Lewis Structures, Resonance, and Molecular Geometry: Study Notes for General Chemistry

Chemical Bonding I: The Lewis Model

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They are essential for understanding molecular structure, formal charges, and resonance.

• Valence Electrons: The electrons in the outermost shell of an atom, involved in bonding.

• Bond Skeleton: The arrangement of atoms connected by single bonds before adding lone pairs or multiple bonds.

• Resonance Structures: Multiple valid Lewis structures for a molecule, differing only in the placement of electrons.

• Formal Charge: The charge assigned to an atom in a molecule, calculated by: $\text{Formal Charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2}\text{Bonding electrons})$

Example: For CO2, the Lewis structure shows double bonds between C and each O, with no formal charge on any atom.

Table: Lewis Structure Analysis

Molecule/Ion	Valence Electrons	Bond Skeleton	Resonance Electrons	Lewis Structure with Formal Charges	Resonance Structure (Y/N)
Cl2	14	Cl–Cl	None	Each Cl has 3 lone pairs	N
CO2	16	O=C=O	None	No formal charge	N
PO43–	32	P surrounded by 4 O	Yes	Formal charges on O	Y
NO3–	24	N surrounded by 3 O	Yes	Formal charge on N and O	Y
SO3	24	S surrounded by 3 O	Yes	Formal charge on S and O	Y
XeF2	22	Xe–F–F	None	Formal charge on Xe	N
I2	14	I–I	None	Each I has 3 lone pairs	N
NO2	17	N–O–O	Yes	Formal charge on N and O	Y
CaBr2	Not applicable (ionic)	Ca2+ and 2 Br–	None	Ionic structure	N

Additional info: Table entries inferred from standard Lewis structure rules.

Molecular Shapes & Valence Bond Theory

VSEPR Theory and Molecular Geometry

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on the repulsion between electron pairs around a central atom.

• Electron Domains: Regions of electron density (bonds or lone pairs) around a central atom.

• Molecular Geometry: The arrangement of atoms in space, determined by the number of bonding and nonbonding electron pairs.

• Bond Angles: The angles between adjacent bonds, affected by lone pairs and multiple bonds.

• Dipole Moment: A measure of molecular polarity; molecules with polar bonds and asymmetric shapes have a dipole moment.

Example: H2O has a bent geometry with a bond angle of approximately 104.5°, and a net dipole moment.

Table: VSEPR and Molecular Geometry Analysis

Molecule	No. of Electron Domains	No. of Lone Pairs	Molecular Geometry	Bond Angle	Dipole Moment (Y/N)
H2S	4	2	Bent	~104.5°	Y
NO2	3	1	Bent	~120°	Y
NO3–	3	0	Trigonal planar	120°	N
CF4	4	0	Tetrahedral	109.5°	N
PF5	5	0	Trigonal bipyramidal	90°, 120°	N
OCl2	4	2	Bent	~104.5°	Y
PbBr4	4	0	Tetrahedral	109.5°	N
SeCl4	5	1	Seesaw	~90°, ~120°	Y
I2	2	6	Linear	180°	N

Additional info: Table entries inferred from VSEPR theory and standard molecular geometries.

Chemical Bonding II: Resonance and Bond Classification

Resonance Structures

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. The actual structure is a hybrid of all resonance forms.

• PO43–: Has multiple resonance structures with different locations for double bonds and formal charges.

• CO2: No resonance; only one valid Lewis structure.

Bond Classification

Bonds between atoms can be classified as:

• Pure Covalent: Electrons shared equally (e.g., H2, Cl2).

• Polar Covalent: Electrons shared unequally due to differences in electronegativity (e.g., HCl, CO).

• Ionic: Electrons transferred from one atom to another (e.g., NaCl, CaBr2).

Example: N–Cl and S–O are polar covalent; Cl–Cl and O–O are pure covalent.

Nitrogen-Nitrogen Bond Strength

The strength of a nitrogen-nitrogen bond depends on the bond order (single, double, triple). A triple bond (as in N2) is the strongest.

• Strongest N–N bond: N2 (triple bond)

• Weakest N–N bond: N2H4 (single bond)

Comparing Bond Types

• CO: Polar covalent

• CO2: Polar covalent (overall molecule is nonpolar)

• O2: Pure covalent

• CF: Polar covalent

Additional info: Bond classifications inferred from electronegativity differences and molecular structure.