Chemical Bonding and Structure

Overview

This lecture introduces the chapter "Chemical Bonding and Molecular Structure," covering basic bonding concepts, major bonding theories, types of chemical bonds, hybridization, molecular geometry, and related properties.

Introduction to Chemical Bonding

• Chemical bonding is the attractive force that holds atoms or ions together in compounds.
• Chemical bonds form due to the force of attraction between atoms, ions, or molecules.
• The main types of bonds are ionic, covalent, and coordinate bonds.

Theories of Chemical Bonding

• Lewis Approach: Bonding involves sharing or transfer of valence shell electrons.
• Octet Rule: Atoms achieve stability when their outer shell has 8 electrons (or 2 for H and He).
• Exceptions to the octet rule exist (e.g., H, He, B, expanded octets for P, S, etc.).

Lewis Structures & Formal Charge

• Lewis symbols represent valence electrons as dots around element symbols.
• Lewis dot structures depict how atoms share or transfer electrons to complete the octet/duplet.
• Formal charge = valence electrons – (lone pair electrons + ½ shared electrons); the lowest formal charge gives the most stable structure.

Types of Chemical Bonds

• Ionic (Electrovalent) bond: Transfer of electrons between metals and non-metals, forming cations and anions.
• Covalent bond: Sharing of electron pairs between two non-metals.
• Coordinate (Dative) bond: Both shared electrons come from the same atom (lone pair donation).
• Polar and non-polar covalent bonds depend on electronegativity differences.

Bond Parameters & Properties

• Bond length: Equilibrium distance between nuclei of bonded atoms.
• Bond angle: Angle between two bonds at the central atom.
• Bond enthalpy: Energy needed to break one mole of bonds in gaseous state.
• Bond order: Number of bonds between two atoms (single, double, triple), higher bond order means greater strength.
• Dipole moment: Measure of bond polarity; higher with greater charge separation and bond length.

Resonance & Limitations

• Resonance: When one Lewis structure fails to describe a molecule, multiple canonical structures are used.
• Real structure is a resonance hybrid, not any single structure.
• Octet rule has many exceptions; resonance helps explain molecular properties.

Hybridization & Molecular Shapes

• Hybridization: Mixing of atomic orbitals to form new, equivalent hybrid orbitals (sp, sp², sp³, sp³d, sp³d², etc.).
• The number and type of hybrid orbitals determine molecular geometry (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, etc.).
• VSEPR theory: Molecular shape depends on repulsions among all electron pairs (bonding + lone pairs) around the central atom.
• Lone pairs exert more repulsion than bonding pairs; repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.

Valence Bond Theory & Molecular Orbital Theory

• Valence Bond Theory: Bonds form by overlap of half-filled atomic orbitals; explains sigma and pi bonds.
• Molecular Orbital Theory: Atomic orbitals combine to form bonding and anti-bonding molecular orbitals.
• Bond order = (number of bonding electrons – number of antibonding electrons)/2.
• Magnetic properties: Unpaired electrons yield paramagnetism; all pairs yield diamagnetism.

Hydrogen Bonding & Fajans’ Rule

• Hydrogen bonding occurs between H and highly electronegative atoms (F, O, N); can be intermolecular or intramolecular.
• Fajans’ rule predicts polarization—small, highly charged cations, and large anions lead to covalent character in ionic bonds.

Key Terms & Definitions

• Valence Electron — electron in the atom’s outer shell involved in bonding.
• Octet Rule — atoms gain, lose, or share electrons to attain 8 electrons in their valence shell.
• Formal Charge — apparent charge on an atom in a structure, calculated to determine stability.
• Hybridization — mixing of atomic orbitals to form new hybrid orbitals for bonding.
• Resonance Structure — different Lewis structures for a molecule that cannot be represented by one structure.
• Bond Order — number of bonds between two atoms.
• Dipole Moment — measure of separation of charge in a molecule.
• VSEPR Theory — explains molecular shape based on electron pair repulsions.
• Sigma Bond (σ) — bond formed by head-on orbital overlap.
• Pi Bond (π) — bond formed by sideways orbital overlap.

Action Items / Next Steps

• Practice drawing Lewis structures and calculating formal charges for given molecules.
• Complete homework: Comment the number of sigma and pi bonds in ethanoic acid; determine the shape for the example with three lone pairs on six groups.
• Review VSEPR chart for predicting molecular geometry.
• Prepare for questions on the application of bond theories and resonance.