Chemical Bonding I: Lewis Theory – Study Notes

Chemical Bonding I: Lewis Theory

Chapter Outline

• Bonding Models and Applications

• Types of Chemical Bonds

• Representing Valence Electrons with Dots

• Lewis Structures: Introduction to Ionic & Covalent Bonding

• The Ionic Bonding Model

• Covalent Bond Energies, Lengths, and Vibrations

• Electronegativity and Bond Polarity

• Resonance and Formal Charge

• Exceptions to the Octet Rule

• Lewis Structures for Hypercoordinate Compounds

Types of Chemical Bonds

Formation of a Chemical Bond

Chemical bonds form when atoms interact in such a way that their total energy is lowered. This typically involves the attraction between nuclei and electrons, balanced by repulsions between like charges.

• Step 1: Atoms approach each other.

• Step 2: Interactions occur:

  • Nucleus of one atom attracts the electrons of the other.

  • Electrons of each atom repel those of the other.

  • Nuclei of both atoms repel each other.

• Step 3: If the net effect is a reduction in energy, a chemical bond forms.

Types of Chemical Bonds

• Ionic Bond: Occurs between metals and nonmetals. Electrons are transferred from the metal (low ionization energy) to the nonmetal (high electron affinity), forming cations and anions. The resulting electrostatic attraction is described by Coulomb's law.

• Covalent Bond: Occurs between nonmetals. Electrons are shared between atoms, and the shared electrons are attracted by both nuclei.

• Metallic Bond: Occurs between metals. Valence electrons are pooled and delocalized over a lattice of metal atoms (electron sea model).

Comparison of Bond Types

Types of Atoms	Type of Bond	Characteristic of Bond
Metal & Non-metal	Ionic	Electrons transferred
Non-metal & Non-metal	Covalent	Electrons shared
Metal & Metal	Metallic	Electrons pooled

Note: Bonds can be pure covalent (electrons shared equally), polar covalent (electrons shared unequally), or ionic (electrons transferred).

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is the ability of an atom to attract electrons to itself in a chemical bond. It generally increases across a period and decreases down a group in the periodic table.

Bond Polarity and Dipole Moment

• Nonpolar Covalent Bond: Electrons are shared equally; no dipole moment.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (δ+ and δ−) and a dipole moment.

• Ionic Bond: Electrons are transferred, resulting in full charges and a large dipole moment.

The dipole moment (μ) is calculated as:

$\mu = q r$

where $q$ is the magnitude of the charge and $r$ is the distance between charges.

Classifying Bonds by 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

Representing Valence Electrons with Dots (Lewis Symbols)

Valence Electrons

• For main-group elements, valence electrons are those in the shell of highest principal quantum number ($n$).

• For main-group elements, valence electrons are in the $s$ and $p$ orbitals of the highest $n$.

• There is one $s$ orbital and three $p$ orbitals in the highest shell, for a total of 8 electrons (the octet rule).

Lewis Symbols

Lewis symbols use dots to represent valence electrons around the chemical symbol of an element. For example, oxygen ($O$) with 6 valence electrons is shown as:

O: •• • O • ••

Lewis Structures: Introduction to Ionic & Covalent Bonding

Lewis Structures

Lewis structures depict the arrangement of valence electrons among atoms in a molecule. They show bonding pairs (shared electrons) and lone pairs (non-bonding electrons).

• Hydrogen Duet: Hydrogen achieves stability with 2 electrons.

• Octet Rule: Most main-group elements achieve stability with 8 valence electrons.

Steps for Drawing Lewis Structures

1. Count the total number of valence electrons.

2. Distribute electrons as bonds and lone pairs.

3. Complete octets (or duets for H) by moving lone pairs into bonds if necessary.

4. Check for formal charges to find the best structure.

Example: For $O_2$:

1. Total valence electrons: 12

2. Distribute as bonds and lone pairs

3. Complete octets

4. Check formal charges

The Ionic Bonding Model

Formation of Ionic Compounds

• Metals lose electrons to form cations; nonmetals gain electrons to form anions.

• The resulting ions are held together by electrostatic attraction (Coulomb's law).

Lattice Energy

Lattice energy is the energy released when gaseous ions form an ionic solid. It is a measure of the strength of the ionic bond and is governed by Coulomb's law:

$E = k \frac{Q_1 Q_2}{r}$

where $Q_1$ and $Q_2$ are the charges on the ions, $r$ is the distance between ion centers, and $k$ is a proportionality constant.

Born-Haber Cycle

The Born-Haber cycle is a thermochemical cycle used to calculate lattice energy by considering all steps in the formation of an ionic compound from its elements.

• Enthalpy of sublimation

• Bond dissociation energy

• Ionization energy

• Electron affinity

• Lattice energy

Example: For NaCl, the sum of these steps gives the lattice energy.

Trends in Lattice Energy

• Lattice energy becomes less exothermic (less negative) with increasing ionic radius.

• Lattice energy becomes more exothermic (more negative) with increasing magnitude of ionic charge.

Compound	Lattice Energy (kJ mol-1)
LiCl	-834
NaCl	-788
KCl	-701
CsCl	-657

Covalent Bond Energies, Lengths, and Vibrations

Bond Energy

Bond energy is the energy required to break 1 mole of a bond in the gas phase. It is a measure of bond strength.

Bond Length and Order

• The higher the bond order, the shorter the bond length and the greater the bond energy.

Bond	Bond Length (pm)	Bond Strength (kJ mol-1)
C≡C	120	837
C=C	134	611
C–C	154	347

Estimating Enthalpy Changes Using Bond Energies

The enthalpy change for a reaction can be estimated using average bond energies:

$\Delta H = \sum (\Delta H_{\text{bonds broken}}) + \sum (\Delta H_{\text{bonds formed}})$

• Bonds broken: positive values (energy input)

• Bonds formed: negative values (energy released)

Resonance and Formal Charge

Resonance Structures

Some molecules cannot be adequately represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for the same molecule, differing only in the arrangement of electrons.

Formal Charge

Formal charge is a bookkeeping tool to determine the most stable Lewis structure. It is calculated as:

$\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2} \times \text{Bonding electrons})$

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

Summary Table: Key Bonding Concepts

Concept	Definition/Key Point
Octet Rule	Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons
Ionic Bond	Electrons transferred from metal to nonmetal
Covalent Bond	Electrons shared between nonmetals
Metallic Bond	Electrons pooled among metal atoms
Electronegativity	Ability of an atom to attract electrons in a bond
Bond Energy	Energy required to break 1 mole of a bond in the gas phase
Lattice Energy	Energy released when gaseous ions form an ionic solid
Resonance	Multiple valid Lewis structures for a molecule
Formal Charge	Valence electrons minus nonbonding electrons minus half bonding electrons

Additional info: These notes cover the foundational aspects of Lewis theory, types of chemical bonds, and the energetics and representations of bonding, as outlined in a typical General Chemistry curriculum.