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

Electronegativity and Bond Polarity

Electronegativity is a measure of an atom’s ability to attract electrons in a chemical bond. The difference in electronegativity between two atoms determines the type and polarity of the bond formed.

• Nonpolar Covalent Bonds: Formed between atoms with similar electronegativities; electrons are shared equally.

• Polar Covalent Bonds: Formed when there is a moderate difference in electronegativity; electrons are shared unequally.

• Ionic Bonds: Formed when the difference in electronegativity is large; electrons are transferred from one atom to another.

Example: The O–H bond is more polar than N–H or C–H because oxygen is more electronegative than nitrogen or carbon.

Lewis Structures and Valence Electrons

Lewis structures represent the arrangement of valence electrons in molecules and polyatomic ions. Mastery of Lewis structures is essential for understanding molecular geometry, formal charge, resonance, and hybridization.

1. Choose a skeletal structure (least electronegative atom is central; hydrogen is always terminal).

2. Count total valence electrons (add for anions, subtract for cations).

3. Connect atoms with single bonds (2 electrons per bond).

4. Distribute remaining electrons as lone pairs to terminal atoms first, then to the central atom.

5. If the central atom lacks an octet, form double or triple bonds as needed.

6. Check that the total number of electrons matches the initial count.

Example: For formaldehyde (H2CO):

• Carbon is central, hydrogens are terminal.

• Total valence electrons: 12.

• Single bonds: H–C–H and C–O (6 electrons).

• Lone pairs on oxygen; convert one lone pair to a double bond for carbon’s octet.

• Final structure: H2C=O with two lone pairs on O.

Resonance and Formal Charge

Some molecules cannot be represented by a single Lewis structure. Resonance structures differ only in the placement of electrons, not atoms. The actual molecule is a resonance hybrid of all valid structures.

• Formal Charge: Used to determine the most stable resonance structure.

Formal Charge Formula:

Example: In ClO3−, all three Cl–O bonds are equivalent with a bond order of 5/3.

Exceptions to the Octet Rule

Not all atoms follow the octet rule. Recognize the following exceptions:

• Odd-Electron Species: Molecules with an odd number of electrons (free radicals).

• Incomplete Octets: Atoms like Be and B can have fewer than 8 electrons.

• Expanded Octets: Atoms in period 3 or beyond (e.g., P, S) can have more than 8 electrons.

Example: Phosphorus (P) can form compounds with expanded octets.

Bond Order, Length, and Energy

Bond order is the number of chemical bonds between a pair of atoms. It is related to bond length and bond energy.

• Higher bond order → shorter bond length and stronger bond.

• Single < double < triple bond in terms of strength and bond order.

Example: CH3CH3 (single C–C bond) has the weakest C–C bond compared to CH2CH2 (double) and HCCH (triple).

VSEPR: Electron and Molecular Geometry

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on electron group repulsions.

• Count electron groups (bonds and lone pairs) around the central atom.

• Determine electron geometry (arrangement of all electron groups).

• Determine molecular geometry (arrangement of atoms only).

• Lone pairs affect molecular geometry but not electron geometry.

Example: SiI4 has 4 bonding groups, no lone pairs: both electron and molecular geometry are tetrahedral.

Molecular Polarity

A molecule is polar if it contains polar bonds and the molecular geometry does not cancel out the bond dipoles.

• Symmetrical molecules with identical surrounding atoms are nonpolar, even if bonds are polar.

• Asymmetrical molecules or those with different surrounding atoms are often polar.

Example: Of KrBr2, OCCl2, PBr5, and SCl6, only OCCl2 is polar.

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

Valence Bond Theory and Hybridization

Valence bond theory explains bonding as the overlap of atomic orbitals. Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals for bonding.

• Each electron group (single, double, triple bond, or lone pair) counts as one group for hybridization.

• Electron geometry determines hybridization:

Example: A central atom with 5 electron groups (e.g., 4 single bonds and 1 double bond) is sp3d hybridized.

Sigma (σ) and Pi (π) Bonds

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

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

• Pi (π) bonds: Formed by side-on overlap; present in double and triple bonds.

• Double bond = 1 σ + 1 π; Triple bond = 1 σ + 2 π.

Example: Propyne (HCCCH3) has 6 σ bonds and 2 π bonds.

Rotation Around Bonds

Rotation about bonds depends on the type of bond:

• Rotation is allowed around single (σ) bonds.

• Rotation is restricted around double (π) bonds because rotation would break the π bond.

Molecular Orbital (MO) Theory

MO theory describes bonding in terms of molecular orbitals formed by the linear combination of atomic orbitals. It explains bond order, magnetism, and the distribution of electrons in molecules.

• Bonding orbitals: Lower energy, stabilize the molecule.

• Antibonding orbitals: Higher energy, destabilize the molecule.

• Nonbonding orbitals: Localized on one atom, do not affect bond order.

• Bond Order Formula:

• Higher bond order means a stronger, shorter bond.

• MO diagrams can be used to predict if a molecule is paramagnetic (unpaired electrons) or diamagnetic (all electrons paired).

Isomers

Isomers are compounds with the same molecular formula but different arrangements of atoms or bonds.

• Structural isomers: Differ in the connectivity of atoms.

• Stereoisomers: Same connectivity, different spatial arrangement.

Exam Readiness Checklist

• Draw Lewis structures quickly and accurately.

• Assign formal charges correctly.

• Predict electron geometry, molecular geometry, polarity, and hybridization.

• Distinguish between σ and π bonding.

• Use MO diagrams to predict bond order and magnetism.