Chemical Bonding I: The Lewis Model – Study Notes

Chemical Bonding I: The Lewis Model

Introduction to Bonding Theories

Chemical bonding theories explain how and why atoms attach together to form molecules, and why some combinations of atoms are stable while others are not. These theories allow us to predict molecular shapes, as well as the chemical and physical properties of compounds.

• Bonding theories help explain molecular stability and predict molecular properties.

• They are essential for understanding why water is H2O and not HO or H3O.

• Applications include drug design, such as HIV-protease inhibitors.

The Lewis Model

The Lewis model is one of the simplest bonding theories, emphasizing the role of valence electrons in bonding. In this model, valence electrons are represented as dots around the element symbol, forming what are known as Lewis structures or electron dot structures.

• Lewis structures help predict molecular stability and appearance.

• They focus on the behavior of valence electrons, which are most important in bonding.

Why Do Atoms Bond?

Atoms form chemical bonds to lower their potential energy. A bond forms when the potential energy of the bonded atoms is less than that of the separate atoms. The interactions to consider include:

• Nucleus-to-nucleus repulsions

• Electron-to-electron repulsions

• Nucleus-to-electron attractions

Types of Chemical Bonds

Classification of Bonds

Bonds are classified based on the types of atoms involved:

Bond Type	Example	Description
Ionic	NaCl	Transfer of electrons from metal to nonmetal, forming cations and anions
Covalent	H2O	Sharing of electrons between nonmetals
Metallic	Na (metal)	Delocalized electrons shared among metal atoms

Ionic Bonds

Formed when a metal atom loses electrons (becoming a cation) and a nonmetal atom gains electrons (becoming an anion). The resulting oppositely charged ions are attracted to each other.

• Metals have low ionization energy; nonmetals have high electron affinity.

• Example: NaCl (table salt)

Covalent Bonds

Formed when nonmetal atoms share valence electrons. The shared electrons are attracted by the nuclei of both atoms, resulting in a stable bond.

• Nonmetals have high ionization energies, making electron sharing favorable.

• Example: H2O (water)

Metallic Bonds

In metals, valence electrons are delocalized and shared among all atoms, forming a "sea of electrons" that holds the metal cations together.

• Explains properties like conductivity and malleability in metals.

Valence Electrons and Lewis Structures

Valence Electrons

Valence electrons are the outermost electrons and are most involved in chemical bonding. The main-group column number on the periodic table indicates the number of valence electrons for main-group elements.

• Transition elements typically have two valence electrons.

Lewis Structures of Atoms

Lewis structures represent valence electrons as dots around the element symbol. The first two dots are paired for s orbital electrons, and the next three are placed singly for p electrons before pairing the rest.

Octet Rule

Atoms tend to bond in such a way that they each obtain an outer shell with eight electrons, similar to noble gases. This is known as the octet rule. There are exceptions, such as hydrogen (duet rule), boron, and elements with expanded octets.

Ionic Bonding and Lattice Energy

Lewis Theory and Ionic Bonding

Lewis symbols can represent the transfer of electrons from a metal to a nonmetal, forming ions that are attracted to each other and bond.

Crystal Lattice and Lattice Energy

Ionic compounds form a crystal lattice, a repeating pattern of cations and anions. The lattice is held together by electrostatic attraction, maximizing stability. The energy released when the lattice forms is called lattice energy.

Born–Haber Cycle

The Born–Haber cycle is a hypothetical series of steps used to calculate the lattice energy of an ionic compound using Hess's law. It involves summing enthalpy changes for each step in the formation of the compound from its elements.

Trends in Lattice Energy

• Lattice energy decreases (becomes less exothermic) as ion size increases.

• Lattice energy increases (becomes more exothermic) as ion charge increases.

• Ion charge is generally more important than ion size.

Properties of Ionic Compounds

Physical Properties

• High melting and boiling points (generally > 300 °C)

• Hard and brittle crystalline solids

• Conduct electricity in the liquid state or when dissolved in water, but not as solids

• Many are soluble in water

Electrical Conductivity

Ionic solids do not conduct electricity because ions are locked in place. When melted or dissolved, ions are free to move and conduct electricity.

Covalent Bonding and Lewis Structures

Covalent Bonds

Covalent bonds involve the sharing of electron pairs between atoms. Shared electrons are called bonding pairs, while non-shared electrons are lone pairs or nonbonding pairs.

• Single bond: one pair of shared electrons

• Double bond: two pairs of shared electrons

• Triple bond: three pairs of shared electrons

Lewis Structures for Molecules

To draw a Lewis structure:

1. Write the correct skeletal structure (hydrogen and more electronegative atoms are terminal).

2. Sum the valence electrons for all atoms.

3. Distribute electrons to give octets (or duets for hydrogen).

4. If atoms lack octets, form double or triple bonds as needed.

Resonance

Some molecules have more than one valid Lewis structure, differing only in the position of electrons. These are called resonance structures. The actual molecule is a resonance hybrid, a combination of all resonance forms.

Formal Charge

Formal charge is a tool to evaluate the best Lewis structure. It is calculated as:

$\text{Formal charge} = \text{valence electrons} - [\text{nonbonding electrons} + \frac{1}{2} \text{bonding electrons}]$

• The sum of all formal charges in a molecule must be zero (or equal to the ion charge).

• Small or zero formal charges are preferred.

• Negative formal charge should reside on the most electronegative atom.

Exceptions to the Octet Rule

• Odd-electron species (free radicals)

• Incomplete octets (e.g., B, Al)

• Expanded octets (elements with empty d orbitals)

Bond Polarity and Electronegativity

Electronegativity

Electronegativity is the ability of an atom to attract bonding electrons. It increases across a period and decreases down a group. The difference in electronegativity between atoms determines bond polarity:

• 0: pure covalent (nonpolar)

• 0.1–0.4: nonpolar covalent

• 0.4–1.9: polar covalent

• ≥2.0: ionic

Bond Dipole Moments

A dipole moment ($\mu$) measures bond polarity and is calculated as:

$\mu = q \times r$

where $q$ is the magnitude of the charge and $r$ is the distance between charges. Measured in Debyes (D).

Percent Ionic Character

The percent ionic character compares the measured dipole moment to the value if electrons were completely transferred. It indicates the degree of electron transfer in a bond.

Bond Energies and Bond Lengths

Bond Energy

Bond energy (or bond enthalpy) is the energy required to break one mole of a bond in the gas phase. It is always positive (endothermic). The overall enthalpy change for a reaction can be estimated using average bond energies:

$\Delta H_{\text{rxn}} = \sum \text{(bond energies of bonds broken)} - \sum \text{(bond energies of bonds formed)}$

• Bond breaking is endothermic (+), bond making is exothermic (−).

Bond Length

Bond length is the distance between the nuclei of bonded atoms. In general:

• More shared electrons = shorter bond

• Bond length decreases across a period and increases down a group

• Shorter bonds are stronger

Summary Table: Types of Bonds and Properties

Bond Type	Formation	Properties
Ionic	Transfer of electrons	High melting/boiling points, conduct electricity when molten or dissolved, hard and brittle
Covalent	Sharing of electrons	Low melting/boiling points, do not conduct electricity, can be soft or hard
Metallic	Delocalized electrons	Conduct electricity, malleable, ductile