Bond Properties and Valence Bond Theory

Introduction to Bond Properties

Chemical bonds are the attractive forces that hold atoms together in molecules and compounds. Understanding bond properties is essential for predicting molecular structure and reactivity.

• Bond length: The average distance between the nuclei of two bonded atoms. Shorter bonds are generally stronger.

• Bond strength (bond energy): The energy required to break a bond in a molecule. Stronger bonds have higher bond energies.

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

• Bond polarity: A measure of how equally electrons are shared between two atoms. Determined by the difference in electronegativity.

Valence Bond Theory explains bonding as the overlap of atomic orbitals, forming sigma (σ) and pi (π) bonds.

• Sigma (σ) bonds: Formed by head-on overlap of orbitals; all single bonds are sigma bonds.

• Pi (π) bonds: Formed by side-on overlap of p orbitals; present in double and triple bonds.

Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals (e.g., sp, sp2, sp3).

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

Resonance and Delocalization

Resonance Structures

Some molecules cannot be represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for the same molecule, differing only in the placement of electrons.

• Actual structure is a resonance hybrid, with delocalized electrons.

• Resonance increases stability by spreading out electron density.

Example: The nitrate ion (NO3-) has three resonance structures, each with a different N–O double bond.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the shapes of molecules based on the repulsion between electron pairs around a central atom.

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

• Geometry: Electron domains arrange themselves to minimize repulsion, determining molecular shape.

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: Water (H2O) has four electron domains (two bonds, two lone pairs), resulting in a bent geometry.

Bond Polarity and Molecular Polarity

Electronegativity and Dipole Moments

Electronegativity is the ability of an atom to attract electrons in a bond. A difference in electronegativity creates a polar bond with a dipole moment.

• Dipole moment (μ): A measure of bond polarity, calculated as (charge × distance).

• Molecular polarity: Determined by both bond polarity and molecular geometry.

Example: Carbon dioxide (CO2) has polar bonds but is nonpolar overall due to its linear geometry.

Hybridization and Molecular Orbital Theory

Hybridization

Hybridization explains the observed shapes of molecules by combining atomic orbitals into hybrid orbitals.

• sp: Linear geometry (e.g., BeCl2)

• sp2: Trigonal planar geometry (e.g., BF3)

• sp3: Tetrahedral geometry (e.g., CH4)

Molecular Orbital (MO) Theory describes bonding as the combination of atomic orbitals to form molecular orbitals, which can be bonding or antibonding.

• Bonding orbitals: Lower energy, increase electron density between nuclei.

• Antibonding orbitals: Higher energy, decrease electron density between nuclei.

Example: In O2, MO theory explains its paramagnetism due to unpaired electrons in π* orbitals.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules, influencing physical properties like boiling and melting points.

• London dispersion forces: Weak, present in all molecules due to temporary dipoles.

• Dipole-dipole interactions: Occur between polar molecules.

• Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

Example: Water has strong hydrogen bonds, resulting in high boiling point.

Functional Groups and Organic Molecules

Common Functional Groups

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

• 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.

Polymers and Macromolecules

Addition and Condensation Polymers

Polymers are large molecules made by joining many small units (monomers).

• Addition polymers: Formed by addition reactions (e.g., polyethylene from ethylene).

• Condensation polymers: Formed by condensation reactions, releasing small molecules like water (e.g., nylon, proteins).

Example: Proteins are condensation polymers made from amino acids.

Biological Macromolecules

DNA and Lipids

DNA is a nucleic acid that stores genetic information. It consists of nucleotides with a sugar, phosphate, and nitrogenous base.

• Double helix structure, complementary base pairing (A–T, G–C).

Lipids are hydrophobic molecules, including fats, oils, and phospholipids. They are important for energy storage and cell membranes.

• Fats are esters of glycerol and fatty acids.

• Phospholipids form bilayers in cell membranes.

Additional info: Where diagrams of molecular shapes, orbital overlaps, or polymer structures were referenced in the notes, these have been described in text for clarity. No images were included, as none were directly relevant or sufficiently clear to reinforce the explanations above.