Chemical Bonding and Molecular Structure

Learning Objectives

• Understand the Kössel-Lewis approach to chemical bonding, explaining the formation of chemical bonds through electron transfer or sharing to achieve stable octet configurations.
• Explain the octet rule and its limitations, including exceptions such as incomplete octets, odd-electron molecules, and expanded octets.
• Calculate and interpret bond order in covalent bonds, understanding its effect on bond strength and length.
• Analyze resonance structures to understand the concept of resonance and its role in representing molecules as hybrids of multiple Lewis structures.
• Describe the formation of ionic or electrovalent bonds through electron transfer and the stabilization of lattice structures by lattice enthalpy.
• Calculate lattice enthalpy and understand its significance in the stability of ionic compounds.
• Apply VSEPR theory to predict molecular shapes based on electron pair repulsions, considering both bonding and lone pairs around the central atom.
• Use valence bond theory to explain chemical bonding through the overlap of atomic orbitals, considering hybridization and electron pair sharing.
• Quantify molecular polarity using dipole moments and understand how molecular geometry influences polarity.
• Draw Lewis structures and calculate formal charges for molecules and ions, ensuring accurate representation of valence electrons and electron pair distribution.
• Analyze bond parameters such as bond length, bond angle, bond enthalpy, and their relation to bond order, strength, and length.
• Differentiate between sigma and pi bonds, understanding their formation through head-on and sidewise overlap, and why sigma bonds are generally stronger than pi bonds.
• Describe hybridization and molecular shapes for various hybridizations such as sp, sp², sp³, sp³d, and sp³d², using examples like $BeCl_2$, $BF_3$, $CH_4$, $NH_3$, $H_2O$, $PCl_5$, and $SF_6$.
• Explain hydrogen bonding, both intermolecular and intramolecular, involving hydrogen with fluorine, oxygen, or nitrogen, and its effect on the structure and properties of compounds.

Chapter Notes

Kössel-Lewis Approach to Chemical Bonding

• Kössel and Lewis (1916): First to provide a logical explanation of valence based on noble gas inertness.
• Lewis Symbols: Represent valence electrons as dots around the element symbol.
• Octet Rule: Atoms aim to achieve a stable octet configuration by electron transfer or sharing.
• Examples:
  • $Na + Cl \rightarrow Na^+ + Cl^- \rightarrow NaCl$
  • $Ca + F_2 \rightarrow Ca^{2+} + 2F^- \rightarrow CaF_2$

Octet Rule and Its Limitations

• Octet Rule: Atoms combine to have eight electrons in their valence shell.
• Limitations:
  • Incomplete octets (e.g., $BeH_2$, $BF_3$).
  • Odd-electron molecules (e.g., $NO$, $NO_2$).
  • Expanded octets (e.g., $PF_5$, $SF_6$).

Covalent Bond and Bond Order

• Covalent Bond: Formed by sharing of electron pairs.
• Bond Order: Number of shared electron pairs.
  • Single bond: Bond order 1
  • Double bond: Bond order 2
  • Triple bond: Bond order 3

Resonance Structures

• Resonance: Multiple Lewis structures represent a molecule.
• Example: Ozone ($O_3$) has resonance structures with bond lengths intermediate between single and double bonds.

Ionic or Electrovalent Bond

• Formation: Electron transfer from a metal to a non-metal.
• Example: $Na^+ + Cl^- \rightarrow NaCl$

Lattice Enthalpy

• Definition: Energy required to separate one mole of an ionic solid into gaseous ions.
• Example: Lattice enthalpy of $NaCl$ is 788 kJ/mol.

VSEPR Theory

• Predicts Molecular Shapes: Based on electron pair repulsions.
• Geometries:
  • Linear: 180°
  • Trigonal Planar: 120°
  • Tetrahedral: 109.5°
  • Trigonal Bipyramidal: 90°, 120°
  • Octahedral: 90°

Valence Bond Theory

• Overlap of Atomic Orbitals: Explains bond formation.
• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.

Dipole Moment and Molecular Polarity

• Dipole Moment ($\mu$): Product of charge and distance between charges.
• Polarity: Depends on molecular geometry and bond dipoles.

Lewis Structures and Formal Charge

• Lewis Structures: Show valence electrons as dots.
• Formal Charge: Difference between valence electrons in free atom and electrons assigned in structure.

Bond Parameters

• Bond Length: Distance between nuclei of bonded atoms.
• Bond Angle: Angle between two bonds at an atom.
• Bond Enthalpy: Energy required to break a bond.

Sigma and Pi Bonds

• Sigma Bonds: Formed by head-on overlap of orbitals.
• Pi Bonds: Formed by sidewise overlap.

Hybridisation and Molecular Shapes

• Types: $sp$, $sp^2$, $sp^3$, $sp^3d$, $sp^3d^2$
• Examples:
  • $sp$: $BeCl_2$
  • $sp^2$: $BF_3$
  • $sp^3$: $CH_4$
  • $sp^3d$: $PCl_5$
  • $sp^3d^2$: $SF_6$

Hydrogen Bonding

• Involves H with F, O, or N: Affects structure and properties.
• Types: Intermolecular and intramolecular hydrogen bonding.