Chemical Bonding I: Basic Concepts – Study Notes

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Chemical Bonding I: Basic Concepts

Introduction

Chemical bonding is a fundamental concept in chemistry that explains how atoms combine to form molecules and compounds. This chapter introduces the basic principles of chemical bonding, focusing on Lewis theory, covalent and ionic bonds, resonance, exceptions to the octet rule, molecular shapes, bond order, bond length, and bond energies.

Lewis Theory: An Overview

Valence Electrons and Bonding

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding.
Atoms achieve stability by attaining a noble gas configuration, often an octet (eight electrons) in their valence shell.
Ionic bonds form by the transfer of electrons from one atom to another, resulting in oppositely charged ions.
Covalent bonds form by the sharing of electrons between atoms.

Portrait of Gilbert Newton Lewis

Lewis Symbols and Lewis Structures

A Lewis symbol represents the nucleus and core electrons of an atom with its chemical symbol, and valence electrons as dots around it.
Lewis structures show how valence electrons are arranged among atoms in a molecule or ion.

Lewis structure examples

Ionic and Covalent Bonding Examples

Ionic bonding: Electron transfer from Na to Cl forms Na+ and Cl- ions.
Covalent bonding: Electron sharing between H and Cl forms HCl.

Ionic crystal structure

Covalent Bonding: An Introduction

Single, Double, and Triple Bonds

A single bond involves one shared pair of electrons (e.g., H2).
A double bond involves two shared pairs (e.g., O2).
A triple bond involves three shared pairs (e.g., N2).

Multiple covalent bonds

Coordinate Covalent Bonds

A coordinate covalent bond (dative bond) forms when both electrons in a shared pair come from the same atom.
Example: Formation of the ammonium ion (NH4+).

Lewis structure of ammonium ion

Polar Covalent Bonds and Electrostatic Potential Maps

Bond Polarity and Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a bond.
A polar covalent bond has unequal sharing of electrons, resulting in partial charges (δ+ and δ−).
The greater the difference in electronegativity, the more polar the bond.

Electrostatic potential maps for NaCl, HCl, and Cl2

Electrostatic Potential Maps

These maps visually represent the distribution of electron density and partial charges in molecules.
Red regions indicate electron-rich (negative potential), blue regions indicate electron-poor (positive potential).

Electrostatic potential map for ammonia

Electronegativity Trends

Electronegativity increases across a period and decreases down a group in the periodic table.

Electronegativity values in the periodic table

Ionic Character and Bond Polarity

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

Percent ionic character vs. electronegativity difference

Writing Lewis Structures

General Rules

All valence electrons must be shown.
Usually, electrons are paired, and each atom (except H) achieves an octet.
Hydrogen requires only 2 electrons.
Multiple bonds may be needed for some atoms (C, N, O, S, P).

Steps for Drawing Lewis Structures

Count total valence electrons.
Draw a skeletal structure, placing the least electronegative atom in the center (except H, which is always terminal).
Complete octets for terminal atoms, then central atom.
Form multiple bonds if necessary to complete octets.

Flowchart for drawing Lewis structures

Formal Charge

Formal charge (FC) helps determine the most stable Lewis structure.
FC = (valence electrons) − (lone pair electrons) − (1/2 × bonding electrons)
The sum of formal charges equals the overall charge of the molecule or ion.
Structures with the smallest formal charges and negative charges on the most electronegative atoms are preferred.

Resonance

Resonance Structures

Some molecules cannot be represented by a single Lewis structure; instead, they have resonance structures.
The actual structure is a resonance hybrid, an average of all possible resonance forms.

Electrostatic potential map of ozone (resonance)

Exceptions to the Octet Rule

Odd-Electron Species

Some molecules have an odd number of electrons and cannot achieve octets for all atoms (e.g., NO).

Incomplete Octets

Some atoms (e.g., B, Be) are stable with fewer than 8 electrons.

Expanded Valence Shells

Atoms in period 3 or beyond can have more than 8 electrons (e.g., SF6, PCl5).

Normal octet vs. expanded valence shell

Shapes of Molecules

Bond Length and Bond Angle

Bond length is the distance between the nuclei of two bonded atoms.
Bond angle is the angle between adjacent bonds.

Bond length and bond angle in water

Valence-Shell Electron Pair Repulsion (VSEPR) Theory

Electron pairs (bonding and lone pairs) repel each other and arrange themselves to minimize repulsion.
Electron group geometry describes the arrangement of electron groups around a central atom.
Molecular geometry describes the arrangement of atoms (nuclei) in a molecule.

Balloon analogy for VSEPR Electron-group geometries

Common Electron-Group Geometries

2 groups: Linear (180°)
3 groups: Trigonal planar (120°)
4 groups: Tetrahedral (109.5°)
5 groups: Trigonal bipyramidal (90°, 120°)
6 groups: Octahedral (90°)

Molecular Shapes and Dipole Moments

The shape of a molecule affects its polarity and dipole moment.
Polar molecules align in an electric field due to their dipole moments.

Polar molecules in an electric field Molecular shapes and dipole moments

Bond Order and Bond Length

Bond Order

Bond order is the number of shared electron pairs between two atoms.
Single bond: bond order = 1; double bond: bond order = 2; triple bond: bond order = 3.
For resonance structures, bond order is averaged over all structures.

Bond Length

Bond length decreases as bond order increases.
Bond length can be estimated as the sum of the covalent radii of the two atoms.

Table of average bond lengths

Bond Energies

Bond Dissociation Energy

Bond energy is the energy required to break one mole of a specific bond in a molecule in the gas phase.
Bond energies can be used to estimate the enthalpy change (ΔH) of a reaction:

Table of average bond energies

Predicting Reaction Enthalpy

If more energy is released in forming bonds than is required to break bonds, the reaction is exothermic.
If more energy is required to break bonds than is released in forming bonds, the reaction is endothermic.

Additional info: These notes cover the essential concepts of chemical bonding, including Lewis structures, resonance, exceptions to the octet rule, molecular geometry, bond order, bond length, and bond energies, as outlined in a standard general chemistry curriculum.