Chemical Bonding: Lewis Structures, Molecular Shapes, and Bond Theories

Chapter 5: Chemical Bonding I – Lewis Structures and Molecular Shapes

Bond Polarity and Electronegativity

Bond polarity arises from differences in electronegativity between atoms in a bond. Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond.

• Electronegativity Trend: Increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Bond Polarity: A bond is polar if the electronegativity difference between the two atoms is significant (typically > 0.4).

• Example: The H–Cl bond is polar because Cl is much more electronegative than H.

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms (central atom usually least electronegative).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons to complete octets (or duets for H).

  5. Form double/triple bonds if necessary.

• Example: CO2 has two double bonds between C and O.

Resonance Structures

Resonance structures are alternative Lewis structures for a molecule that differ only in the placement of electrons.

• Key Point: Resonance stabilizes molecules by delocalizing electrons.

• Example: Ozone (O3) has two resonance structures.

Formal Charge Assignment

Formal charge helps identify the most stable Lewis structure.

• Formula: $\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2} \text{Bonding electrons})$

• Example: In NO3-, formal charges help determine the best resonance structure.

Exceptions to the Octet Rule

Some molecules have atoms with less or more than eight valence electrons.

• Less than 8: Common for H (duet), Be, B.

• More than 8: Expanded octet possible for elements in period 3 or higher (e.g., SF6).

• Example: BF3 has only 6 electrons around B.

Bond Strength and Length

The strength and length of a bond depend on the number of shared electron pairs.

• Single Bond: Longest and weakest.

• Double Bond: Intermediate length and strength.

• Triple Bond: Shortest and strongest.

• Example: N≡N in N2 is a triple bond, very strong and short.

VSEPR Shapes and Bond Angles

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Common Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Bond Angles: Determined by shape and presence of lone pairs (lone pairs decrease bond angles).

• Example: Water (H2O) is bent with a bond angle of ~104.5° due to two lone pairs on O.

Predicting Molecular Shape and Polarity

Molecular shape and bond polarity together determine if a molecule is polar or nonpolar.

• Shape: Use VSEPR to predict geometry.

• Polarity: If the shape is asymmetrical and bonds are polar, the molecule is polar.

• Example: CO2 is linear and nonpolar; H2O is bent and polar.

Chapter 6: Chemical Bonding II – Valence Bond and Molecular Orbital Theory

Hybridization of Atoms in Molecules

Hybridization explains the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Types: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), sp3d, sp3d2.

• Example: Methane (CH4) has sp3 hybridization.

Sigma and Pi Bonds

Sigma (σ) and pi (π) bonds are types of covalent bonds formed by orbital overlap.

• Sigma Bond: Formed by head-on overlap; every single bond is a sigma bond.

• Pi Bond: Formed by side-on overlap; present in double and triple bonds (double = 1 σ + 1 π, triple = 1 σ + 2 π).

• Example: Ethylene (C2H4) has a double bond: 1 σ and 1 π.

Molecular Orbital (MO) Theory for Homonuclear Diatomic Molecules/Ions

MO theory describes bonding by combining atomic orbitals to form molecular orbitals, which can be bonding or antibonding.

• Bonding MO: Lower energy, stabilizes molecule.

• Antibonding MO: Higher energy, destabilizes molecule.

• Example: O2 has both bonding and antibonding molecular orbitals.

Bond Order and Magnetic Behavior

Bond order indicates the strength and stability of a bond; magnetic behavior depends on unpaired electrons in molecular orbitals.

• Bond Order Formula: $\text{Bond order} = \frac{1}{2} (\text{Number of electrons in bonding MOs} - \text{Number in antibonding MOs})$

• Magnetic Behavior: Molecules with unpaired electrons are paramagnetic; those with all electrons paired are diamagnetic.

• Example: O2 is paramagnetic due to two unpaired electrons.

Bond Type	Bond Order	Bond Strength	Bond Length
Single	1	Weakest	Longest
Double	2	Intermediate	Intermediate
Triple	3	Strongest	Shortest

Additional info: Figure 6.10 typically shows molecular orbital diagrams for diatomic molecules, illustrating the arrangement of electrons in bonding and antibonding orbitals.