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Electronegativity Difference

• As the electronegativity difference (ΔEN) increases between two atoms, the bond becomes more ionic.

Intramolecular Forces

• Intramolecular forces refer to the chemical bonds holding the atoms together in molecules.

• The three major types of chemical bonds:

  • Metallic bond

  • Ionic bond

  • Covalent bond

Metallic Bond

• Represents the strongest type of chemical bond.

• Over 80 elements out of 118 in the periodic table are metals.

• Metals are characterized as hard solids with closely packed atoms.

• Metallic bond defined as the force holding atoms closely together, consisting of positively charged ions fixed in a crystal lattice surrounded by delocalized electrons.

Ionic Bond

• An ionic bond occurs when the electronegativity difference between bonded atoms is significant (typically > 1.9), leading to the complete transfer of bonding electrons.

• This results in the formation of cations and anions.

• Electrostatic interactions occur between the oppositely charged ions, orienting in a 3D crystal lattice to maximize attractive interactions and minimize repulsive ones.

• Ionic bonds are usually weaker than metallic bonds but stronger than other types of bonds.

Covalent Bond

• A covalent bond occurs when the electronegativity difference between bonded atoms is moderate to zero (typically < 1.9).

• Bonding electrons are shared, with attractive forces formed between the bonding electrons and the atomic nuclei.

• Generally weaker than metallic and ionic bonds but stronger than intermolecular forces.

Intermolecular Forces

• Intermolecular forces are electrostatic interactions between molecules and are typically much weaker than intramolecular forces.

• Major types of intermolecular forces include:

  • Dipole-dipole interactions

  • Hydrogen bonding

  • London dispersion forces

Dipole-Dipole Interactions

• Polar molecules exhibit permanent dipoles, with one end being partial positive (δ+) and the other partial negative (δ-).

• These polar molecules interact through their δ+ and δ- ends in what are known as dipole-dipole interactions, which are weaker than ionic bonds.

• Molecules orient themselves to maximize attractive forces between opposite charges while minimizing repulsion between like charges.

Hydrogen Bonds

• A specialized type of dipole-dipole interaction involving hydrogen bound to highly electronegative atoms (O, N, F).

• Hydrogen bonding is stronger than standard dipole-dipole interactions and plays vital roles in biomolecules like DNA and proteins.

• Key features include:

  • Higher electronegativity differences with O, N, or F than other polar bonds.

  • Greater charge density on hydrogen due to its smaller size, enabling closer interaction with δ- atoms.

London Dispersion Forces

• Present in all molecules, including nonpolar ones, creating temporary charge imbalances in electron clouds.

• These temporary dipoles can induce dipoles in neighboring molecules, resulting in transient dipole-induced dipole interactions.

Hydrocarbons

• Organic compounds consisting exclusively of carbon and hydrogen.

• Hydrocarbons can be classified based on carbon-carbon bond types into:

  • Saturated hydrocarbons (only single bonds)

  • Unsaturated hydrocarbons (double or triple bonds present)

  • Aromatic hydrocarbons

• Saturated Hydrocarbons (Alkanes): Contain only single C-C bonds and can be straight-chain or cyclic (cycloalkanes).

Functional Groups with Carbon Bonded to Electronegative Atoms

• Alkyl halides (C-X)

• Alcohols (C-OH)

• Ethers (C-O-C)

• Amines (C-N)

• Thiols (C-SH)

• Sulfides (C-S-C)

Groups with Carbon-Oxygen Double Bonds

• Aldehyde (C=O with H)

• Ketone (C=O with two Cs)

• Carboxylic acid (–OH bonded to C=O)

• Ester (–O bonded to C=O)

• Amide (C-N bonded to C=O)

Alkanes and Alkane Isomers

• Alkanes are hydrocarbons with C-C single bonds and the formula CnH2n+2.

• They are saturated and called aliphatic compounds.

• Isomerism occurs in alkanes with four or more carbons, leading to different structural formulas despite the same molecular formula.

Physical Properties of Alkanes

• Alkanes are non-polar and generally insoluble in polar solvents but soluble in non-polar solvents.

• Alkanes' physical states: gaseous (C1-C4), liquid (C5-C17), and solid (C18+).

• Boiling points depend on isomerism; branched alkanes typically have lower boiling points than their straight-chain counterparts.

• Alkanes are lighter than water, contributing to their buoyancy.

Reactions of Alkanes

• Alkanes, being relatively inert, mainly undergo substitution reactions, often involving halogenation.

• Free Radical Chlorination of Methane:

  1. Initiation: Cl2 + Light → 2 Cl.

  2. Propagation: Cl + CH4 → .CH3 + HCl; .CH3 + Cl2 → CH3-Cl + .Cl.

  3. Termination: .CH3 + .Cl → CH3-Cl.

Combustion of Hydrocarbons

• Complete combustion produces carbon dioxide and water, whereas incomplete combustion may produce carbon monoxide alongside carbon dioxide and water.

Grignard Reaction

• Involves reacting an organic halide (R-X) with magnesium in dry ether to form a Grignard reagent (R-MgX) which acts as a nucleophile in further reactions.

Unsaturated Hydrocarbons

• Hydrocarbons containing at least one C=C or C≡C bond, classified as alkenes and alkynes.

Addition Reactions of Alkenes

• Reactions predominantly occur when a double bond is converted to single bonds, leading to more stable products.

• Markovnikov's Rule: Describes the regioselectivity in reactions involving asymmetrical alkenes.

Polymerization of Alkenes

• Alkenes can undergo self-addition reactions leading to polymers, marked by long-chain molecules from repeating units.

Methods of Preparation of Alkenes

1. Reduction of alkynes.

2. Dehydrohalogenation of alkyl halides.

3. Dehydration of alcohols using concentrated sulfuric acid.

Alkynes

• Unsaturated hydrocarbons with at least one C≡C bond, showing chain, position, and functional isomerism.

• Ethyne is an important compound with industrial applications, notably in welding.