Bond Properties and Valence Bond Theory

Bond Properties

Chemical bonds are the attractive forces holding atoms together in molecules. The properties of these bonds, such as bond length, bond energy, and bond order, are fundamental to understanding molecular structure and reactivity.

• Bond Length: The average distance between the nuclei of two bonded atoms. Shorter bonds are generally stronger and involve more shared electrons.

• Bond Energy (Bond Dissociation Energy): The energy required to break one mole of a bond in a molecule in the gas phase. Higher bond energy indicates a stronger bond.

• Bond Order: The number of chemical bonds between a pair of atoms. Higher bond order typically means a shorter and stronger bond.

• Polarity: Determined by the difference in electronegativity between bonded atoms. Polar bonds have unequal sharing of electrons.

Example: In a carbon-oxygen double bond (C=O), the bond order is 2, the bond is shorter and stronger than a single C–O bond, and the bond is polar due to the higher electronegativity of oxygen.

Valence Bond Theory

Valence Bond (VB) Theory explains the formation of chemical bonds as the overlap of atomic orbitals from different atoms. The strength of the bond depends on the extent of orbital overlap.

• Sigma (σ) Bonds: Formed by head-on overlap of orbitals (e.g., s-s, s-p, or p-p). All single bonds are sigma bonds.

• Pi (π) Bonds: Formed by side-to-side overlap of p orbitals. Present in double and triple bonds, in addition to a sigma bond.

• Hybridization: Atomic orbitals mix to form new, equivalent hybrid orbitals (e.g., sp, sp2, sp3).

Example: In methane (CH4), the carbon atom is sp3 hybridized, forming four equivalent sigma bonds with hydrogen atoms.

Bonding Theories and Molecular Geometry

Lewis Structures and Resonance

Lewis structures represent the arrangement of valence electrons in molecules. Resonance structures are alternative Lewis structures for the same molecule, showing delocalization of electrons.

• Octet Rule: Atoms tend to form bonds to achieve eight electrons in their valence shell (exceptions exist).

• Formal Charge: Used to determine the most stable Lewis structure. Calculated as:

• Resonance: Actual structure is a hybrid of all resonance forms.

Example: The carbonate ion (CO32–) has three resonance structures, each with a double bond to a different oxygen atom.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) Theory predicts the shapes of molecules based on electron pair repulsions around a central atom.

• Electron Domains: Regions of electron density (bonds or lone pairs) around a central atom.

• Geometry: Determined by minimizing repulsions between electron domains.

• Common Geometries:

  • 2 domains: Linear (180°)

  • 3 domains: Trigonal planar (120°)

  • 4 domains: Tetrahedral (109.5°)

  • 5 domains: Trigonal bipyramidal (90°, 120°)

  • 6 domains: Octahedral (90°)

Example: Ammonia (NH3) has four electron domains (three bonds, one lone pair), resulting in a trigonal pyramidal shape.

Hybridization and Molecular Orbital Theory

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• sp Hybridization: Linear geometry, 180° bond angles (e.g., BeCl2).

• sp2 Hybridization: Trigonal planar geometry, 120° bond angles (e.g., BF3).

• sp3 Hybridization: Tetrahedral geometry, 109.5° bond angles (e.g., CH4).

Example: Ethene (C2H4) has sp2 hybridized carbons, with a sigma bond and a pi bond between the carbons.

Molecular Orbital (MO) Theory

MO Theory describes the combination of atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.

• Bonding Orbitals: Lower energy, electrons in these orbitals stabilize the molecule.

• Antibonding Orbitals: Higher energy, electrons in these orbitals destabilize the molecule.

• Bond Order (MO Theory):

Example: In O2, the presence of unpaired electrons in antibonding orbitals explains its paramagnetism.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces (IMFs) are forces of attraction between molecules, influencing physical properties such as boiling and melting points.

• London Dispersion Forces: Weakest IMF, present in all molecules due to temporary dipoles.

• Dipole-Dipole Interactions: Occur between polar molecules with permanent dipoles.

• Hydrogen Bonding: Strongest IMF, occurs when H is bonded to N, O, or F.

Example: Water (H2O) exhibits hydrogen bonding, leading to its high boiling point relative to its molar mass.

Functional Groups and Organic Molecules

Common Functional Groups

Functional groups are specific groups of atoms within molecules that determine the chemical reactivity of those molecules.

• Alcohols: –OH group

• Aldehydes: –CHO group

• Ketones: C=O group within a carbon chain

• Carboxylic Acids: –COOH group

• Amines: –NH2 group

Example: Ethanol (CH3CH2OH) is an alcohol, while acetone (CH3COCH3) is a ketone.

Polymers and Macromolecules

Addition and Condensation Polymers

Polymers are large molecules made by linking repeating units (monomers). There are two main types of polymerization: addition and condensation.

• Addition Polymers: Formed by the addition of monomers with double bonds (e.g., polyethylene).

• Condensation Polymers: Formed by the reaction of monomers with the elimination of a small molecule (e.g., water), such as in nylon or polyester.

Example: Polyethylene is made from ethylene monomers, while nylon is made from diamine and dicarboxylic acid monomers.

Biological Macromolecules

DNA and Lipids

Biological macromolecules are essential for life and include nucleic acids (DNA, RNA) and lipids.

• DNA: Composed of nucleotides (sugar, phosphate, nitrogenous base). Stores genetic information.

• Lipids: Include fats, oils, and phospholipids. Important for energy storage and cell membrane structure.

Example: Triglycerides are formed from glycerol and three fatty acids; phospholipids are key components of cell membranes.

Summary Table: Types of Bonds and Forces