Chemical Bonding I: Lewis Structures, Molecular Shapes, and Bond Properties

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Chemical Bonding I

Introduction

This chapter explores the fundamental concepts of chemical bonding, including the differences between ionic and covalent bonds, how to draw Lewis structures, resonance, formal charge, exceptions to the octet rule, bond energies and lengths, and the Valence Shell Electron-Pair Repulsion (VSEPR) theory for predicting molecular shapes and polarity.

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a chemical bond. It varies across the periodic table, increasing from left to right and decreasing from top to bottom. Fluorine is the most electronegative element, assigned a value of 4.0.

Definition: Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.
Periodic Trend: Electronegativity increases across a period and decreases down a group.

Periodic table showing electronegativity trends

Bond Polarity and Types

Bonds can be classified based on the difference in electronegativity (ΔEN) between the atoms involved:

Nonpolar Covalent: Electrons are shared equally (ΔEN ≈ 0).
Polar Covalent: Electrons are shared unequally (ΔEN between 0.4 and 2.0).
Ionic: Electrons are transferred (ΔEN > 2.0).

Continuum of bond types based on electronegativity difference

Electronegativity Difference (ΔEN)	Bond Type	Example
Small (0–0.4)	Covalent	Cl2
Intermediate (0.4–2.0)	Polar covalent	HCl
Large (2.0+)	Ionic	NaCl

Table showing effect of electronegativity difference on bond type

Bond Polarity and Dipole Moments

Polar covalent bonds have a dipole moment, which is a measure of the separation of charge. The dipole moment (μ) occurs when there is a difference in electronegativity between two atoms, resulting in partial positive (δ+) and partial negative (δ−) charges.

HF molecules align with an electric field due to their dipole moment Electron-rich and electron-poor regions in a polar bond

Percent Ionic Character

The percent ionic character of a bond increases with the electronegativity difference between the atoms.

Graph showing percent ionic character vs. electronegativity difference

Writing Lewis Structures for Molecules and Polyatomic Ions

Steps for Drawing Lewis Structures

Skeletal Structure: Place the least electronegative atom in the center. Hydrogen and electronegative atoms are always terminal.
Count Valence Electrons: Sum the valence electrons for all atoms, adjusting for charge if the species is an ion.
Place Electrons: Assign two electrons (a single bond) between each atom, then complete duets for hydrogen and octets for other atoms.
Octet Rule: If any atom lacks an octet, move lone pairs to form multiple bonds.
Check Electron Count: Ensure the total number of electrons matches the calculated sum.

Examples of Lewis Structures

Lewis structures for SF2, HCOOH, PH3, O22−, and nitrogen trifluoride (NF3) are commonly used for practice.

Possible Lewis structures for nitrogen trifluoride

Resonance Structures

Definition and Explanation

A resonance structure is one of two or more valid Lewis structures for a molecule that differ only in the arrangement of electrons. The actual molecule is a resonance hybrid, an intermediate between the contributing structures.

Resonance Hybrid: The true structure is a blend of all resonance forms.

Resonance hybrid structure example

Formal Charge

Calculating Formal Charge

Formal charge (FC) is calculated to determine the most stable Lewis structure:

Formula:
Guidelines:
- Sum of all FC in a neutral molecule = 0
- Sum of all FC in an ion = charge of ion
- The smaller the FC, the more stable the structure
- Negative FC should reside on the more electronegative atom

Table showing formal charge calculations for different resonance structures

Choosing the Best Lewis Structure

Use formal charge to select the most stable Lewis structure for molecules such as CH3SOCH3.

Possible Lewis structures for CH3SOCH3

Exceptions to the Octet Rule

Types of Exceptions

Odd Electron Species: Molecules with an odd number of electrons (free radicals), e.g., ClO.
Incomplete Octets: Atoms such as B in BF3 may have fewer than eight electrons.
Expanded Octets: Elements in the third period and beyond can have more than eight electrons, e.g., AsF5, SF6, XeF2.

Lewis structures showing expanded octets for AsF5 and SF6

Bond Energies and Bond Lengths

Bond Energy

Bond energy is the energy required to break one mole of a bond in the gas phase. Stronger bonds require more energy to break.

Example: HCl bond energy = 431 kJ/mol; HF bond energy = 565 kJ/mol.

Bond	Bond Energy (kJ/mol)
H–H	436
H–F	565
H–Cl	431
H–Br	364
H–I	297

Table of average bond energies

Bond Length

Bond length is the distance between the nuclei of two bonded atoms. Generally, shorter bonds are stronger.

Bond	Bond Length (pm)
H–H	74
H–F	92
H–Cl	127
H–Br	141
H–I	161

Bond lengths for halogen molecules Table of average bond lengths

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

Basic Principles

VSEPR theory states that electron groups (lone pairs, single, double, or triple bonds) around a central atom repel each other and arrange themselves to minimize repulsion. The number of electron groups determines the molecular shape.

2 groups: Linear
3 groups: Trigonal planar
4 groups: Tetrahedral
5 groups: Trigonal bipyramidal
6 groups: Octahedral

Common molecular geometries

Linear Geometry

Linear geometry occurs when there are two electron groups around the central atom, resulting in a bond angle of 180°.

Linear geometry for BeCl2 Linear geometry for CO2

Trigonal Planar Geometry

Three electron groups around the central atom form a trigonal planar geometry with bond angles of 120°.

Trigonal planar geometry for BF3 Trigonal planar geometry for CH2O

Tetrahedral Geometry

Four electron groups form a tetrahedral geometry with bond angles of 109.5°.

Tetrahedral geometry and tetrahedron shape Tetrahedral geometry for CH4

Trigonal Bipyramidal Geometry

Five electron groups form a trigonal bipyramidal geometry with bond angles of 90° and 120°.

Trigonal bipyramidal geometry and shape

Octahedral Geometry

Six electron groups form an octahedral geometry with bond angles of 90°.

Octahedral geometry and shape Octahedral geometry for SF6

VSEPR Theory: The Effect of Lone Pairs

Lone Pairs and Molecular Geometry

Lone pairs affect molecular geometry by reducing bond angles. For example, in tetrahedral electron geometry:

4 bonding pairs: tetrahedral (CH4)
3 bonding pairs, 1 lone pair: trigonal pyramidal (NH3)
2 bonding pairs, 2 lone pairs: bent (H2O)

Effect of lone pairs on molecular geometry Effect of lone pairs on bond angles for CH4, NH3, and H2O

Expanded Octets and Molecular Geometry

For molecules with five or six electron groups, lone pairs can lead to geometries such as see-saw, T-shaped, square pyramidal, and square planar.

Electron and molecular geometry for XeF2

VSEPR Theory: Predicting Molecular Geometries

Summary Table of Geometries

The following table summarizes electron groups, lone pairs, electron geometry, molecular geometry, bond angles, and examples:

Table summarizing electron and molecular geometries

Molecular Shape and Polarity

Polarity of Molecules

The polarity of a molecule depends on the presence of polar bonds and the molecular shape. If the dipoles do not cancel, the molecule is polar and has a net dipole moment.

Diatomic molecules: If the bond is polar, the molecule is polar.
Adding Dipoles: Vector addition of bond dipoles determines overall molecular polarity.

Adding dipoles to determine molecular polarity

Practice: Determining Polarity

Determine whether the following molecules are polar or nonpolar:

CH2F2
SiCl4
SeCl6
SF2

Example: Glycine

Assign a geometry to each interior atom in glycine, a simple amino acid.

Lewis structure of glycine with interior atoms

Conclusion

This chapter provides a comprehensive overview of chemical bonding, including how to draw and interpret Lewis structures, understand resonance and formal charge, recognize exceptions to the octet rule, and predict molecular shapes and polarity using VSEPR theory. Mastery of these concepts is essential for understanding the structure and properties of molecules in chemistry.