Chemical Bonding, Molecular Structure, and Intermolecular Forces

Chemical Bonding and Lewis Structures

Types of Chemical Bonds

Chemical bonds are the attractive forces that hold atoms together in compounds. The three main types are:

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions.

• Covalent Bonds: Formed by the sharing of electrons between two nonmetals.

• Metallic Bonds: Involve a 'sea' of delocalized electrons shared among metal atoms.

Lewis Dot Structures

Lewis structures represent the valence electrons of atoms and show how atoms are bonded in molecules and ions.

• Atoms and Monatomic Ions: Dots represent valence electrons around the element symbol. For ions, add or remove electrons according to the charge.

• Polyatomic Molecules and Ions: Arrange atoms, connect with single bonds, distribute remaining electrons to satisfy the octet rule, and adjust for multiple bonds if needed.

Exceptions to the Octet Rule

• Odd-Electron Species: Molecules with an odd number of electrons (e.g., NO) cannot have all atoms with octets.

• Incomplete Octets: Some elements (e.g., Be, B) are stable with fewer than 8 electrons.

• Expanded Octets: Elements in period 3 or higher (e.g., P, S, Xe) can have more than 8 electrons.

Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula:

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

• Structures with formal charges closest to zero are generally preferred.

Resonance

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

Bond Polarity and Electronegativity

• Bond Polarity: A bond is polar if electrons are shared unequally due to differences in electronegativity.

• Electronegativity: The ability of an atom to attract shared electrons. Increases across a period and decreases down a group.

Bond Order

• Definition: The number of shared electron pairs between two atoms.

• Calculation: For resonance structures, bond order = (total number of bonds)/(number of bond locations).

Estimating Reaction Enthalpy from Bond Energies

• Bond energy is the energy required to break one mole of a bond in the gas phase.

• Formula:

$\Delta H_{\text{rxn}} = \sum \text{(Bond energies of bonds broken)} - \sum \text{(Bond energies of bonds formed)}$

Molecular Geometry and Bonding Theories

VSEPR Theory (Valence Shell Electron Pair Repulsion)

VSEPR theory predicts the shape of molecules based on the repulsion between electron groups around a central atom.

• Steric Number: The total number of electron groups (bonding and lone pairs) around the central atom.

• Electron Geometry: The arrangement of all electron groups.

• Molecular Geometry: The arrangement of only the atoms (ignoring lone pairs).

Common Geometries for 2–5 Electron Groups

Steric Number	Electron Geometry	Molecular Geometry (examples)	Bond Angles
2	Linear	Linear	180°
3	Trigonal planar	Trigonal planar, Bent	120°
4	Tetrahedral	Tetrahedral, Trigonal pyramidal, Bent	109.5°
5	Trigonal bipyramidal	Trigonal bipyramidal, Seesaw, T-shaped, Linear	90°, 120°, 180°

Additional info: Group 6 geometries (octahedral) are typically provided separately.

Bond Angles

• Bond angles decrease as lone pairs increase due to greater repulsion.

Polarity of Molecules

• Diatomic Molecules: Polar if the two atoms have different electronegativities.

• Polyatomic Molecules: Polarity depends on both bond polarity and molecular geometry. If dipoles cancel, the molecule is nonpolar.

Valence Bond Theory

Describes covalent bonding as the overlap of atomic orbitals, forming sigma (σ) and pi (π) bonds.

• Sigma (σ) Bonds: Formed by head-on overlap; all single bonds are sigma bonds.

• Pi (π) Bonds: Formed by side-on overlap; present in double and triple bonds.

Identifying and Counting Sigma/Pi Bonds

• Single bond: 1 σ

• Double bond: 1 σ + 1 π

• Triple bond: 1 σ + 2 π

Hybridization

• Atomic orbitals mix to form hybrid orbitals for bonding.

• Types: sp (2 groups), sp2 (3 groups), sp3 (4 groups), sp3d (5 groups), sp3d2 (6 groups).

Molecular Orbital (MO) Theory

Describes bonding using molecular orbitals formed from the combination of atomic orbitals.

• MO Configuration: Fill molecular orbitals according to energy levels and the number of electrons.

• Homonuclear Molecules and Ions: MO diagrams are especially important for diatomic molecules of the same element (e.g., O2, N2).

• Bond Order:

$\text{Bond order} = \frac{1}{2} (\text{number of bonding electrons} - \text{number of antibonding electrons})$

Intermolecular Forces and Properties of Liquids and Solids

Intermolecular Forces (IMFs)

IMFs are forces of attraction between molecules, influencing physical properties.

• Types:

  • Dispersion (London) Forces: Present in all molecules; strongest in large, polarizable molecules.

  • Dipole-Dipole Forces: Between polar molecules.

  • Hydrogen Bonding: Strong dipole-dipole interaction when H is bonded to N, O, or F.

  • Ion-Dipole Forces: Between ions and polar molecules (important in solutions).

Vaporization and Vapor Pressure

• Vaporization: The process by which molecules escape from the liquid phase to the gas phase.

• Vapor Pressure: The pressure exerted by a vapor in equilibrium with its liquid at a given temperature.

• Stronger IMFs result in lower vapor prebssure and higher boiling points.

Sublimation and Fusion

• Sublimation: Direct transition from solid to gas (e.g., dry ice).

• Fusion (Melting): Transition from solid to liquid.

Phase Diagrams

Phase diagrams show the state of a substance as a function of temperature and pressure.

• Regions: Solid, liquid, and gas phases.

• Lines: Boundaries between phases (equilibrium lines).

• Triple Point: All three phases coexist in equilibrium.

• Critical Point: The end point of the liquid-gas boundary; above this, the substance is a supercritical fluid.