Chemical Bonding I: Lewis Theory

TYPES OF CHEMICAL BONDS

Lewis Theory (Section 9-2)
- Valence electrons play a crucial role in chemical bonding.
- Electrons can be transferred from one atom to another, forming positive and negative ions.
- Ions attract each other through electrostatic forces, forming ionic bonds.
- One or more electrons can be shared between atoms, resulting in covalent bonds.
- Electrons are shared such that each atom acquires a stable electron configuration, often achieving an octet.
- Metallic bonding involves delocalized valence electrons throughout the metal, leading to a 'sea of electrons' that attracts positively charged metal ions.

REPRESENTING VALENCE ELECTRONS WITH DOTS

Lewis Symbols and Structures

Lewis Symbol: A representation of an atom's valence electrons.
Lewis Structure: A depiction of valence electrons' arrangement among atoms in a molecule.

LEWIS STRUCTURES: AN INTRODUCTION TO IONIC & COVALENT BONDING

Covalent Bonds (Section 9-4)

Bonding or shared electron pair: Electrons involved in a covalent bond.
Lone or unshared electron pair: Electrons not involved in bonding.
Lone Pair of Electrons: Non-bonding valence electrons in an atom.
Bond Pair of Electrons: Electrons that are shared between atoms in a bond.
Coordinate Covalent Bonds: A covalent bond where one atom contributes both electrons to a shared pair.

Multiple Covalent Bonds

Single Bond: Consists of one bonding pair of electrons (e.g., H2, HF, F2).
Double Bond: Comprises two bonding pairs of electrons (e.g., CO2), sharing four electrons between two atoms.
Triple Bond: Involves three bonding pairs (e.g., N2), sharing six electrons between two atoms.
Multiple bonds typically occur when C, O, N, or S atoms are bonded with each other.

ELECTRONEGATIVITY AND BOND POLARITY

Electronegativity (χ or EN)

Definition: The ability of an atom to attract shared electrons in a bond.
In HCl, the electrons are shared unequally, making it a polar covalent bond.
Dipole Moment (µ): A measure of the separation of positive and negative charges in a bond: ext{µ} = q imes r
- Where $q$ is the magnitude of charge and $r$ is the distance between the charges.
Electronegativity Difference (ΔEN):
- Small (0 - 0.4): Covalent (e.g., Cl2)
- Intermediate (0.4 - 2.0): Polar covalent (e.g., HCl)
- Large (2.0+): Ionic (e.g., NaCl)
- ΔEN allows for classification of bonds.
Electrostatic Potential Map: A visualization of charge distribution within a molecule, showing ranging densities from high (red) to low (blue).

WRITING LEWIS STRUCTURES

Fundamental Concepts

Lewis Structure: A way to represent electron pairs in a molecule.
- Octet Rule: Atoms tend to bond to achieve a filled outer shell of 8 electrons.
- Duet Rule: Hydrogen atoms should have 2 electrons.
Skeletal Structure: Arrangement of atoms reflecting actual spatial distribution.
- Central Atom: The atom bonded to two or more other atoms.
- Terminal Atom: An atom bonded to only one atom.

Rules for Writing Lewis Structures

Count Valence Electrons:
- For molecules, sum the valence electrons.
- For ions, add charge for negative ions and subtract for positive ions.
Predict Atom Arrangement:
- H is always a terminal atom.
- O and halogens are usually terminal unless bound to another halogen.
- C typically acts as a central atom.
Draw Sigma Bonds:
- Place two electrons to represent bonds.
Distribute Electrons:
- Assign lone pairs to terminal atoms to satisfy their octet (or duet for H).
- If electrons remain, assign them to a central atom with available electron pairs.
- If not eight electrons surround the central atom, form pi bonds by converting lone pairs from terminal atoms.
- Only C, N, O, and S can form pi bonds with each other.

Example Structures

H2O:
- Skeletal structure: H-O-H.
- Sum of valence electrons = 8.
- Bonds drawn: H-O (4 electrons used). Remaining = 4, assigned as lone pairs to O.
CO2:
- Predict arrangement: C (lowest EN) central, O (terminal).
- Valence electrons = 16; draw O=C=O bonds. Remaining electrons = 12; assign as lone pairs.

THE IONIC BONDING MODEL

Lattice Energy

Definition: The energy associated with forming ionic compounds from gaseous ions.
Lattice energy is proportional to the product of the charges and inversely proportional to the sum of the radii of cations and anions: ext{lattice energy} ext{ (ion)} imes k rac{qA imes qB}{( ext{radius of cation} + ext{radius of anion})}
As ionic size increases, lattice energy decreases.
- Example Lattice Energies for Metal Chlorides:
- LiCl: -834 kJ/mol
- NaCl: -788 kJ/mol
- KCl: -701 kJ/mol
- CsCl: -657 kJ/mol
- As ionic charge increases, lattice energy increases.

Born-Haber Cycle

A thermodynamic cycle used to analyze the stages in an ionic solid's formation.
Example Calculation for LiF:
- Sublimation of lithium: ΔH°_{ ext{sublimation}} = 161 ext{ kJ}
- Ionization energy of Li: IE_{ ext{Li}} = 520 ext{ kJ}
- Bond energy for F2: E_{dissociation} = 79.5 ext{ kJ}
- Electron affinity of F: EA = -328 ext{ kJ}
- Lattice energy of LiF: ΔH°_{ ext{lattice}} = -1049.5 ext{ kJ}
- Overall energy change: ΔH°{overall} = ΔH°{f} = -617 ext{ kJ}

COVALENT BOND ENERGIES, LENGTHS, AND VIBRATIONS

Bond Energy

Definition: The energy required to break one mole of covalent bonds in gaseous species.
Bond breakage leads to homolytic cleavage, an endothermic process (positive energy change).
Bond formation is exothermic (negative energy change).

Average Bond Energy

Average of bond dissociation energies for various related species.

Estimation of Enthalpy Changes

Estimated as ΔH ≈ RBE{reactants} - RBE{products}
For example, using bond energies to calculate enthalpy of combustion for methanol.

BOND ORDER & BOND LENGTHS

Bond Order

The number of electron pairs shared between two bonded atoms:
- Single bonds: bond order = 1
- Double bonds: bond order = 2
- Triple bonds: bond order = 3

Bond Length

The distance between the nuclei of two bonded atoms.
Covalent bond length approximately equals the sum of the covalent radii of the bonded atoms.

BOND VIBRATIONS

Concept: Bonds are dynamic and can stretch or contract.
Vibrations require energy input and can be studied using infrared (IR) spectroscopy.
IR spectra units are in wavenumbers (cm-1).

RESONANCE & FORMAL CHARGE

Resonance

More than one Lewis structure can describe the electron distribution in a molecule.
The actual structure is an average of resonance structures, connected by double-headed arrows (↔).
Example:
- Nitrate ion (NO3-):
- Structure depicting resonance where the bonds are distributed among the N and O atoms.

Formal Charge

Definition: Given by the equation:
ext{Formal Charge} = ext{(number of valence electrons)} - ext{(assigned valence electrons)}
General rules for assessing Lewis Structures favoring lower formal charges and proper distributions based on electronegativity

EXCEPTIONS TO THE OCTET RULE

Molecules with Odd Electron Atoms:
- Free radicals possess unpaired electrons and can be particularly reactive.
- Examples include NO and ClO2.
Molecules with Electron Deficient Atoms:
- Elements like B and Be may have fewer than 8 electrons in their covalent compounds.
Molecules with Expanded Valence Shells:
- Non-metals in the third row or higher can utilize d orbitals for bonding.

Lewis Structures for Hypercoordinate Compounds

Examples include xenon tetrafluoride (XeF4) and phosphoric acid (H3PO4).

OBJECTIVES

Students should be able to:
- Differentiate between ionic, covalent, and metallic bonds.
- Write Lewis symbols and structures.
- Describe covalent bond formation and properties.
- Calculate lattice energies and apply the Born-Haber cycle.
- Understand bond energies, lengths, and vibrations.
- Explain electronegativity and bond polarity.
- Analyze resonance and formal charge implications in chemical structures.
- Address exceptions to the octet rule and construct valid Lewis structures.
- Review hypercoordinate compounds and their Lewis structures.