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 bonded atoms, leading to unequal sharing of electrons.

• Electronegativity: The ability of an atom to attract shared electrons in a chemical bond.

• Trend in Electronegativity:

  • Increases across a period (left to right) on the periodic table.

  • Decreases down a group (top to bottom).

  • Fluorine is the most electronegative element.

• Bond Polarity:

  • Nonpolar covalent: Electronegativity difference < 0.5

  • Polar covalent: Electronegativity difference between 0.5 and 1.7

  • Ionic: Electronegativity difference > 1.7

• Example: In HCl, Cl is more electronegative than H, so the bond is polar.

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules.

• Steps to Draw Lewis Structures:

  1. Count total valence electrons for all atoms.

  2. Draw a skeletal structure, connecting atoms with single bonds.

  3. Distribute remaining electrons as lone pairs to complete octets (or duets for H).

  4. Form double or triple bonds if necessary to satisfy the octet rule.

• Example: For CO2, carbon forms two double bonds with oxygen.

Resonance Structures

Some molecules cannot be represented by a single Lewis structure; instead, they have multiple resonance forms.

• Resonance: The actual structure is a hybrid of all possible resonance forms.

• Example: Ozone (O3) has two resonance structures with different O=O and O–O bonds.

Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula:

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

• Lowest formal charges (closest to zero) are preferred.

• Example: In the nitrate ion (NO3–), formal charges are minimized by delocalizing electrons.

Exceptions to the Octet Rule

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

• Incomplete octet: Elements like Be and B can have fewer than 8 electrons (e.g., BF3).

• Expanded octet: Elements in period 3 or higher (e.g., SF6) can have more than 8 electrons.

• Odd-electron species: Molecules with an odd number of electrons (e.g., NO).

Bond Strength and Bond 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 (triple bond) is much stronger and shorter than N–N (single bond).

VSEPR Theory: Shapes and Bond Angles

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

• Electron domains (bonding and lone pairs) arrange to minimize repulsion.

• Common shapes and bond angles:

  • Linear: 180°

  • Trigonal planar: 120°

  • Tetrahedral: 109.5°

  • Trigonal bipyramidal: 90°, 120°

  • Octahedral: 90°

• Lone pairs decrease bond angles due to greater repulsion.

• Example: NH3 (ammonia) is trigonal pyramidal with a bond angle of ~107°.

Predicting Molecular Shape and Polarity

The shape and distribution of polar bonds determine the overall polarity of a molecule.

• Nonpolar molecules: Symmetrical shape with polar bonds canceling out (e.g., CO2).

• Polar molecules: Asymmetrical shape or uneven distribution of polar bonds (e.g., H2O).

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

Hybridization of Atoms in Molecules

Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals for bonding.

• Types of hybridization:

  • sp: Linear geometry (2 electron domains)

  • sp2: Trigonal planar geometry (3 electron domains)

  • sp3: Tetrahedral geometry (4 electron domains)

  • sp3d: Trigonal bipyramidal geometry (5 electron domains)

  • sp3d2: Octahedral geometry (6 electron domains)

• Example: In methane (CH4), carbon is sp3 hybridized.

Sigma (σ) and Pi (π) Bonds

Sigma and pi bonds are types of covalent bonds formed by different orbital overlaps.

• Sigma (σ) bond: Formed by head-on overlap of orbitals; all single bonds are sigma bonds.

• Pi (π) bond: Formed by side-to-side overlap of p orbitals; present in double and triple bonds.

• Example: In ethene (C2H4), the C=C bond consists of one sigma and one pi bond.

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

MO theory describes the combination of atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.

• Bonding and antibonding orbitals: Constructive overlap forms bonding orbitals (lower energy), destructive overlap forms antibonding orbitals (higher energy).

• Order of filling: For elements up to N2, the order is: σ1s, σ1s*, σ2s, σ2s*, π2p, σ2p, π2p*, σ2p*.

• Example: O2 has unpaired electrons in π2p* orbitals, making it paramagnetic.

Bond Order and Magnetic Behavior

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

• Bond order formula:

  $\text{Bond order} = \frac{1}{2} (\text{Number of bonding electrons} - \text{Number of antibonding electrons})$

• Magnetic behavior:

  • Paramagnetic: Molecules with unpaired electrons (attracted to magnetic fields).

  • Diamagnetic: Molecules with all electrons paired (repelled by magnetic fields).

• Example: N2 has a bond order of 3 and is diamagnetic; O2 has a bond order of 2 and is paramagnetic.

Table: Summary of Bond Types and Properties

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

Additional info: For MO diagrams, refer to provided figures in your course materials for specific orbital energy ordering in diatomic molecules.