Overview
This lecture introduces the chapter on chemical bonding and molecular structure, covering core theories, types of bonds, Lewis structures, hybridization, molecular geometry, and special topics like resonance, bond parameters, and hydrogen bonding.
Introduction to Chemical Bonding
- Chemical bonding is the attractive force that holds atoms or ions together in molecules or compounds.
- The main focus is on how and why atoms form bonds and the theories that explain bonding.
Theories of Chemical Bonding
- Lewis approach: Only valence (outermost shell) electrons participate in bonding.
- Lewis dot structures represent only valence electrons using dots around element symbols.
- Octet rule: Atoms tend to have eight electrons in their valence shell for stability (exceptions: H, He).
- Atoms achieve stability by sharing (covalent) or transferring (ionic) electrons.
Types of Chemical Bonds
- Covalent bond: Sharing of electron pairs between two nonmetals; can form single, double, or triple bonds.
- Polar covalent bond: Shared electrons are unequally distributed due to differences in electronegativity.
- Ionic (electrovalent) bond: Complete transfer of electrons from a metal to a nonmetal, forming ions.
- Coordinate (dative) bond: Both electrons in the bond come from the same atom (e.g., in NH4⁺).
Lewis Structures, Bonding Rules, and Formal Charge
- Select the central atom (usually least electronegative and in lowest quantity).
- Draw skeletal structure with single bonds; complete octets by adding lone pairs or forming double/triple bonds.
- Calculate total valence electrons (Q), bond pair electrons, and distribute remaining as lone pairs.
- Formal charge = (Valence electrons) - (Lone pair electrons) - ½(Bond pair electrons).
- The most stable structure has the lowest formal charges.
Limitations of Lewis and Octet Rule
- Cannot explain molecules with incomplete or expanded octets or odd-electron species.
- Octet rule exceptions: Less than eight electrons (e.g., BCl3), more than eight (e.g., SF6), odd electrons (e.g., NO).
Bond Parameters
- Bond length: Equilibrium distance between nuclei of bonded atoms.
- Bond angle: Angle between two adjacent bonds at the central atom.
- Bond enthalpy: Energy required to break one mole of bonds in gaseous molecules.
- Bond order: Number of shared electron pairs (single = 1, double = 2, triple = 3).
- Bond strength is inversely related to bond length.
Resonance Structures
- Resonance occurs when more than one valid Lewis structure can be drawn.
- Actual structure is a resonance hybrid; charges and bonds are delocalized.
Bond Polarity and Dipole Moment
- Bond polarity occurs due to differences in electronegativity.
- Dipole moment measures the separation of charges and gives information about molecular polarity.
- Dµ = q × r (charge × separation).
Fajan’s Rule and Polarization
- Small, highly charged cations and large anions lead to greater polarization and more covalent character in otherwise ionic bonds.
Valence Bond Theory and Types of Overlapping
- Bonds result from overlapping atomic orbitals: head-on (sigma bond, strong) or sideways (pi bond, weak).
- Sigma bonds: Head-on overlap (s-s, s-p, p-p).
- Pi bonds: Side-wise (parallel) p-p overlap.
Hybridization and Molecular Shapes
- Hybridization: Mixing of atomic orbitals to form equivalent hybrid orbitals (e.g., sp, sp2, sp3).
- Geometry determined by the number of sigma bonds and lone pairs (e.g., sp – linear, sp2 – trigonal planar, sp3 – tetrahedral).
- Use Z = number of sigma bonds + lone pairs to determine hybridization.
VSEPR Theory (Valence Shell Electron Pair Repulsion)
- Molecular shape determined by repulsions among electron pairs (lone pair-lone pair > lone pair-bond pair > bond pair-bond pair).
- Shape and geometry may differ if there are lone pairs on the central atom.
Molecular Orbital Theory
- Atomic orbitals combine to form bonding and antibonding molecular orbitals.
- Electron filling follows Hund’s rule and Pauli exclusion principle.
- Bond order = ½(Number of bonding electrons – Number of antibonding electrons).
- Magnetic properties: Paired electrons (diamagnetic), unpaired (paramagnetic).
Hydrogen Bonding
- Hydrogen bonding occurs when H is bonded to highly electronegative atoms (N, O, F).
- Two types: intermolecular (between molecules), intramolecular (within a molecule).
- Stronger than van der Waals forces but weaker than ionic/covalent bonds.
Key Terms & Definitions
- Valence electrons — electrons in the outermost shell of an atom, important in bonding.
- Octet rule — tendency of atoms to attain eight electrons in their valence shell.
- Lewis structure — diagram showing bonding and lone pairs using dots.
- Covalent bond — bond formed by sharing electron pairs between atoms.
- Ionic bond — bond formed by complete transfer of electrons, resulting in ions.
- Bond order — number of chemical bonds between a pair of atoms.
- Resonance — delocalization of electrons across multiple structures.
- Hybridization — mixing of atomic orbitals to form new hybrid orbitals.
- VSEPR theory — model to predict molecular shapes based on electron pair repulsion.
- Formal charge — hypothetical charge assigned to atoms in a molecule.
Action Items / Next Steps
- Practice drawing Lewis dot structures and calculating formal charges for given molecules.
- Memorize typical geometries and hybridization types (sp, sp2, sp3, etc.).
- Complete homework: Draw and analyze structures for NO3⁻ and ethanoic acid; comment your answers.
- Review the VSEPR chart for predicting molecular shapes.
- Prepare for next lecture by revising molecular orbital diagrams and bond order calculations.