Chemical Bonding I: Lewis Theory – Comprehensive Study Notes

Chemical Bonding I: Lewis Theory

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. The three main types of chemical bonds are ionic, covalent, and metallic bonds.

• Ionic bonds: Formed when electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other through electrostatic forces.

• Covalent bonds: Formed when two or more electrons are shared between atoms, creating a stable electron configuration for each atom.

• Metallic bonds: Involve delocalized valence electrons that are shared among a lattice of metal atoms, resulting in a 'sea of electrons' that gives metals their characteristic properties.

Example: Sodium chloride (NaCl) forms via ionic bonding, while water (H2O) forms via covalent bonding.

Representing Valence Electrons with Dots: Lewis Symbols and Structures

Lewis symbols and structures are visual representations of valence electrons in atoms and molecules. They help predict bonding patterns and molecular structure.

• Lewis symbol: Shows the element's symbol surrounded by dots representing valence electrons.

• Lewis structure: Depicts how atoms share or transfer electrons to achieve stable configurations, showing bonds and lone pairs.

Example: The Lewis structure for water (H2O) shows two single bonds and two lone pairs on oxygen.

Lewis Structures: Introduction to Ionic & Covalent Bonding

Lewis structures illustrate the arrangement of electrons in molecules, distinguishing between bonding pairs (shared electrons) and lone pairs (unshared electrons).

• Bonding pair: Electrons shared between two atoms, forming a covalent bond.

• Lone pair: Electrons not involved in bonding, localized on a single atom.

• Coordinate covalent bond: A covalent bond in which one atom provides both electrons for the shared pair.

Example: In the ammonium ion (NH4+), a coordinate covalent bond forms when NH3 donates a lone pair to H+.

Multiple Covalent Bonds

Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs.

• Single bond: One pair of shared electrons (e.g., H2, F2).

• Double bond: Two pairs of shared electrons (e.g., O2, CO2).

• Triple bond: Three pairs of shared electrons (e.g., N2).

Example: The Lewis structure for N2 shows a triple bond between two nitrogen atoms.

Electronegativity and Bond Polarity

Electronegativity (EN) is a measure of an atom's ability to attract shared electrons in a chemical bond. Differences in EN between bonded atoms lead to bond polarity.

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges.

• Dipole moment: Quantifies the polarity of a bond; calculated as $\mu = q \times d$ where $q$ is the magnitude of charge and $d$ is the distance between charges.

Example: In HCl, chlorine is more electronegative than hydrogen, resulting in a polar covalent bond.

Electronegativity Difference ($\Delta EN$)

The absolute difference in EN values of bonded atoms determines bond type:

Bond Type	$\Delta EN$ Range	Example
Nonpolar covalent	0 – 0.4	Cl2
Polar covalent	0.4 – 2.0	HCl
Ionic	> 2.0	NaCl

Electrostatic Potential Maps

Electrostatic potential maps visualize the charge distribution within a molecule, showing regions of high (red) and low (blue) electron density.

• Helps predict molecular reactivity and polarity.

Writing Lewis Structures

Lewis structures are constructed by following systematic rules to represent the arrangement of valence electrons in molecules and ions.

• Count the total number of valence electrons (use the periodic table).

• Add or subtract electrons for ions (add for negative, subtract for positive).

• Predict the arrangement of atoms (central atom is usually the least electronegative).

• Place electron pairs between atoms to form bonds.

• Distribute remaining electrons as lone pairs to satisfy the octet rule.

Example: For H2O, oxygen is the central atom, with two single bonds to hydrogen and two lone pairs on oxygen.

The Ionic Bonding Model and Lattice Energy

Ionic compounds are held together by electrostatic forces between cations and anions. Lattice energy quantifies the strength of these interactions.

• Lattice energy ($U_{lattice}$): The energy required to separate one mole of an ionic solid into its gaseous ions.

• Calculated using the formula: $U_{lattice} \propto \frac{\text{charge of cation} \times \text{charge of anion}}{\text{cation radius} + \text{anion radius}}$

• Lattice energy increases with higher ionic charges and decreases with larger ionic radii.

Compound	Cation Charge	Anion Charge	Lattice Energy (kJ/mol)
NaCl	+1	-1	-786
MgO	+2	-2	-3795
CaO	+2	-2	-3414

Example: MgO has a higher lattice energy than NaCl due to higher ionic charges.

Covalent Bond Energies, Lengths, and Vibrations

Covalent bond energy is the energy required to break a bond, while bond length is the distance between nuclei of bonded atoms. Bonds can vibrate, and these vibrations can be detected using IR spectroscopy.

• Bond energy (BE): Quantity of energy required to break one mole of a covalent bond in a gaseous species.

• Bond length: Approximately equal to the sum of the covalent radii of the bonded atoms.

• Bond vibrations: Bonds stretch and contract, and their frequencies are measured in wavenumbers (cm-1).

Example: The O–H bond in water has a bond energy of about 463 kJ/mol and a bond length of 0.96 Å.

Bond Order and Bond Lengths

Bond order refers to the number of electron pairs shared between two atoms. Higher bond order generally means shorter bond length and greater bond strength.

• Single bond: Bond order = 1

• Double bond: Bond order = 2

• Triple bond: Bond order = 3

Example: In N2, the triple bond results in a very short and strong bond.

Resonance and Formal Charge

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. The actual structure is an average of all resonance forms. Formal charge helps determine the most reasonable Lewis structure.

• Resonance hybrid: The true structure is a blend of all resonance forms.

• Formal charge: Calculated as: $\text{Formal charge} = \text{valence electrons} - (\text{lone pair electrons} + \frac{1}{2} \times \text{bonding electrons})$

• Structures with the smallest formal charges and negative charges on the most electronegative atoms are preferred.

Example: The nitrate ion (NO3-) has three resonance forms, each with different formal charges distributed among the oxygen atoms.

Exceptions to the Octet Rule

Some molecules and ions do not follow the octet rule due to the presence of odd electrons, electron-deficient atoms, or expanded valence shells.

• Odd-electron species: Molecules with an odd number of electrons (e.g., NO, NO2).

• Electron-deficient atoms: Elements like B and Be often have fewer than 8 electrons.

• Expanded valence shells: Nonmetals in the third period and beyond can accommodate more than 8 electrons (e.g., SF6, XeF4).

Example: Phosphorus pentachloride (PCl5) has 10 electrons around phosphorus.

Lewis Structures for Hypercoordinate Compounds

Elements in the third period and beyond can form compounds with more than eight electrons in their valence shell, known as hypercoordinate compounds.

• Atoms can be bonded to more than four atoms or have multiple bonds to four or fewer atoms.

Example: The Lewis structure for XeF4 shows xenon bonded to four fluorine atoms with two lone pairs.

Summary Table: Key Concepts in Lewis Theory

Concept	Definition	Example
Ionic Bond	Transfer of electrons	NaCl
Covalent Bond	Sharing of electrons	H2O
Metallic Bond	Delocalized electrons	Cu
Resonance	Multiple valid Lewis structures	NO3-
Formal Charge	Electron accounting for atoms	CO2
Octet Rule Exception	Odd electrons, electron-deficient, expanded shells	BF3, SF6

Additional info:

• Bond energies and enthalpy changes can be estimated using the formula: $\Delta H = \sum BE(\text{reactants}) - \sum BE(\text{products})$

• Infrared (IR) spectroscopy is used to detect bond vibrations, with absorption frequencies reported in wavenumbers (cm-1).

• Lewis structures are essential for predicting molecular geometry, reactivity, and physical properties.