Covalent Bonding and Atomic Structure

Fundamentals of Chemical Bonds

Chemical bonds are the forces that hold atoms together in molecules and compounds. The most important interactions involve electrons, which are arranged in orbitals around the nucleus. Orbitals are defined by quantum numbers and are grouped into electron shells. The arrangement of electrons in these shells determines how atoms bond.

• Atom Structure: Atoms consist of a nucleus (protons and neutrons) surrounded by electrons.

• Orbitals: Regions in space where electrons are likely to be found. Types include s, p, d, and f.

• Electron Shells: Each shell can hold a specific number of orbitals and electrons.

• Valence Electrons: Electrons in the outermost s and p orbitals, crucial for bonding.

Electron Configuration

The electron configuration of an atom is written by listing each orbital type with a superscript indicating the number of electrons. For example, lithium is 1s22s1.

• Valence Electron Counting: For s and d block elements, count from the previous noble gas. For p block elements, subtract 10 from the group number.

Types of Chemical Bonds

Bonding Interactions

Within a molecule, three main interactions occur:

• Electrons and nuclei attract each other.

• Electrons repel each other.

• Nuclei repel each other.

Balance among these interactions leads to molecular stability. When electron density is shared between nuclei, a covalent bond forms.

Bond Length and Bond Energy

Bond length is the distance between nuclei where the molecule is most stable. Bond energy is the energy required to break the bond.

• Bond length and energy are characteristic for each bond type.

• Bond length increases as atom size increases.

• Bond energy increases with more shared electrons and greater electronegativity difference.

Covalent Bonding

Formation of Covalent Bonds

Covalent bonds are formed by sharing one or more pairs of electrons between atoms. Atoms achieve a stable noble gas configuration (octet rule) through sharing.

• Single Bond: One pair of electrons shared (e.g., H–H).

• Double Bond: Two pairs shared (e.g., O=O).

• Triple Bond: Three pairs shared (e.g., N≡N).

Polar Covalent Bonds and Electronegativity

Polar Covalent Bonds

When two different atoms share electrons unequally, a polar covalent bond forms. The atom with higher electronegativity attracts electrons more strongly, resulting in partial charges (δ+ and δ–).

Electronegativity

Electronegativity (𝜒) is a measure of an atom's ability to attract electrons in a bond. It increases across a period and decreases down a group.

• If Δ𝜒 > 2.0: Ionic bond

• If Δ𝜒 < 0.5: Covalent bond

• If 0.5 < Δ𝜒 < 2.0: Polar covalent bond

Properties of Covalent Bonds

Dipole Moments

Bond polarities can create a dipole moment (µ), a measure of charge separation across a molecule. The magnitude depends on the electronegativity difference.

Bond Length

Bond length increases with atom size. For example, halogen bond lengths increase from F2 to I2 due to increasing atomic radius.

Bond Energy

• Bond energy increases with more shared electrons and greater electronegativity difference.

• Bond energy decreases as bond length increases.

Polarity of Molecules

Polar and Nonpolar Molecules

A molecule is polar if it has a net dipole moment due to bond polarities and molecular shape. Asymmetric molecules like water and chloroform are polar, while symmetric molecules like CO2 and CCl4 are nonpolar.

Effect of Polarity

• Polarity affects melting points, boiling points, and solubility.

• Water's polarity is responsible for its liquid state at room temperature.

Ionic Bonds and Lattice Energies

Ionic Bonds

Ionic bonds form between elements with large differences in electronegativity. Ionic compounds are typically solids with high melting points, held together by electrostatic attraction between cations and anions.

Lattice Energy

• Lattice energy is the energy required to break an ionic lattice.

• It decreases as ion size increases.

Lewis Structures and Resonance

Lewis Structures

Lewis structures show how atoms are bonded and the locations of lone-pair valence electrons. Only valence electrons are depicted, with lines for bonds and dots for lone pairs.

Drawing Lewis Structures: Five-Step Procedure

1. Count valence electrons in all atoms.

2. Assemble the bonding framework using single bonds.

3. Place three nonbonding electron pairs on each outer atom except H/He.

4. Assign remaining valence electrons to inner atoms.

5. Minimize formal charges on all atoms.

Formal Charge Calculation

• Formal charge = (valence electrons of free atom) – (electrons assigned in structure)

• Lone pair electrons are assigned to the atom; bonding electrons are shared.

Resonance Structures

Resonance occurs when more than one valid Lewis structure exists for a molecule. The true structure is a blend of all resonance forms, with electrons delocalized.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR Theory

VSEPR theory states that molecular shape is determined by repulsions between pairs of valence electrons, both bonding and nonbonding. Electron pairs arrange themselves as far apart as possible to minimize repulsion.

Geometry vs. Shape

• Geometry: Arrangement of all electron sets (bonding and lone pairs) around the central atom.

• Shape: Arrangement of atoms in space, considering only bonding pairs.

• Lone pairs cause deviations from ideal bond angles due to stronger repulsion.

Examples of Molecular Shapes

Linear Geometry (2 sets)

• Bond angle: 180°

• Example: BeH2, CO2

Trigonal Planar Geometry (3 sets)

• Bond angle: 120°

• Example: BF3 (trigonal planar), NO2– (bent)

Tetrahedral Geometry (4 sets)

• Bond angle: 109.5°

• Example: CH4 (tetrahedral), NH3 (trigonal pyramidal), H2O (bent)

Trigonal Bipyramidal Geometry (5 sets)

• Bond angles: 90°, 120°

• Example: PCl5 (trigonal bipyramidal), SF4 (seesaw), ClF3 (T-shaped), I3– (linear)

Octahedral Geometry (6 sets)

• Bond angle: 90°

• Example: SF6 (octahedral), ClF5 (square pyramidal), XeF4 (square planar)

Representing 3D Structures

3D molecular structures are often represented in 2D using wedges: solid wedges for bonds coming out of the page, dashed wedges for bonds going into the page.

Summary Table: VSEPR Geometry and Shape