3.2 Lewis structures and molecular geometry

Lewis structures and molecular geometry are crucial concepts in understanding chemical bonding. They help us visualize how atoms connect and arrange themselves in molecules. This knowledge is key to predicting molecular properties and reactivity.

VSEPR theory and hybridization explain why molecules have specific shapes. These ideas build on Lewis structures, showing how electron arrangement influences molecular geometry. Understanding these concepts is essential for grasping more complex chemical interactions.

Lewis Structures and the Octet Rule

Drawing Lewis Structures

• Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule
• To draw a Lewis structure:
  1. Determine the total number of valence electrons in the molecule or ion by adding the valence electrons of each atom and accounting for any overall charge
  2. Place the least electronegative atom in the center and connect it to the other atoms with single bonds
  3. Distribute the remaining electrons as lone pairs to complete the octets of the outer atoms, and then place any leftover electrons on the central atom
  4. If there are not enough electrons to give the outer atoms an octet, try multiple bonds (double or triple) between the central atom and outer atoms
• The octet rule states that atoms tend to have eight electrons in their valence shell, except for hydrogen, which tends to have two

Exceptions to the Octet Rule

• In some cases, the octet rule is not followed:
  • Molecules with an odd number of electrons (NO, ClO2)
  • Molecules with less than an octet (BF3, BeCl2)
  • Molecules with more than an octet (SF6, PCl5)
• These exceptions occur due to the presence of incomplete octets, expanded octets, or the inability to achieve an octet configuration

Formal Charge in Molecules and Ions

Calculating Formal Charge

• Formal charge is the charge assigned to an atom in a molecule or ion, assuming that electrons in a chemical bond are shared equally between atoms
• To calculate formal charge, use the formula:

$\text{Formal charge} = [\text{Number of valence electrons in free atom}] - [\text{Number of non-bonding electrons}] - \frac{1}{2}[\text{Number of bonding electrons}]$

Formal Charge and Molecular Stability

• The sum of the formal charges on all atoms in a molecule or ion should equal the overall charge of the species
• Atoms in molecules and ions should have formal charges as close to zero as possible
• When drawing Lewis structures, if there are multiple valid structures, the one with the formal charges closest to zero is usually the most stable and preferred (CO2, SO3)

Molecular Geometry with VSEPR Theory

Predicting Molecular Geometry

• VSEPR theory predicts the geometry of a molecule based on the number of electron pairs (bonding and non-bonding) around the central atom
• Electron pairs repel each other and arrange themselves to minimize repulsion, leading to specific geometric arrangements
• The five basic geometric arrangements are:
  • Linear (2 electron pairs)
  • Trigonal planar (3 electron pairs)
  • Tetrahedral (4 electron pairs)
  • Trigonal bipyramidal (5 electron pairs)
  • Octahedral (6 electron pairs)

Effect of Lone Pairs on Molecular Geometry

• Lone pairs occupy more space than bonding pairs, so they have a greater influence on the geometry
• The presence of lone pairs can distort the basic geometries, resulting in:
  • Bent or angular (H2O, SO2)
  • Trigonal pyramidal (NH3, PCl3)
  • Seesaw (SF4)
  • T-shaped (ClF3)
  • Square pyramidal (BrF5)

Geometry and Bond Angles

Ideal Bond Angles

• Bond angles are the angles formed between the imaginary lines connecting the nuclei of atoms in a molecule
• Molecular geometry determines the ideal bond angles in a molecule based on the arrangement of electron pairs around the central atom
• In the five basic geometric arrangements, the ideal bond angles are:
  • Linear: 180°
  • Trigonal planar: 120°
  • Tetrahedral: 109.5°
  • Trigonal bipyramidal: 90° and 120°
  • Octahedral: 90°

Effect of Lone Pairs on Bond Angles

• The presence of lone pairs can distort these ideal bond angles due to their greater repulsive effect
• Lone pairs push the bonding pairs closer together, resulting in slightly smaller bond angles than the ideal angles in the basic geometries (H2O: 104.5°, NH3: 107°)

Hybridization of Atomic Orbitals

Types of Hybridization

• Hybridization is the mixing of atomic orbitals to form new hybrid orbitals with specific geometries and energies
• Hybrid orbitals are formed to minimize electron repulsion and create a more stable bonding arrangement
• The common types of hybridization are:
  • sp hybridization: 2 electron pairs, linear geometry
  • sp2 hybridization: 3 electron pairs, trigonal planar geometry
  • sp3 hybridization: 4 electron pairs, tetrahedral geometry
  • sp3d hybridization: 5 electron pairs, trigonal bipyramidal geometry
  • sp3d2 hybridization: 6 electron pairs, octahedral geometry

Determining Hybridization

• To determine the hybridization of the central atom in a molecule or ion:
  1. Predict the molecular geometry using VSEPR theory
  2. Match the geometry with the corresponding hybridization type
• Examples:
  • BeCl2: Linear geometry, sp hybridization
  • BF3: Trigonal planar geometry, sp2 hybridization
  • CH4: Tetrahedral geometry, sp3 hybridization
  • PCl5: Trigonal bipyramidal geometry, sp3d hybridization
  • SF6: Octahedral geometry, sp3d2 hybridization