⁻ → CH₃CH₂OH + Br⁻

17.2.2 Discuss carbocation and its stability.

• Carbocation Stability Order:
  • 3° (tertiary) > 2° (secondary) > 1° (primary) > CH₃⁺
  • Stability increases due to inductive effect and hyperconjugation.

17.2.3 Compare SN1 and SN2 reactions.

17.2.4 Deduce the mechanism of nucleophilic substitution (SN1 and SN2) reaction for the given alkyl halide.

• Given an alkyl halide, identify whether it follows SN1 or SN2 mechanism.
  • Example: (CH₃)₃CBr undergoes SN1, while CH₃CH₂Br undergoes SN2.

17.2.5 Identify nucleophile (base), substrate, and leaving group in the given nucleophilic substitution reactions.

• Example: CH₃CH₂Br + OH⁻ → CH₃CH₂OH + Br⁻
  • Nucleophile: OH⁻
  • Substrate: CH₃CH₂Br
  • Leaving Group: Br⁻

17.3 Elimination Reactions

17.3.1 Describe the mechanism of elimination (E1 and E2) reactions.

• E1 (Unimolecular Elimination):
  • Two-step process:
    1. Formation of carbocation (slow step).
    2. Elimination of H⁺ to form an alkene (fast step).
  • Example: CH₃CHBrCH₃ → CH₃CH=CH₂ + HBr
• E2 (Bimolecular Elimination):
  • One-step process:
    • Base removes β-hydrogen, causing simultaneous leaving group departure.
  • Example: CH₃CH₂Br + OH⁻ → CH₂=CH₂ + H₂O + Br⁻

17.3.2 Compare E1 and E2 reactions.

17.3.3 Deduce the mechanism of elimination (E1 and E2) reaction for the given alkyl halide.

• Given an alkyl halide, determine if it follows E1 or E2 mechanism.
  • Example: (CH₃)₃CBr undergoes E1, while CH₃CH₂Br undergoes E2.

17.3.4 Compare substitution reaction with elimination reaction.

• Substitution: Replaces halogen with nucleophile.
• Elimination: Forms alkene by removing HX.
• SN1 vs. E1: Compete in similar conditions.
• SN2 vs. E2: Compete in strong base conditions.

17.4 Organo-Metallic Compounds (Grignard Reagent)

17.4.1 Describe the preparation and reactivity of Grignard reagent.

• Preparation:
  • Grignard reagent (RMgX) is formed by reacting alkyl or aryl halides (RX) with magnesium metal in dry ether.
  • Equation: R-X + Mg → R-Mg-X (in dry ether)
• Reactivity:
  • Highly reactive due to carbon-metal bonding, making it a strong nucleophile.
  • Readily reacts with electrophilic compounds like CO₂, aldehydes, ketones, esters.

17.4.2 Describe chemical reaction of Grignard reagent with aldehydes, ketones, esters, and carbon dioxide.

• Reaction with Aldehydes & Ketones: Forms alcohols.
  • Equation (Aldehyde): R-Mg-X + HCHO → R-CH₂OH (Primary alcohol)
  • Equation (Ketone): R-Mg-X + R’CO → R-CHOH-R’ (Tertiary alcohol)
• Reaction with Esters: Forms tertiary alcohols.
  • Equation: R-Mg-X + R’-COOR” → R-COH-R’R”
• Reaction with CO₂: Forms carboxylic acids.
  • Equation: R-Mg-X + CO₂ → R-COOH

17.5 Amines

17.5.1 Apply IUPAC and trivial system for naming amines.

• IUPAC Naming:
  • Primary amines (R-NH₂): Named by adding “amine” to the alkyl group.
    • Example: CH₃NH₂ → Methylamine
  • Secondary amines (R₂NH): Named by using prefix “N-” for the smaller alkyl group.
    • Example: N-Methylethanamine
  • Tertiary amines (R₃N): Named by listing all alkyl groups in alphabetical order.
    • Example: Trimethylamine (N,N-Dimethylethanamine)

17.5.2 Discuss physical properties of amines (melting point, boiling point, and solubility).

• Melting and Boiling Point:
  • Higher than hydrocarbons due to hydrogen bonding.
  • Primary amines > Secondary amines > Tertiary amines (in boiling point, due to reduced hydrogen bonding).
• Solubility:
  • Lower amines (1° and 2°) are soluble in water due to hydrogen bonding.
  • Tertiary amines are less soluble due to the absence of N-H bonds.

17.5.3 Draw the structure of amines (primary, secondary, and tertiary) from their formulae.

• Primary Amine (R-NH₂):
  • Example: CH₃NH₂ (Methylamine)
• Secondary Amine (R₂NH):
  • Example: CH₃-NH-CH₃ (Dimethylamine)
• Tertiary Amine (R₃N):
  • Example: N(CH₃)₃ (Trimethylamine)

17.5.4 Explain basicity (basic character) of amines.

• Amines are basic due to the presence of a lone pair on nitrogen.
• Order of basicity in aqueous solutions:
  • Aliphatic amines > Ammonia > Aromatic amines.
• Factors affecting basicity:
  • Electron-donating groups increase basicity.
  • Electron-withdrawing groups decrease basicity.

17.5.5 Describe the preparation of amines by:

• Alkylation of NH₃:
  • Ammonia reacts with alkyl halides to form amines.
  • Equation: NH₃ + RX → RNH₂ + HX
• Reduction of Nitriles:
  • Nitriles (R-C≡N) are reduced to primary amines using LiAlH₄ or H₂/Ni.
  • Equation: R-C≡N + 4[H] → R-CH₂NH₂
• Reduction of Nitro Compounds:
  • Nitroarenes (Ar-NO₂) are reduced to aromatic amines using Sn/HCl.
  • Equation: Ar-NO₂ + 6[H] → Ar-NH₂ + 2H₂O
• Reduction of Amides:
  • Amides (R-CONH₂) are reduced to amines.
  • Equation: R-CONH₂ + 4[H] → R-CH₂NH₂ + H₂O

17.5.6 Describe the chemical reaction of amines, i.e.:

• Alkylation with RX:
  • Amines react with alkyl halides to form higher amines.
  • Equation: R-NH₂ + RX → R-NH-R + HX
• Reaction with Aldehydes and Ketones:
  • Amines form Schiff bases (imines) with carbonyl compounds.
  • Equation: R-NH₂ + R’CHO → R-CH=NR’ + H₂O

17.5.7 Describe the preparation of amides and diazonium salts.

• Preparation of Amides:
  • Carboxylic acids react with ammonia or amines to form amides.
  • Equation: R-COOH + NH₃ → R-CONH₂ + H₂O
• Preparation of Diazonium Salts:
  • Aromatic amines react with NaNO₂ and HCl at 0-5°C to form diazonium salts.
  • Equation: Ar-NH₂ + NaNO₂ + HCl → Ar-N₂⁺Cl⁻ + H₂O

18.1 Alcohols

18.1.1 Apply IUPAC and trivial system for naming different alcohols.

• Alcohols are named using the IUPAC system, where the suffix “-ol” is added to the longest carbon chain containing the hydroxyl (-OH) group.
• Examples:
  • Methanol (CH₃OH) – The simplest alcohol, commonly used as an industrial solvent.
  • Ethanol (CH₃CH₂OH) – Found in alcoholic beverages and used as a fuel additive.
  • 2-Propanol (CH₃CHOHCH₃) – Also known as isopropanol, used as rubbing alcohol.

18.1.2 Describe the physical properties and structure of alcohol.

• Structure: Alcohols consist of a hydroxyl (-OH) group attached to a saturated carbon.
• Physical Properties:
  • Polarity: Alcohols are polar due to the presence of the hydroxyl (-OH) group.
  • Hydrogen Bonding: Alcohol molecules form strong hydrogen bonds, leading to higher boiling points than alkanes of similar molecular weight.
  • Solubility: Lower alcohols (e.g., methanol, ethanol) are highly soluble in water, but solubility decreases with increasing hydrocarbon chain length.

18.1.3 Distinguish among primary, secondary, and tertiary alcohols using Lucas reagent test.

• Lucas Test: Differentiates alcohols based on their reactivity with Lucas reagent (ZnCl₂ + HCl):
  • Primary alcohols – No immediate reaction, remain clear.
  • Secondary alcohols – Cloudiness appears within 5 minutes.
  • Tertiary alcohols – Immediate turbidity (cloudy solution).

18.1.4 Differentiate between methanol and ethanol using iodoform test (haloform reaction).

• Iodoform Test: Used to identify ethanol and other compounds with a CH₃CO- group.
  • Ethanol gives a yellow precipitate (CHI₃ – iodoform).
  • Methanol does not react, as it lacks the necessary structure.

18.1.5 Describe the preparation of alcohol by reduction of aldehydes, ketones, carboxylic acids, and esters using chemical equations.

• Reduction Reactions:
  • Aldehydes → Primary Alcohols
    • CH₃CHO + 2[H] → CH₃CH₂OH (Ethanol from Acetaldehyde using NaBH₄/LiAlH₄)
  • Ketones → Secondary Alcohols
    • CH₃COCH₃ + 2[H] → CH₃CH(OH)CH₃ (Isopropanol from Acetone)
  • Carboxylic Acids → Primary Alcohols
    • CH₃COOH + 4[H] → CH₃CH₂OH + H₂O (Ethanol from Acetic Acid)
  • Esters → Alcohols
    • CH₃COOCH₃ + 4[H] → CH₃CH₂OH + CH₃OH (Ethanol and Methanol from Methyl Acetate)

18.1.6 Discuss the acidic character of alcohols.

• Alcohols are weak acids, with acidity decreasing in the order:
  • Methanol > Primary Alcohol > Secondary Alcohol > Tertiary Alcohol
• The hydroxyl group (-OH) can donate a proton (H⁺), but this occurs weakly compared to strong acids.
• Reaction with sodium metal demonstrates acidity:
  • 2CH₃OH + 2Na → 2CH₃O⁻Na⁺ + H₂ (Methanol reacting with sodium to form sodium methoxide and hydrogen gas)

18.1.7 Describe the reactions of alcohol, i.e.

• Preparation of Ethers: Alcohols react under acidic conditions to form ethers.
  • CH₃CH₂OH + CH₃CH₂OH → CH₃CH₂OCH₂CH₃ + H₂O (Diethyl ether formation)
• Preparation of Esters: Alcohols react with carboxylic acids in the presence of acid catalysts (H₂SO₄).
  • CH₃CH₂OH + CH₃COOH → CH₃COOCH₂CH₃ + H₂O (Ethyl Acetate formation)
• Oxidative Cleavage of 1,2-diols:
  • CH₂(OH)-CH₂(OH) + [O] → 2HCHO (Glycol cleavage to form formaldehyde)

18.1.8 Define thiols (RSH).

• Thiols (mercaptans) are sulfur analogs of alcohols, where oxygen is replaced by sulfur (-SH group).
• Example: Methanethiol (CH₃SH) – responsible for the smell in natural gas.

18.1.9 Describe the uses of alcohol as disinfectant and antiseptic.

• Ethanol and Isopropanol are used as disinfectants and antiseptics due to their ability to denature proteins and dissolve lipid membranes of bacteria and viruses.
• Common Applications:
  • Hand sanitizers (contain 60–70% ethanol/isopropanol).
  • Medical antiseptics for skin preparation before injections.
  • Disinfection of medical instruments.

18.2 Phenols

18.2.1 Apply IUPAC and trivial system for naming different phenols.

• Phenols are aromatic compounds containing a hydroxyl (-OH) group directly attached to a benzene ring.
• IUPAC Naming: The hydroxyl group is assigned position 1 in the benzene ring.
  • Example: Phenol (C₆H₅OH)
  • Example: 2-Methylphenol (o-Cresol)
• Trivial Names:
  • Hydroxybenzene = Phenol
  • Hydroxytoluene = Cresol

18.2.2 Discuss the physical properties, structure, and acidic characteristics of phenol (with reference to its resonance only).

• Physical Properties:
  • State: Colorless crystalline solid
  • Boiling Point: High due to hydrogen bonding
  • Solubility: Soluble in water due to hydrogen bonding, but decreases with alkyl substitution.
• Structure of Phenol:
  • The -OH group in phenol participates in resonance, increasing electron density at ortho and para positions.
  • This affects its reactivity in electrophilic substitution reactions.
• Acidic Character:
  • Phenol is more acidic than alcohols due to resonance stabilization of the phenoxide ion (C₆H₅O⁻).
  • Example: C₆H₅OH ⇌ C₆H₅O⁻ + H⁺ (Phenol ionizes partially in water)

18.2.3 Describe the preparation of phenols from the given compounds using chemical equations.

• From Benzene Sulphonic Acid:
  • C₆H₆SO₃H + NaOH → C₆H₅ONa + H₂O (Fusion)
  • C₆H₅ONa + HCl → C₆H₅OH + NaCl
• From Chlorobenzene:
  • C₆H₅Cl + NaOH → C₆H₅ONa + H₂O (High temperature & pressure)
  • C₆H₅ONa + HCl → C₆H₅OH + NaCl
• Acidic Oxidation of Cumene:
  • C₆H₅CH(CH₃)₂ + O₂ → C₆H₅C(OOH)(CH₃)₂ (Cumene hydroperoxide formation)
  • C₆H₅C(OOH)(CH₃)₂ + H₂SO₄ → C₆H₅OH + CH₃COCH₃ (Phenol and Acetone formation)
• Hydrolysis of Diazonium Salts:
  • C₆H₅N₂⁺Cl⁻ + H₂O → C₆H₅OH + N₂ + HCl

18.2.4 Discuss the reactivity of phenol with reference to electrophilic aromatic substitution, reaction with Na metal, and oxidation.

• Electrophilic Aromatic Substitution:
  • Phenol undergoes nitration, halogenation, and sulfonation due to increased electron density at ortho-para positions.
  • Example: C₆H₅OH + Br₂ → 2,4,6-Tribromophenol (White precipitate)
• Reaction with Sodium Metal:
  • Phenol reacts with sodium to form sodium phenoxide and hydrogen gas.
  • 2C₆H₅OH + 2Na → 2C₆H₅ONa + H₂
• Oxidation of Phenol:
  • Phenol oxidizes to form quinones in the presence of oxidizing agents.
  • C₆H₅OH + [O] → C₆H₄O₂ + H₂O (Benzoquinone formation)
• Comparison with Alcohols:
  • Phenols are more reactive due to the aromatic system.
  • Alcohols lack resonance stabilization and do not undergo electrophilic substitution.

18.3 Ethers

18.3.1 Apply IUPAC and trivial system for naming different ethers.

• Ethers are organic compounds with an oxygen atom bonded to two alkyl or aryl groups (R-O-R’).
• IUPAC Naming:
  • Smaller alkyl group is named as an alkoxy group.
  • Larger alkyl group is considered the main chain.
  • Example: Methoxyethane (CH₃OCH₂CH₃)
• Trivial Names:
  • Dimethyl Ether (CH₃OCH₃)
  • Diethyl Ether (C₂H₅OC₂H₅)

18.3.2 Describe the physical and chemical properties of ethers.

• Physical Properties:
  • Boiling Point: Lower than alcohols due to absence of hydrogen bonding.
  • Solubility: Slightly soluble in water due to weak dipole interactions.
  • Density: Less dense than water.
• Chemical Properties:
  • Inert Nature: Ethers are generally unreactive due to the absence of a reactive functional group.
  • Peroxide Formation: Ethers react with oxygen to form explosive peroxides.
  • Acidic Cleavage: Strong acids like HI or HBr cleave ethers into alkyl halides.
  • Example: RO-R’ + HI → RI + R’OH

18.3.3 Describe the preparation of ethers by the following methods using chemical equations.

• Williamson Ether Synthesis:
  • Reaction of alkoxide ions (RO⁻) with alkyl halides (R’-X).
  • RO⁻ + R’X → R-O-R’ + X⁻
  • Example: CH₃ONa + CH₃CH₂Br → CH₃OCH₂CH₃ + NaBr
• Reaction of Alkyl Halides with Dry Silver Oxide:
  • 2R-X + Ag₂O → R-O-R + 2AgX
  • Example: C₂H₅Br + Ag₂O → C₂H₅OC₂H₅ + AgBr
• Reaction of Alcohols with Excess H₂SO₄:
  • 2R-OH + H₂SO₄ → R-O-R + H₂O
  • Example: 2CH₃OH → CH₃OCH₃ + H₂O (Methanol to Dimethyl Ether)

18.3.4 Describe the use of ether in the field of medicine.

• Anesthetic:
  • Diethyl ether (C₂H₅OC₂H₅) was widely used as a general anesthetic.
• Solvent:
  • Used in pharmaceutical extractions and drug formulations.
• Chemical Reagent:
  • Used in the Grignard reaction as a solvent.
• Fuel Additive:
  • Methyl tert-butyl ether (MTBE) is used as an octane booster in gasoline.

19.1 Nomenclature and Structure

19.1.1 Apply IUPAC and trivial system for naming aldehydes and ketones.

• Aldehydes contain the -CHO functional group, and their IUPAC names end in “-al”.
  • Example: Methanal (Formaldehyde) → HCHO
  • Example: Ethanal (Acetaldehyde) → CH₃CHO
• Ketones contain the C=O (carbonyl group) bonded to two carbon atoms, with IUPAC names ending in “-one”.
  • Example: Propanone (Acetone) → CH₃COCH₃
  • Example: Butanone → CH₃COCH₂CH₃

19.1.2 Draw the structure of given aldehydes and ketones.

• Structures:
  • Aldehydes: Carbonyl group (C=O) at the end of the carbon chain.
    • Example: HCHO (Methanal), CH₃CHO (Ethanal)
  • Ketones: Carbonyl group (C=O) within the carbon chain.
    • Example: CH₃COCH₃ (Propanone), CH₃COCH₂CH₃ (Butanone)

19.1.3 Describe glucose and fructose as examples of aldehydes and ketones.

• Glucose is an aldose (contains an aldehyde group at C1).
  • Molecular Formula: C₆H₁₂O₆
  • Structure: Pyranose form in aqueous solutions.
• Fructose is a ketose (contains a ketone group at C2).
  • Molecular Formula: C₆H₁₂O₆
  • Structure: Furanose form in aqueous solutions.

19.2 Physical Properties

19.2.1 Explain the physical properties of aldehydes and ketones.

• Boiling Point:
  • Higher than alkanes but lower than alcohols due to dipole-dipole interactions.
  • Example: Propanal (48°C) < Propanol (97°C)
• Solubility:
  • Lower aldehydes and ketones are soluble in water due to hydrogen bonding.
  • Solubility decreases as the carbon chain increases.
• Odor:
  • Aldehydes have strong, pungent smells.
  • Ketones have sweet, pleasant smells.

19.3 Preparation of Aldehydes and Ketones

19.3.1 Describe the preparation of aldehydes and ketones by:

• Ozonolysis of Alkenes:
  • Reaction: Alkene reacts with ozone (O₃), followed by Zn/H₂O workup to form aldehydes or ketones.
  • Example:
    • CH₂=CH₂ + O₃ → CH₂O + CH₂O (Ethene to Formaldehyde)
    • CH₃CH=CH₂ + O₃ → CH₃CHO + HCHO (Propene to Acetaldehyde & Formaldehyde)
• Hydration of Alkynes:
  • Reaction: Alkyne reacts with H₂O in presence of Hg²⁺ and H₂SO₄, forming an enol, which tautomerizes into an aldehyde or ketone.
  • Example:
    • CH≡CH + H₂O → CH₃CHO (Ethyne to Acetaldehyde)
    • CH≡CCH₃ + H₂O → CH₃COCH₃ (Propyne to Acetone)
• Oxidation of Alcohols:
  • Primary alcohols oxidize to aldehydes.
    • Example: CH₃CH₂OH + [O] → CH₃CHO + H₂O (Ethanol to Acetaldehyde)
  • Secondary alcohols oxidize to ketones.
    • Example: CH₃CHOHCH₃ + [O] → CH₃COCH₃ + H₂O (Propan-2-ol to Acetone)
  • Common oxidizing agents: PCC, K₂Cr₂O₇, KMnO₄.
• Friedel-Crafts Acylation of Aromatic Compounds:
  • Reaction: Aromatic compounds react with acid chlorides (RCOCl) in the presence of AlCl₃, forming ketones.
  • Example:
    • C₆H₆ + CH₃COCl → C₆H₅COCH₃ + HCl (Benzene to Acetophenone)

19.4 Reactions of Aldehydes and Ketones

19.4.1 Describe the base-catalyzed nucleophilic addition reaction of aldehydes and ketones, i.e.

• Addition of Hydrogen Cyanide (HCN):
  • Forms cyanohydrins by adding CN⁻ to the carbonyl carbon.
  • Reaction: CH₃CHO + HCN → CH₃CH(OH)CN (Acetaldehyde to Cyanohydrin).
• Addition of Grignard Reagent:
  • Grignard reagents (RMgX) add to the carbonyl group, forming alcohols.
  • Reaction: CH₃CHO + CH₃MgBr → CH₃CH(OH)CH₃.
• Addition of Sodium Bisulphate (NaHSO₃):
  • Aldehydes and ketones form bisulfite addition compounds, useful in purification.
  • Reaction: CH₃CHO + NaHSO₃ → CH₃CH(OH)SO₃Na.
• Aldol Condensation:
  • Two aldehyde or ketone molecules react in the presence of a base to form a β-hydroxy aldehyde or ketone, which can further dehydrate.
  • Reaction: 2CH₃CHO → CH₃CH(OH)CH=CH₂ (Aldol product from Acetaldehyde).
• Cannizzaro’s Reaction:
  • Non-enolizable aldehydes undergo disproportionation in a strong base to give an alcohol and a carboxylate.
  • Reaction: 2HCHO + NaOH → CH₃OH + HCOONa (Formaldehyde to Methanol & Sodium Formate).
• Haloform (Iodoform) Reaction:
  • Methyl ketones react with iodine and a base to form a yellow iodoform precipitate (CHI₃).
  • Reaction: CH₃COCH₃ + 3I₂ + 4NaOH → CHI₃ (yellow ppt.) + CH₃COONa + 3NaI + H₂O.

19.4.2 Describe the acid-catalyzed nucleophilic addition reaction of aldehydes and ketones, i.e.

• Polymerization:
  • Aldehydes undergo acid-catalyzed polymerization to form cyclic or linear polymers.
  • Example: Formaldehyde polymerizes into paraformaldehyde.
• Addition of Ammonia Derivatives:
  • Aldehydes and ketones react with hydroxylamine (NH₂OH), hydrazine (NH₂NH₂), and semicarbazide (NH₂CONHNH₂).
  • Reaction: CH₃CHO + NH₂OH → CH₃CH=NOH (Oxime Formation).
• Addition of Alcohols:
  • Forms hemiacetals and acetals under acid catalysis.
  • Reaction: CH₃CHO + CH₃OH → CH₃CH(OH)OCH₃ (Hemiacetal Formation).

19.4.3 Describe the reduction of aldehydes and ketones using:

• Clemmensen Reduction Method:
  • Uses Zn(Hg) + HCl to reduce carbonyl compounds to alkanes.
  • Reaction: CH₃COCH₃ + Zn(Hg) + HCl → CH₃CH₂CH₃ (Acetone to Propane).
• Wolff-Kishner Reduction Method:
  • Uses NH₂NH₂ + KOH (heat) to reduce aldehydes/ketones to alkanes.
  • Reaction: CH₃CHO + NH₂NH₂ + KOH → CH₃CH₃ + N₂ + H₂O (Acetaldehyde to Ethane).
• Hydride Reagents:
  • Uses NaBH₄ or LiAlH₄ to reduce aldehydes to primary alcohols and ketones to secondary alcohols.
  • Reaction: CH₃CHO + NaBH₄ → CH₃CH₂OH (Ethanol from Acetaldehyde).
• Carbon Nucleophiles:
  • Grignard reagents (RMgX) add to the carbonyl group forming alcohols.
  • Reaction: CH₃CHO + CH₃MgBr → CH₃CH(OH)CH₃.
• Nitrogen Nucleophiles:
  • Aldehydes and ketones react with hydroxylamine, hydrazine, and semicarbazide.
  • Reaction: CH₃CHO + NH₂OH → CH₃CH=NOH (Oxime formation).
• Oxygen Nucleophiles:
  • Alcohols form hemiacetals and acetals under acid catalysis.
  • Reaction: CH₃CHO + CH₃OH → CH₃CH(OH)OCH₃.

19.4.4 Describe the oxidation reactions of aldehydes and ketones.

• Aldehydes:
  • Easily oxidized to carboxylic acids using KMnO₄ or K₂Cr₂O₇.
  • Reaction: CH₃CHO + [O] → CH₃COOH (Acetaldehyde to Acetic Acid).
• Ketones:
  • Resistant to oxidation but undergo cleavage with strong oxidizers.
  • Reaction: CH₃COCH₃ + [O] → CH₃COOH + HCOOH (Acetone cleavage).

19.5 Uses and Effects

19.5.1 Discuss the uses of formaldehyde in daily life.

• Preservative: Used in formalin (40% HCHO solution) for preserving biological specimens.
• Disinfectant: Used in hospitals and industry.
• Plastic Manufacturing: Essential in Bakelite and Urea-Formaldehyde Resins.

19.5.2 Discuss the health hazards associated with exposure to formalin.

• Respiratory Issues: Can cause coughing, irritation, and asthma.
• Carcinogenic Effects: Long-term exposure may increase the risk of cancer.
• Skin and Eye Irritation: Causes redness, burns, and allergic reactions.

20.1 Nomenclature

20.1.1 Apply IUPAC and trivial system for naming carboxylic acids and their derivatives.

• Carboxylic Acids: Named by replacing the -e of the parent alkane with -oic acid.
  • Example: Methanoic acid (Formic acid), Ethanoic acid (Acetic acid).
• Acid Derivatives:
  • Acyl Halides: Replace -ic acid with -yl halide.
    • Example: Ethanoyl chloride (Acetyl chloride).
  • Esters: Replace -ic acid with -ate, naming the alkyl group first.
    • Example: Methyl ethanoate (Methyl acetate).
  • Amides: Replace -oic acid with -amide.
    • Example: Ethanamide (Acetamide).
  • Acid Anhydrides: Named by replacing acid with anhydride.
    • Example: Ethanoic anhydride (Acetic anhydride).

20.2 Structure and Physical Properties

20.2.1 Describe the structure and physical properties (solubility, melting point, and boiling point) of carboxylic acids.

• Structure: Carboxylic acids contain a carboxyl (-COOH) functional group, with a polar C=O and O-H bond.
• Solubility:
  • Lower carboxylic acids are soluble in water due to hydrogen bonding.
  • Solubility decreases with increasing carbon chain length.
• Melting and Boiling Points:
  • Higher than alcohols due to strong hydrogen bonding.
  • Increase with molecular weight due to increased van der Waals forces.

20.2.2 Draw the structure of given compounds of carboxylic acids and their derivatives.

• Examples:
  • Carboxylic Acid: CH₃COOH (Ethanoic acid)
  • Acid Chloride: CH₃COCl (Ethanoyl chloride)
  • Ester: CH₃COOCH₃ (Methyl ethanoate)
  • Amide: CH₃CONH₂ (Ethanamide)
  • Anhydride: (CH₃CO)₂O (Ethanoic anhydride)

20.3 Acidity

20.3.1 Discuss the acidity of carboxylic acids.

• Carboxylic acids are weak acids, partially ionizing in water to form R-COO⁻ and H⁺.
  • Equation: R-COOH ⇌ R-COO⁻ + H⁺
• Factors affecting acidity:
  • Resonance Stabilization: Carboxylate ion (R-COO⁻) is stabilized by delocalization of electrons, increasing acidity.
  • Inductive Effect:
    • Electron-withdrawing groups (-Cl, -NO₂) increase acidity.
    • Electron-donating groups (-CH₃, -OH) decrease acidity.
• Comparison:
  • Stronger acids than alcohols and phenols.
  • Acidity trend: Formic acid > Acetic acid > Propionic acid.

20.4 Preparation of Carboxylic Acids

20.4.1 Describe the preparation of carboxylic acids by:

• From Grignard Reagents:
  • Reaction of Grignard reagent (RMgX) with CO₂, followed by acidic hydrolysis.
  • Equation: RMgX + CO₂ → RCOO⁻MgX → RCOOH
• Hydrolysis of Nitriles:
  • Nitriles hydrolyze in the presence of acid or base to form carboxylic acids.
  • Equation: R-CN + H₂O + H⁺ → RCOOH + NH₄⁺
• Oxidation of Primary Alcohols and Aldehydes:
  • Strong oxidizing agents (KMnO₄, K₂Cr₂O₇) oxidize primary alcohols or aldehydes to carboxylic acids.
  • Equation: RCH₂OH → RCHO → RCOOH
• Oxidation of Alkyl Benzenes:
  • Benzyl groups (-CH₃) in aromatic compounds oxidize to carboxyl groups (-COOH).
  • Equation: C₆H₅CH₃ + [O] → C₆H₅COOH

20.5 Reactivity

20.5.1 Describe the reactions of carboxylic acid involving:

• Hydrogen atom of the carboxyl group:
  • Carboxylic acids undergo reduction to form primary alcohols using LiAlH₄.
  • Equation: RCOOH + 4[H] → RCH₂OH + H₂O
• Hydroxyl group of carboxyl group:
  • Formation of acid chlorides with thionyl chloride (SOCl₂).
  • Equation: RCOOH + SOCl₂ → RCOCl + SO₂ + HCl
• Carboxyl group as a whole:
  • Carboxylic acids react with alcohols to form esters (esterification reaction).
  • Equation: RCOOH + R’OH ⇌ RCOOR’ + H₂O (in the presence of H₂SO₄)

20.5.2 Compare the reactivity of different derivatives of carboxylic acid (i.e. acyl halides, acid anhydrides, esters, and amides).

• Reactivity order: Acyl halides > Acid anhydrides > Esters > Amides.
• Acyl halides are the most reactive due to strong electron-withdrawing effects.
• Amides are the least reactive due to resonance stabilization.

20.6 Reactions of Carboxylic Acid Derivatives

20.6.1 Describe the preparation of acyl halides, acid anhydrides, esters, and amides.

• Acyl Halides: Prepared by reacting carboxylic acid with SOCl₂.
• Equation: RCOOH + SOCl₂ → RCOCl + SO₂ + HCl
• Acid Anhydrides: Formed by dehydration of two carboxylic acid molecules.
• Equation: 2RCOOH → RCO-O-COR + H₂O
• Esters: Prepared by esterification reaction between carboxylic acid and alcohol.
• Amides: Formed by reaction of ammonia or amines with carboxylic acids.

20.6.2 Describe the inter-conversion reactions of carboxylic acid derivatives.

• Acyl halides react with water to form carboxylic acids.
• Esters hydrolyze into carboxylic acids in acidic or basic medium.
• Amides undergo hydrolysis to form carboxylic acids in the presence of H₂SO₄.

20.6.3 Describe the reactions of carboxylic acid derivatives, i.e.:

• Friedel-Crafts acylation using acyl halide:
  • Acyl halides react with benzene in the presence of AlCl₃ catalyst.
  • Equation: RCOCl + C₆H₆ → C₆H₅COR + HCl
• Hydrolysis of acid anhydrides, esters, and amides:
  • Acid anhydrides react with water to form carboxylic acids.
  • Esters hydrolyze to form alcohols and acids.
  • Equation: RCOOR’ + H₂O → RCOOH + R’OH
• Reduction of esters and amides:
  • Esters reduce to primary alcohols using LiAlH₄.
  • Amides reduce to amines.
• Reaction of Grignard reagent with esters:
  • Esters react with Grignard reagent (RMgX) to form tertiary alcohols.

20.7 Uses

20.7.1 Identify carboxylic acids present in fruits and vegetables.

• Citric acid: Found in lemons and oranges.
• Malic acid: Found in apples and grapes.
• Oxalic acid: Found in spinach and rhubarb.
• Tartaric acid: Found in tamarind and grapes.

20.7.2 Describe the uses of carboxylic acids in industries.

• Plastic Industry: Used in polyesters and synthetic fibers.
• Leather Industry: Used in tanning processes.
• Rubber Industry: Used for vulcanization.
• Soap Industry: Fatty acids are used in soap production.
• Food Industry: Used as preservatives and flavoring agents (e.g., acetic acid in vinegar).

21.1 Carbohydrates, Proteins, and Lipids

21.1.1 Describe the basis of classification of carbohydrates, proteins, and lipids.

• Carbohydrates: Classified into monosaccharides, disaccharides, and polysaccharides based on the number of sugar units.
  • Monosaccharides: Simple sugars (e.g., glucose, fructose).
  • Disaccharides: Two monosaccharides linked together (e.g., sucrose, lactose).
  • Polysaccharides: Long chains of monosaccharides (e.g., starch, cellulose).
• Proteins: Classified based on structure and function.
  • Simple proteins: Composed only of amino acids (e.g., albumin, globulin).
  • Conjugated proteins: Contain a non-protein part (e.g., hemoglobin).
• Lipids: Classified into simple, compound, and derived lipids.
  • Simple lipids: Fats and oils.
  • Compound lipids: Contain additional groups like phosphates (e.g., phospholipids).
  • Derived lipids: Products of hydrolysis of simple and compound lipids (e.g., steroids).

21.1.2 Describe the structure of carbohydrates, proteins, and lipids.

• Carbohydrates: Contain C, H, and O, usually in a 1:2:1 ratio.
  • Example: Glucose (C₆H₁₂O₆) has a six-membered ring structure.
• Proteins: Composed of amino acids linked by peptide bonds.
  • Primary structure: Sequence of amino acids.
  • Secondary structure: Alpha helices and beta sheets.
  • Tertiary structure: Three-dimensional folding.
• Lipids: Consist of glycerol and fatty acids.
  • Saturated fatty acids: No double bonds.
  • Unsaturated fatty acids: Contain double bonds.

21.1.3 Explain the role of carbohydrates in health and disease.

• Energy Source: Glucose is the primary source of energy for cells.
• Dietary Fiber: Polysaccharides like cellulose aid in digestion.
• Diabetes: Excess carbohydrate consumption can lead to insulin resistance and Type 2 diabetes.
• Obesity: High sugar intake contributes to weight gain.

21.1.4 Discuss the nutritional importance of proteins and lipids.

• Proteins: Essential for growth, repair, and enzyme formation.
  • Deficiency causes kwashiorkor and marasmus.
• Lipids: Provide long-term energy storage, insulation, and hormone production.
  • Essential fatty acids like omega-3 and omega-6 support heart health.

21.1.5 Explain different types of lipids (simple, compound, derived or associated including steroids).

• Simple Lipids: Fats and oils (e.g., triglycerides).
• Compound Lipids: Phospholipids, glycolipids.
• Derived Lipids: Steroids, cholesterol, fat-soluble vitamins (A, D, E, K).

21.1.6 Describe the effect of lowering pH (using lemon juice) on milk proteins.

• Acidic pH causes protein denaturation, leading to curdling.
• Casein, the primary milk protein, precipitates in acidic conditions.

21.1.7 Describe the role of biochemical compounds such as insulin and cholesterol in the human body.

• Insulin: Regulates blood glucose levels by facilitating glucose uptake into cells.
• Cholesterol: Essential for cell membrane stability, hormone production, and vitamin D synthesis.

21.2 Enzymes

21.2.1 Describe the role of enzymes as biological catalysts, i.e., in digestion of food.

• Enzymes speed up biochemical reactions without being consumed.
• Example: Amylase breaks down starch into maltose.
• Example: Pepsin digests proteins in the stomach.

21.2.2 Explain the factors that affect enzyme activity.

• Temperature: Optimal range (e.g., 37°C in humans).
• pH: Each enzyme works best at a specific pH (e.g., pepsin at pH 2).
• Substrate concentration: Higher substrate levels increase activity up to saturation point.
• Inhibitors: Competitive and non-competitive inhibitors reduce enzyme function.

21.2.3 Explain the role of inhibitors in enzyme-catalyzed reactions.

• Competitive Inhibitors: Bind to the active site, blocking substrate binding (e.g., sulfa drugs against bacteria).
• Non-Competitive Inhibitors: Bind elsewhere, altering enzyme shape (e.g., heavy metals like lead).

21.3 Nucleic Acids

21.3.1 Differentiate between the structure of DNA and RNA.

• DNA: Double-stranded, contains deoxyribose and thymine (T).
• RNA: Single-stranded, contains ribose and uracil (U) instead of thymine.

21.3.2 Describe the role of:

• DNA in storing genetic information:
  • Contains genes that determine traits.
  • Replicates before cell division.
• RNA in protein synthesis:
  • mRNA: Carries genetic code from DNA to ribosomes.
  • tRNA: Transfers amino acids to ribosomes.
  • rRNA: Forms ribosomes for protein assembly.

21.4 Minerals of Biological Significance

21.4.1 Describe the role of iron and phosphorus as nutrients.

• Iron: Essential for hemoglobin production, oxygen transport, and enzyme function.
  • Deficiency causes anemia.
• Phosphorus: Key component of ATP, DNA, and bones.
  • Found in dairy, meat, and nuts.

22.1 Introduction

22.1.1 Discuss the importance of chemical industries for the economy of Pakistan.

• Chemical industries play a vital role in economic development by providing raw materials for various sectors such as agriculture, pharmaceuticals, textiles, and manufacturing.
• Major contributions include job creation, foreign exchange earnings, and industrial growth.
• Pakistan has a growing chemical sector, with industries focusing on fertilizers, cement, plastics, petrochemicals, and pharmaceuticals.

22.1.2 List the raw materials available in Pakistan for various chemical and petrochemical industries.

• Natural Gas: Used in fertilizer and petrochemical industries.
• Crude Oil: Processed into petroleum products.
• Limestone: Used in cement manufacturing.
• Rock Salt: Extracted for chlor-alkali production.
• Coal: Utilized in power plants and synthetic fuel production.
• Sulfur: Used in sulfuric acid and fertilizer industries.

22.2 Safety Measurement

22.2.1 List safety precautions that should be followed in chemical industries.

• Personal Protective Equipment (PPE): Use of gloves, masks, safety goggles, and protective clothing.
• Proper Ventilation: Ensuring air circulation to avoid the accumulation of toxic fumes.
• Emergency Procedures: Installation of fire extinguishers, eyewash stations, and emergency exits.
• Chemical Storage: Proper labeling and segregation of hazardous chemicals.
• Waste Management: Safe disposal of toxic by-products and industrial effluents.

22.3 Dyes and Pigments

22.3.1 Describe the types of dyes.

• Natural Dyes: Derived from plants, minerals, and animals (e.g., indigo, turmeric, cochineal).
• Synthetic Dyes: Chemically manufactured, classified as:
  • Acid Dyes: Used for wool, silk, and nylon.
  • Basic Dyes: Used for acrylic fibers and paper.
  • Direct Dyes: Applied to cotton without a mordant.
  • Disperse Dyes: Used for polyester and acetate fabrics.
  • Reactive Dyes: Form covalent bonds with cellulose fibers.

22.3.2 Discuss the importance of dyes and pigments in cosmetic, textile, paints, and food industries.

• Cosmetic Industry: Used in lipsticks, hair dyes, nail polishes.
• Textile Industry: Enhances the color and durability of fabrics.
• Paint Industry: Provides color stability and weather resistance.
• Food Industry: Used in processed foods, soft drinks, and confectionery.

22.4 Petrochemicals

22.4.1 Describe the process of:

• Fractional Distillation:
  • Separation of crude oil into components based on boiling points.
  • Yields petrol, diesel, kerosene, and lubricating oils.
• Refining of Petroleum:
  • Removal of impurities such as sulfur and nitrogen compounds.
  • Converts crude oil into usable fuels and petrochemical feedstocks.

22.4.2 Explain the processes of cracking (with its types) and reforming of petroleum.

• Cracking: Breaking large hydrocarbon molecules into smaller, more useful ones.
  • Thermal Cracking: Uses high temperature and pressure.
  • Catalytic Cracking: Uses catalysts like zeolites to break down hydrocarbons.
• Reforming:
  • Converts low-octane hydrocarbons into high-octane fuels.
  • Uses catalysts like platinum or rhenium.

22.4.3 Identify (in a given equation) the petrochemicals and chemicals derived from them (monomer and polymer).

• Ethylene (C₂H₄): Used to make polyethylene, ethanol, and ethylene glycol.
• Propylene (C₃H₆): Used in polypropylene and acrylonitrile.
• Benzene (C₆H₆): Used for styrene, phenol, and aniline production.
• Toluene (C₆H₅CH₃): Precursor for benzene derivatives and TNT.

22.4.4 List some major petrochemicals.

• Ethylene, Propylene, Benzene, Toluene, Xylene, Methanol, Ammonia.

22.5 Synthetic Polymers (PVC and Nylon)

22.5.1 Describe the chemical processes of addition and condensation polymerisation.

• Addition Polymerisation:
  • Monomers with double bonds join together without the loss of any molecule.
  • Example: Polyvinyl Chloride (PVC) formation from vinyl chloride (C₂H₃Cl).
• Condensation Polymerisation:
  • Monomers join together with the elimination of a small molecule like H₂O or HCl.
  • Example: Nylon-6,6 formation from hexamethylenediamine and adipic acid.

22.5.2 Describe the formation and uses of polyvinyl chloride (PVC) and nylon.

• PVC Formation:
  • Polymerisation of vinyl chloride (CH₂=CHCl) in the presence of a free radical initiator.
  • Hard PVC: Used in pipes, window frames.
  • Soft PVC: Used in flexible tubing, raincoats.
• Nylon Formation:
  • Condensation polymerisation of diamine and dicarboxylic acid.
  • Used in textile fibers, ropes, seat belts, parachutes.

22.6 Synthetic Adhesives

22.6.1 Describe types and applications of synthetic adhesives.

• Thermoplastic Adhesives:
  • Soften when heated and harden on cooling (e.g., hot glue).
• Thermosetting Adhesives:
  • Harden permanently after heating (e.g., epoxy resins, superglue).
• Applications: Used in woodwork, automobile industry, packaging, and electronics.

22.7 Pesticides

22.7.1 Define pesticides.

• Chemical substances used to eliminate or control pests that damage crops or spread diseases.

22.7.2 Discuss the types of pesticides on the basis of their uses in daily life.

• Insecticides: Kill insects (e.g., DDT, Malathion).
• Herbicides: Kill unwanted plants (e.g., Glyphosate).
• Fungicides: Control fungal infections (e.g., Mancozeb).
• Rodenticides: Kill rodents (e.g., Zinc Phosphide).

22.7.3 Discuss the advantages and disadvantages of using pesticides for the environment.

• Advantages:
  • Increase crop yield by protecting plants.
  • Prevent disease spread by controlling harmful organisms.
• Disadvantages:
  • Bioaccumulation in food chains, affecting wildlife.
  • Contaminate soil and water, leading to pollution.

23.1 Chemistry of Troposphere and Stratosphere

23.1.1 Describe the chemical reactions occurring in the atmosphere with reference to formation of acid rain, ozone, ammonium nitrates and sulphates, and carbon dioxide.

• Acid Rain Formation:
  • SO₂ and NO₂ react with water in the atmosphere forming H₂SO₄ and HNO₃.
  • Equations:
    • SO₂ + O₂ → SO₃
    • SO₃ + H₂O → H₂SO₄ (Sulfuric acid)
    • NO₂ + H₂O → HNO₃ (Nitric acid)
• Ozone Formation and Depletion:
  • O₂ + UV → 2O• (Oxygen radicals)
  • O• + O₂ → O₃ (Ozone)
  • Ozone depletion due to CFCs:
    • CFCl₃ → CFCl₂• + Cl• (Chlorine radical attacks ozone)
    • Cl• + O₃ → ClO• + O₂
    • ClO• + O• → Cl• + O₂
• Ammonium Nitrates and Sulphates:
  • NH₃ reacts with NO₃⁻ and SO₄²⁻ forming NH₄NO₃ and (NH₄)₂SO₄, contributing to air pollution.
• Carbon Dioxide:
  • Produced from combustion of fossil fuels and respiration.
  • CO₂ + H₂O ⇌ H₂CO₃ (Forms carbonic acid, affecting ocean pH).

23.1.2 Discuss the release of oxides of carbon, sulphur, nitrogen, and volatile organic compounds (VOCs) which are associated with combustion of hydrocarbon-based fuel.

• Oxides of Carbon (CO, CO₂): Released from burning of coal, oil, gas.
• Oxides of Sulfur (SO₂, SO₃): Formed by burning of sulfur-containing fuels.
• Oxides of Nitrogen (NO, NO₂): Emitted from vehicles, industries, and power plants.
• VOCs: Emitted from paints, solvents, fuels, contributing to smog formation.

23.1.3 Discuss problems associated with the release of pollutants, i.e., oxides of carbon, sulphur, nitrogen, VOCs, and peroxyacetyl nitrate (PAN).

• Health Issues: Respiratory diseases, lung cancer, asthma.
• Environmental Effects: Acid rain damages soil, water bodies, and structures.
• Smog Formation: NO₂ and VOCs form photochemical smog, reducing visibility.

23.1.4 Describe causes and impacts of oxidising and reducing smogs.

• Oxidising Smog (Photochemical Smog):
  • Forms due to NO₂ and VOCs reacting under sunlight.
  • Produces ozone (O₃) and PAN, causing eye irritation and breathing issues.
• Reducing Smog (Sulfurous Smog):
  • Forms in cold, humid conditions due to SO₂ and particulate matter.
  • Causes respiratory diseases and poor air quality.

23.1.5 Describe the role of chlorofluorocarbons (CFCs) in destroying ozone in the stratosphere.

• CFCs release chlorine radicals, which destroy ozone molecules.
• Consequence: Increased UV radiation reaching Earth, leading to skin cancer and crop damage.

23.1.6 List possible alternatives to the use of CFCs.

• Hydrofluorocarbons (HFCs) and Hydrochlorofluorocarbons (HCFCs).
• CO₂ and Propane-based refrigerants.

23.1.7 Explain climate change as a result of the greenhouse effect and global warming.

• Greenhouse Effect: Gases like CO₂, CH₄, and N₂O trap heat in the atmosphere.
• Global Warming Consequences:
  • Rising sea levels.
  • Melting glaciers.
  • Extreme weather events.

23.2 Water Pollution and Water Treatment

23.2.1 Describe the parameters of water analysis.

• pH Level: Indicates acidity or alkalinity.
• Dissolved Oxygen (DO): Essential for aquatic life.
• Biochemical Oxygen Demand (BOD): Measures organic pollution.
• Chemical Oxygen Demand (COD): Determines overall pollution.

23.2.2 Explain the methods of water purification, i.e., raw water treatment, sewage treatment, zeolite process, and reverse osmosis.

• Raw Water Treatment: Sedimentation, filtration, and chlorination.
• Sewage Treatment:
  • Primary Treatment: Removal of solid waste.
  • Secondary Treatment: Bacteria break down organic matter.
  • Tertiary Treatment: Removes nutrients and disinfects water.
• Zeolite Process: Removes hardness by exchanging Ca²⁺ and Mg²⁺ ions.
• Reverse Osmosis (RO): Uses a semi-permeable membrane to remove dissolved salts.

23.3 Green Chemistry

23.3.1 Describe the principles of green chemistry.

• Prevention of Waste: Minimize chemical waste.
• Atom Economy: Maximize reactant conversion into products.
• Use of Renewable Feedstocks: Utilize biodegradable materials.
• Energy Efficiency: Use reactions that consume less energy.
• Use of Catalysts: Reduce reaction time and side reactions.

23.3.2 Discuss the significance of green chemistry.

• Environmental Protection: Reduces pollution and waste.
• Sustainable Industrial Growth: Encourages eco-friendly production.
• Economic Benefits: Saves energy and raw materials.

24.1 Classical and Modern Methods of Analysis

24.1.1 Compare the classical and modern methods of structural analysis of compounds.

• Classical Methods:
  • Include qualitative and quantitative analysis through chemical reactions.
  • Examples: Titration, precipitation, gravimetric analysis.
• Modern Methods:
  • Use instrumental techniques for precise structural analysis.
  • Examples: Spectroscopy (IR, UV-Vis, Mass Spectrometry), Chromatography (HPLC, GC).

24.1.2 Describe the procedure of combustion analysis of hydrocarbons.

• Combustion analysis is used to determine carbon and hydrogen content.
• Steps:
  • Sample combustion in excess oxygen → forms CO₂ and H₂O.
  • Absorption of CO₂ and H₂O in specific compounds.
  • Mass difference determines C and H content.

24.1.3 Define the term ‘spectroscopy’.

• Spectroscopy is the study of the interaction of light with matter.
• Used to determine composition, structure, and functional groups of molecules.

24.1.4 Discuss applications of spectroscopy in analytical chemistry.

• Used for identification of unknown substances.
• Determines purity, functional groups, molecular structure.
• Applied in pharmaceuticals, forensic science, environmental analysis.

24.1.5 Explain the different regions of the electromagnetic spectrum (according to wavelength).

• Radio waves: Used in NMR spectroscopy.
• Microwaves: Used for rotational spectroscopy.
• Infrared (IR): Used for vibrational spectroscopy.
• Visible and Ultraviolet (UV-Vis): Used in electronic transitions.
• X-rays: Used in X-ray crystallography.

24.1.6 Explain atomic emission and atomic absorption spectrum.

• Atomic Emission Spectrum:
  • Formed when atoms absorb energy and emit light at specific wavelengths.
  • Example: Flame test colors of elements.
• Atomic Absorption Spectrum:
  • Formed when atoms absorb specific wavelengths of light.
  • Example: Used in AAS (Atomic Absorption Spectroscopy) for metal analysis.

24.1.7 Describe the basic principles of infrared (IR) spectroscopy (i.e., absorption of IR radiation, molecular rotation, molecular vibrations, vibrational coupling).

• IR Absorption: Molecules absorb IR radiation, causing vibrational transitions.
• Molecular Vibrations: Include stretching and bending of bonds.
• Vibrational Coupling: Interaction of multiple vibrational modes within a molecule.

24.1.8 Interpret the infrared (IR) spectra of benzene, acetone, acetic acid, and ethanol.

• Benzene: Shows C=C stretching (~1600 cm⁻¹).
• Acetone: Exhibits C=O stretching (~1700 cm⁻¹).
• Acetic Acid: Shows O-H stretch (~3300 cm⁻¹) and C=O stretch (~1700 cm⁻¹).
• Ethanol: Displays O-H stretch (~3300 cm⁻¹) and C-O stretch (~1000 cm⁻¹).

24.1.9 Predict whether a given molecule will absorb in the UV-Visible radiation.

• Compounds with conjugated double bonds absorb UV-Vis radiation.
• Example: Benzene absorbs in the UV region (~250 nm).
• Highly conjugated systems absorb in the visible region, leading to colored compounds.

24.1.10 Predict the colors of compounds (methylene blue and [Ti(H₂O)₆]³⁺) from their UV-Visible spectra.

• Methylene blue: Absorbs red light (~660 nm), appears blue.
• [Ti(H₂O)₆]³⁺: Absorbs yellow-green (~500 nm), appears violet.

24.1.11 Explain instrumentation and working of a mass spectrometer (MS).

• Ionization: Sample molecules are ionized into charged fragments.
• Acceleration: Ions are accelerated by an electric field.
• Deflection: Ions are separated based on mass-to-charge ratio (m/z).
• Detection: A detector records the mass spectrum.

24.1.12 Discuss the use of MS in determination of relative isotopic masses.

• Mass spectrometry helps determine isotopic composition of elements.
• Peak intensities indicate abundance of different isotopes.
• Example: Chlorine isotopes (³⁵Cl and ³⁷Cl) produce characteristic peaks.