Chemical Bonding I: Lewis Theory – Structured Study Notes

Ch9: Chemical Bonding I – Lewis Theory

9.2 Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. They form because they lower the potential energy between charged particles in atoms, resulting in a more stable molecule.

• Ionic Bond: Formed by the transfer of electrons from a metal to a non-metal, resulting in positive and negative ions held together by electrostatic attraction.

• Covalent Bond: Formed by the sharing of electrons between non-metal atoms, creating a stable electron configuration for each atom.

• Metallic Bond: Involves a lattice of metal cations surrounded by a sea of delocalized valence electrons, which are not localized to any single atom.

Each bond type results from distinct electron interactions, and understanding the arrangement of valence electrons is key to predicting bond formation.

9.3 Representing Valence Electrons with Dots

Lewis theory uses dot diagrams to represent valence electrons around an element's symbol, helping visualize bonding possibilities.

• Valence Electrons: Electrons in the outermost shell, involved in chemical bonding.

• Lewis Dot Structure: Dots are placed around the element symbol, with a maximum of two per side.

• Octet Rule: Atoms tend to achieve a stable configuration with eight valence electrons (an octet). Helium is an exception, stable with two electrons.

Example: Oxygen (O) has 6 valence electrons, represented as six dots around the symbol.

9.4 Lewis Structures: Ionic and Covalent Bonding

Lewis structures show how atoms share or transfer electrons to achieve stable configurations.

• Single, Double, Triple Bonds: Atoms can share one, two, or three pairs of electrons, forming single, double, or triple bonds, respectively.

• Lone Pairs: Electrons not involved in bonding, which influence molecular shape.

• Polyatomic Ions: Lewis structures for ions include extra electrons and are enclosed in brackets with the charge indicated.

Examples:

• Water (H2O): Oxygen shares electrons with two hydrogens, achieving an octet.

• Chlorine gas (Cl2): Each Cl shares one electron, forming a single bond and three lone pairs.

• Oxygen gas (O2): Each O shares two electrons, forming a double bond and two lone pairs.

• Nitrogen gas (N2): Each N shares three electrons, forming a triple bond and one lone pair.

9.8 Resonance and Formal Charge

Some molecules have multiple valid Lewis structures, called resonance structures. The true structure is a hybrid of these possibilities.

• Resonance: Delocalization of electrons across multiple atoms, as seen in benzene (C6H6) and carbonate ion (CO32−).

• Formal Charge: Hypothetical charge assigned to an atom in a Lewis structure, calculated as:

$\text{FC} = \text{valence electrons} - (\text{non-bonding electrons} + \frac{1}{2} \times \text{bonding electrons})$

• Structures with the lowest formal charges (ideally zero) are preferred.

• Negative formal charges should reside on the most electronegative atoms.

Example Table: Formal Charge Calculation for HCl

	H in HCl	Cl in HCl
# valence e−	1	7
# non-bonding e−	0	6
# bonding e−	2	2
FC	0	0

9.9 Exceptions to the Octet Rule

Some molecules do not follow the octet rule due to odd numbers of electrons, incomplete octets, or expanded octets.

• Odd Electron Species: Molecules with an odd number of electrons (free radicals), e.g., OH•, NO, NO2.

• Incomplete Octets: Some atoms (e.g., boron in BF3) are stable with fewer than eight electrons.

• Expanded Octets: Atoms in period 3 or higher (e.g., S in SO42−, Cl in ClO4−) can have more than eight electrons.

9.10 Hypercoordinate Compounds

Hypercoordinate compounds support expanded octets, often seen in polyatomic ions.

• Example: Sulfate ion (SO42−) and perchlorate ion (ClO4−) have central atoms with more than eight electrons.

• Formal charge calculations help determine the most stable structure.

9.5 The Ionic Bonding Model

The formation of ionic compounds is exothermic, primarily due to lattice energy. The Born-Haber cycle is used to calculate the energetics of ionic compound formation.

• Born-Haber Cycle: A series of steps (ionization, electron affinity, lattice formation) used to determine lattice energy.

• Lattice Energy: Energy released when gaseous ions form an ionic solid. It decreases for larger ions and increases with higher charges.

Example Equation:

$\Delta H_f[\text{NaCl(s)}] = -411 \text{ kJ mol}^{-1}$

9.6 Covalent Bond Energies, Lengths, and Vibrations

Covalent bond energy is the energy required to break one mole of bonds in the gas phase. Bond length and strength are inversely related.

• Bond Energy: Used to estimate reaction enthalpies by summing energies of bonds broken and formed.

• Bond Length: Shorter bonds are generally stronger.

• Bond Vibrations: Bonds vibrate at characteristic frequencies, detectable by infrared (IR) spectroscopy.

Example Equation for Reaction Enthalpy:

$\Delta H^\circ = \sum \text{Energy of bonds broken} - \sum \text{Energy of bonds formed}$

9.7 Electronegativity and Bond Polarity

Electronegativity is the tendency of an atom to attract electrons in a bond. Differences in electronegativity lead to bond polarity.

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

• Electronegativity Trends: Increase across a period and up a group in the periodic table.

• Bonds range from nonpolar covalent (equal sharing) to polar covalent to ionic (complete transfer) depending on electronegativity difference.

Example: In HCl, Cl is more electronegative, so the bond is polar with partial charges.