Chemical Bonding I: Molecular Shapes and Polarity – Study Notes

Chapter 11: Chemical Bonding I – Molecular Shapes and Polarity

Introduction

This chapter introduces the fundamental concepts of chemical bonding with a focus on the structure of molecular (covalent) compounds. Key topics include Lewis structures, the VSEPR model for predicting molecular shapes, and the concept of molecular polarity. Understanding these concepts is essential for predicting the properties and reactivity of chemical compounds.

Learning Goals

• Determine the number of valence electrons for an atom and write its Lewis symbol.

• Recognize how the octet rule applies to the arrangement of electrons in the valence shell.

• Draw Lewis structures for atoms, ions, and covalent compounds, including cases with multiple bonds, resonance structures, expanded valence shells, deficient electrons, and odd electrons.

• Use electronegativity differences between bonding atoms to classify bonds as polar covalent, nonpolar covalent, or ionic.

• Calculate the enthalpy of reaction using bond energies.

• Predict electron pair geometry and molecular geometry using the VSEPR model.

• Predict whether a molecule can have a net dipole moment from the molecular shape.

Key Terms and Definitions

• Chemical bond: The force that holds atoms together in compounds.

• Ionic bond: Electrostatic attraction between a positively charged cation and a negatively charged anion.

• Covalent bond: Sharing of electron pairs between atoms.

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

• Lewis electron dot symbol: Representation of an atom showing valence electrons as dots around the element symbol.

• Lone pair: Pair of valence electrons not involved in bonding.

• Bond pair: Pair of electrons shared between two atoms in a covalent bond.

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Formal charge: The charge assigned to an atom in a molecule, calculated by: $\text{FC} = \text{(valence electrons)} - \text{(nonbonding electrons)} - \frac{1}{2}\text{(bonding electrons)}$

• Resonance: Delocalization of electrons in molecules where more than one valid Lewis structure can be drawn.

• VSEPR model: Valence Shell Electron Pair Repulsion model, used to predict molecular geometry.

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

• Bond polarity: The distribution of electrical charge over the atoms joined by the bond.

• Dipole moment: A measure of the separation of positive and negative charges in a molecule.

• Bond energy: The energy required to break a chemical bond.

• Bond length: The average distance between the nuclei of two bonded atoms.

Valence Electrons and Lewis Symbols

Main Group Elements

Valence electrons are the s- and p-electrons in the outermost shell. The number of valence electrons for main group elements equals the group number.

• Example: Carbon (Group 4A) has 4 valence electrons.

Transition Metals

Valence electrons include the ns and (n–1)d orbitals.

Lewis Symbols for Atoms

Lewis symbols represent the nucleus and core electrons by the element symbol, with valence electrons shown as dots.

Group	Lewis Symbol
1A	Li·
2A	Be··
3A	B···
4A	C····
5A	N·····
6A	O······
7A	F·······
8A	Ne········

An octet of electrons is a stable configuration, corresponding to a noble gas electron configuration.

Types of Chemical Bonds

Ionic Bonding

Ionic bonding is the electrostatic attraction between a cation and an anion. Electrons are transferred from one atom to another, resulting in charged ions that attract each other.

• Example: Formation of NaCl from Na and Cl.

Covalent Bonding

Covalent bonding involves sharing valence electrons between atoms. A single bond consists of a shared pair of electrons, attracted to both nuclei.

• Example: Hydrogen molecule (H2) – electrons are shared equally (nonpolar covalent bond).

• Example: Hydrogen fluoride (HF) – electrons are shared unequally (polar covalent bond).

Electronegativity and Bond Polarity

Periodic Trend in Electronegativity

Electronegativity increases across a period and decreases down a group. Fluorine is the most electronegative element.

Type of Bond	Electronegativity Difference	Example
Nonpolar covalent	0 – 0.4	C–H ($\Delta EN = 0.4$)
Polar covalent	0.4 – 2.0	H–Cl ($\Delta EN = 0.9$)
Ionic	2.0 or greater	KBr ($\Delta EN = 2.2$)

There is a continuum of bonding possibilities from nonpolar covalent to ionic.

Electronegativity Values for Main Group Elements

Element	Electronegativity
H	2.1
C	2.5
N	3.0
O	3.5
F	4.0
Cl	3.0
Br	2.8
I	2.5

Comparison of Ionic and Covalent Compounds

Characteristic	Ionic Compounds	Covalent Compounds
Unit of Composition	Formula unit of cations and anions	Molecule
Bond Formation	Transfer of electrons from one atom to another	Sharing of electron pairs between atoms
Types of atoms or ions involved	Metal cation and nonmetal anion or polyatomic anion	Usually nonmetal atoms
Physical state at room temperature (25°C)	Solid	Solid, liquid, or gas
Melting Point and Boiling Point	High (above 300°C)	Moderate to low
Electrical conductivity of an aqueous solution	Good conductors (soluble ionic compounds are strong electrolytes)	Nonconductors (except acids and bases)

Lewis Structures and the Octet Rule

Lewis Structure Construction

Lewis structures show the arrangement of bonding and nonbonding valence electrons in a molecule. The octet rule states that atoms tend to be surrounded by eight electrons.

• In simple molecules, Lewis structures can be drawn by joining Lewis symbols so the octet rule is obeyed.

• Examples: HCl, Cl2, O2, H2O, NH3

Bond Energies and Bond Lengths

Bond Energy

Bond energy is the energy required to break a bond. It is a measure of bond strength.

• Energy is released when a new bond is formed (exothermic process).

• Bond energies for single bonds range from about 150–550 kJ/mol.

Bond Length

Bond length is the average distance between the nuclei of two bonded atoms. Multiple bonds (double, triple) are shorter and stronger than single bonds.

Calculating Enthalpy Change Using Bond Energies

The overall energy (enthalpy) change for a reaction can be calculated from average bond energies:

$\Delta H = \sum \text{BE (bonds broken)} - \sum \text{BE (bonds made)}$

VSEPR Model and Molecular Geometry

VSEPR Model

The Valence Shell Electron Pair Repulsion (VSEPR) model predicts the 3D shape of molecules based on the repulsion between electron pairs around a central atom.

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

• Common geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Molecular Polarity and Dipole Moments

Bond Dipole and Molecular Dipole

A bond dipole is a measure of charge separation in a polar bond. The molecular dipole moment is the vector sum of all bond dipoles in a molecule.

• Dipole moments are measured in Debye (D).

• Polar molecules align in an external electric field.

• Molecular polarity depends on both bond polarity and molecular geometry.

Predicting Molecular Polarity

• Molecules with symmetrical geometry and identical terminal atoms are generally nonpolar (e.g., CO2, CF4).

• Molecules with lone pairs on the central atom or different terminal atoms are generally polar (e.g., H2O, NH3).

Summary Table: Bond Types and Polarity

Bond Type	Electronegativity Difference	Example
Nonpolar covalent	0 – 0.4	C–H
Polar covalent	0.4 – 2.0	H–Cl
Ionic	2.0 or greater	KBr

Additional info:

• Metallic bonding (delocalized electrons) is another type of chemical bonding, but is not covered in this chapter.

• Formal charge calculations help determine the most reasonable resonance structure.

• Expanded valence shells are possible for elements in the third period and beyond.