Unit 12

1. Why Carbon?

• Carbon has 4 valence electrons → forms 4 covalent bonds (tetravalent).

• It achieves stability by sharing electrons, not gaining or losing them (ions like C⁴⁺ or C⁴⁻ are too unstable).

• Catenation: Carbon atoms bond to each other forming chains and rings.

• Can form single, double, or triple bonds with other carbon or heteroatoms (O, N, etc.).

• C–C and C–H bonds are strong, short, and nearly non-polar (similar electronegativities).

• This makes carbon chemistry extremely versatile — over 90% of known compounds are organic.

• Organic chemistry includes both natural (proteins, lipids, carbohydrates, nucleic acids) and synthetic compounds (plastics, dyes, solvents, explosives, drugs).

2. Organic vs Inorganic Compounds

1. Organic compounds: Contain carbon, usually bonded with hydrogen, oxygen, or nitrogen.

Properties:

• Low melting and boiling point 

• Solubility varies

• Flammable, react with O2 producing H2O and CO2

• Joined by covalent bonds

2. Inorganic compounds: Include carbonates, oxides of carbon, and elemental carbon (graphite, diamond).

Properties:

• High melting and boiling points

• Solid at room temperature

• Non-combustile

• Good conductors

• Dissolve in water to form ions

3. Homologous Series

• A homologous series is a family of organic compounds with:

  • The same functional group,

  • Similar chemical properties,

  • Gradually changing physical properties (e.g., boiling point increases with chain length),

  • A general formula (e.g., alkanes: CₙH₂ₙ₊₂).

• Each successive member differs by a –CH₂– unit.

• Compounds are named according to IUPAC rules, based on their homologous series and number of carbons.

4. Alkanes

• General formula: CₙH₂ₙ₊₂

• Saturated hydrocarbons — only single C–C bonds.

• Unreactive, except for combustion.

• Non-polar

• Colorless and odorless

• Insoluble in water but soluble in organic solvents

• Low boiling points 

• Combustion reaction:

  • Complete: Alkane + O₂ → CO₂ + H₂O + energy

  • Incomplete: Alkane + limited O₂ → CO + C + H₂O

• Uses: Fuels (methane, propane, butane, petrol, diesel, kerosene).

• Source: Crude oil (petroleum).

• Natural examples:

  • Pristane (C₁₉H₄₀) in shark liver oil.

  • Methane produced by microbes in ruminants.

  • Long-chain alkanes in plant waxes (cuticle) protect leaves against water loss, UV, and microbes.

5. Alkenes

• General formula: CₙH₂ₙ

• Unsaturated hydrocarbons — contain a C=C double bond.

• Names end in –ene (e.g., ethene, propene).

• More reactive than alkanes due to the double bond.

• Combustion: Burn in air to form CO₂ and H₂O but are usually used for synthesis, not fuel.

• Colorless 

Addition Reactions

• A molecule adds across the double bond, forming one product.

• Example: Bromine + Ethene → 1,2-Dibromoethane

• Used to test for unsaturation (bromine water decolorizes).

Hydrogenation

• Ethene + H₂ → Ethane

• Requires a Ni catalyst and heat.

• Converts unsaturated oils into saturated fats.

Hydration

• Ethene + H₂O (steam) → Ethanol

• Catalyst: Hot concentrated H₃PO₄ (phosphoric acid).

6. Functional Groups

A functional group determines the chemical properties of a molecule.

Alcohols

• General formula: CₙH₂ₙ₊₁OH

• Polar due to –OH group → soluble in water, react with sodium.

• Example: Methanol, Ethanol.

• Produced by hydration of alkenes.

• Increased water solubility(ability to form hydrogen bonds with water) 

• Higher boiling points 

• Used as solvents, fuels, and intermediates in organic synthesis.

Carboxylic Acids

• General formula: CₙH₂ₙ₊₁COOH

• Weak acids (release H⁺ in water).

• Polar at the –COOH end.

• Solubility decreases with longer hydrocarbon chains.

• High boiling points 

• Example: Methanoic acid, Ethanoic acid (vinegar).

Esters

• General formula: CₙH₂ₙ₊₁COOCₘH₂ₘ₊₁

• Functional group: –COO–

• Formed by reaction between a carboxylic acid and an alcohol (esterification)

• Condensation reaction → produces water as a byproduct

• Contain a carbonyl (C=O) next to an oxygen–carbon bond

• Have a sweet, fruity smell (used in flavorings and perfumes)

• Non-polar- low solubility in water

• Lower boiling points than carboxylic acids (no hydrogen bonding between molecules)

• Volatile(easily evaporated) liquids

• Examples: Ethyl ethanoate, methyl propanoate

7. Isomerism

• Isomers: Same molecular formula, different structural arrangement.

Types:

1. Structural isomers: Different bonding pattern (e.g., butane vs. methylpropane).

2. Geometric (cis/trans) isomers: Occur in alkenes due to restricted rotation around C=C bond.

   • Cis: groups on same side of double bond.

   • Trans: groups on opposite sides.

3. Enantiomers (optical isomers): Non-superimposable mirror images with an asymmetric carbon.

   • Rotate plane-polarized light in opposite directions:

     • d / + = clockwise, l / – = anticlockwise.

   • Have identical physical properties except optical activity.

8. Crude Oil and Its Refining

• Crude oil: Complex mixture of hydrocarbons (mainly alkanes).

• Found under impermeable rock layers, extracted by drilling.

• Major source of fuels and chemical feedstocks.

Refining Process:

1. Fractional Distillation:

   • Separates hydrocarbons by boiling point.

   • Lighter molecules rise to top (gasoline); heavier stay at bottom (bitumen).

   • Main fractions: gases, petrol, kerosene, diesel, lubricating oil, residue.

2. Cracking:

   • Large hydrocarbon molecules are broken into smaller, more useful ones.

   • Produces alkenes (like ethene) for chemical synthesis.

   • Methods:

     • Thermal cracking (high temperature).

     • Catalytic cracking (lower temp, solid catalyst).

   • Ethene decolorizes bromine — test for unsaturation.

3. Reforming:

   • Converts straight-chain alkanes to branched or aromatic hydrocarbons (benzene, toluene).

   • Improves fuel efficiency and octane rating.

   • Uses hydrogen gas and a catalyst.

   • Aromatics (like benzene) burn with sooty yellow flames and are toxic.

9. Organic Materials

Soaps and Detergents

• Soaps: Made from natural fats/oils via saponification (reaction with NaOH).

• Detergents: Synthetic (from petroleum), more soluble, don’t form scum in hard water.

• Both have:

  • Hydrophobic tail (non-polar, attracted to grease).

  • Hydrophilic head (polar, attracted to water).

• Function as emulsifiers — help grease mix with water to be rinsed away.

10. Polymers

• Polymers: Large molecules (macromolecules) formed from many monomers.

• Types:

  1. Addition Polymerization – monomers add without losing atoms.

     • Monomer must be unsaturated (C=C).

     • Example: Ethene → Poly(ethene) (polythene).

     • Other examples: PVC, polystyrene, PTFE.

  2. Condensation Polymerization – monomers join with loss of small molecules (usually H₂O).

     • Example:

       • Nylon (polyamide): from amine + carboxylic acid.

       • Terylene (polyester): from diol + dicarboxylic acid.

     • Occurs in proteins (amino acids) and DNA (nucleotides) in living organisms.

Hydrolysis

• Reverse of condensation polymerization.

• Breaks down polymers into monomers using water.

• Example: Protein → amino acids (during digestion).

11.  Electronegativity difference

Intermolecular forces 

• London forces: Weak, temporary intermolecular attractions that occur between all atoms and molecules due to the constant motion of electrons

• Dipole-Dipole forces: A type of attraction between partially polar molecules

• Hydrogen bonds:  Strong intermolecular attraction between Hydrogen and highly electronegative atom like oxygen, nitrogen, and fluorine. 

12. Emulsifiers

• Substances that help mix 2 immiscible liquids by preventing separation

• Have two parts:

  • Hydrophilic (water-loving) head

  • Hydrophobic (water-hating) tail

• Form stable emulsions