Chemical Bonding I – Lewis Theory (General Chemistry Study Notes)

Topic 3: Chemical Bonding I – Lewis Theory

Overview

This topic introduces the fundamental concepts of chemical bonding, focusing on Lewis Theory. It covers the types of chemical bonds, the ionic bonding model, representation of valence electrons, Lewis structures, electronegativity and bond polarity, covalent bond energies, lengths and vibrations, resonance, formal charge, and exceptions to the octet rule.

Types of Chemical Bonding

Main Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. There are three major types:

• Metallic Bonds: Involve the sharing of free electrons among a lattice of metal atoms.

• Ionic Bonds: Formed by the electrostatic attraction between cations and anions.

• Covalent Bonds: Involve the sharing of electron pairs between atoms, typically non-metals.

Type of bond	Description of interaction	Example
Metallic	Delocalized electrons among metal atoms	Cu, Fe
Ionic	Electrostatic attraction between ions	NaCl
Covalent	Electron sharing between atoms	H2O, CO2

Comparison of Bond Types and Properties

	Metallic Compounds	Ionic Compounds	Covalent Compounds
Constituent Particles	Metals	Crystalline Solids	Gases, liquids, or solids
Melting/Boiling Point	Extremely high	High	Low
Electrical Conductivity	Good conductors	Conductors (melted or aqueous)	Poor conductors
Water Solubility	Not water soluble	Water soluble (many)	Not water soluble (many)

The Ionic Bonding Model

Formation and Structure

Ionic compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions), resulting in a repeating pattern known as a crystal lattice. This typically occurs between metals and non-metals, such as sodium chloride (NaCl).

• Low ionization energy: Metals lose electrons easily.

• High electron affinity: Non-metals gain electrons easily.

• Charge balance: The total positive and negative charges are balanced in the compound.

Naming Ionic Compounds

• The cation is named first, followed by the anion with the suffix -ide.

• For transition metals (d-block), the charge is specified in parentheses using Roman numerals (e.g., CuCl2 is copper(II) chloride).

• Practice: Name the following compounds:

  • KCl: Potassium chloride

  • Na3N: Sodium nitride

  • FeCl3: Iron(III) chloride

Lattice Energy

Lattice energy is the energy released when one mole of an ionic solid is formed from its gaseous ions. It is a measure of the strength of the ionic bond.

• Higher lattice energy indicates a stronger ionic bond.

The lattice energy can be estimated using Coulomb's Law:

$E = \dfrac{C q^+ q^-}{R_0}$

• C: Constant depending on lattice type

• q+, q-: Charges on ions

• R0: Distance between ions

Trends in Lattice Energy

• For ions with the same charge, lattice energy decreases (becomes less negative) as ionic radius increases.

• For ions with similar radii, lattice energy increases (becomes more negative) as ionic charge increases.

• Charge is typically the more important factor in determining lattice energy.

Metal Chloride	Lattice Energy (kJ mol-1)
LiCl	-834
NaCl	-788
KCl	-701
RbCl	-680
CsCl	-657

Compound	Lattice Energy (kJ mol-1)
NaF	-910
CaO	-3414

Representing Valence Electrons

Valence Electrons and Covalent Bonding

Covalent bonds involve the sharing of valence electrons between two atoms, usually non-metals. Non-metals have high ionization energies and high electron affinities, making them likely to share electrons rather than lose or gain them.

• Example: H2O (water)

Lewis Symbols and Structures

Lewis symbols represent atoms with their valence electrons shown as dots around the element symbol. Only electrons in the s and p orbitals are included.

• Lewis structures show how atoms are connected, non-bonding electrons, and possible resonance structures.

• Limitations: Do not show 3D shape, overall polarity, or bond strength.

Lewis Structures

Constructing Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the arrangement of valence electrons in a molecule.

• Shared electron pairs (covalent bonds) are represented as lines.

• Non-bonding electrons are shown as dots.

Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by eight electrons (except for hydrogen, which follows the duet rule).

Steps for Writing Lewis Structures

1. Determine the total number of valence electrons (add for anions, subtract for cations).

2. Choose the central atom (usually the least electronegative) and connect attached atoms with single bonds (2 electrons each).

3. Complete the octets on the outer atoms, then the central atom.

4. If the central atom lacks an octet, use lone pairs from outer atoms to form multiple bonds.

5. For ions, enclose the structure in square brackets with the charge indicated.

Electronegativity and Bond Polarity

Electronegativity

Electronegativity (EN) is a measure of how strongly an atom attracts shared electrons in a covalent bond. It is calculated from thermochemical data and ranges from 0 (low) to 4 (high).

• Electronegativity is different from electron affinity, which is the energy change when an atom gains an electron.

Bond Polarity

Atoms do not always share electrons equally. The greater the difference in electronegativity ($\Delta EN$), the more polar the bond.

• Nonpolar covalent bond: Electrons shared equally ($\Delta EN \approx 0$).

• Polar covalent bond: Electrons shared unequally ($0.5 < \Delta EN < 2.0$).

• Ionic bond: Electrons transferred ($\Delta EN > 2.0$).

The atom with higher EN gains a partial negative charge ($\delta^-$), while the other gains a partial positive charge ($\delta^+$).

Covalent Bond Energies, Lengths, and Vibrations

Bond Energy and Length

Bond energy is the energy required to break a bond. Shorter bonds tend to be stronger, and multiple bonds (double, triple) are stronger than single bonds.

• Bond energy is typically measured in kJ/mol.

Vibrations and Infrared Activity

Absorption of infrared light causes molecular vibrations to become excited, leading to larger amplitude vibrations. For a vibrational mode to be IR active, it must cause a change in the dipole moment of the molecule.

• Different bonds vibrate at characteristic frequencies, which can be detected in IR spectroscopy.

• The greenhouse effect is due to absorption and re-emission of infrared energy by CO2 and other greenhouse gases.

Resonance and Formal Charge

Resonance Structures

Some molecules can be represented by more than one valid Lewis structure. These are called resonance structures. The actual structure is an average of the resonance forms, with delocalized electrons.

• Double-headed arrows indicate resonance structures.

• Major resonance structures have the fewest nonzero formal charges, charges as close to zero as possible, and negative charges on the most electronegative atom.

Formal Charge

Formal charge is the hypothetical charge an atom would have if bonding electrons were shared equally. It helps determine the most appropriate Lewis structure.

Formal charge is calculated as:

$\text{Formal charge} = (\text{valence electrons in free atom}) - (\text{number of lone pair electrons}) - \frac{1}{2}(\text{number of bonding electrons})$

• The best Lewis structure has the lowest formal charges and places negative charges on the most electronegative atoms.

Drawing Resonance Structures

• Move only electrons from lone pairs, double/triple bonds, or free radicals.

• Do not move electrons in single bonds.

• Use curved arrow notation to indicate electron movement.

• Octet rule applies to C, N, O, and F; H is limited to two electrons.

Exceptions to the Octet Rule

Types of Exceptions

• Free radicals: Molecules with an odd number of electrons, often containing nitrogen (e.g., NO).

• Incomplete octets: Molecules with less than eight electrons on the central atom, typically group 2 and 13 elements (e.g., BF3).

• Hypercoordinate molecules: Molecules with more than eight electrons on the central atom, possible for elements in period 3 and below due to available d-orbitals (e.g., PF5).

When drawing Lewis structures for these molecules, minimize formal charges and follow the rules for resonance structures.

*Additional info: Some examples and explanations have been expanded for clarity and completeness, including the use of tables and formulas for lattice energy and formal charge calculations.*