Chemical Bonding and Molecular Structure: Chapters 8 & 9 Study Guide

Chemical Bonding and Molecular Structure

I. Lattice Energy

Lattice energy is a measure of the strength of the forces holding ions together in an ionic solid. It is defined as the energy required to separate one mole of a solid ionic compound into its gaseous ions.

• Definition: Lattice energy quantifies the stability of an ionic solid.

• Factors Affecting Lattice Energy:

  • Charge of ions: Higher charges increase lattice energy.

  • Size of ions: Smaller ions result in higher lattice energy.

• Equation (Born-Haber Cycle): $\Delta H_{lattice} = \sum \text{(steps in the Born-Haber cycle)}$

• Example: NaCl has a high lattice energy due to the small size and high charge of Na+ and Cl- ions.

II. Born-Haber Cycle

The Born-Haber cycle is a thermochemical cycle used to calculate the lattice energy of an ionic compound using Hess's Law.

• Steps in the Cycle:

  1. Sublimation of the metal

  2. Ionization of the metal atom

  3. Dissociation of the nonmetal molecule

  4. Electron affinity of the nonmetal atom

  5. Formation of the ionic solid

• Equation: $\Delta H_{f} = \Delta H_{sublimation} + IE + \frac{1}{2} D + EA + \Delta H_{lattice}$

• Application: Used to determine the lattice energy indirectly.

III. Ionic vs. Covalent Bond

Chemical bonds can be classified as ionic or covalent based on the nature of electron sharing or transfer between atoms.

• Ionic Bond: Involves the transfer of electrons from a metal to a nonmetal, resulting in the formation of ions.

• Covalent Bond: Involves the sharing of electrons between two nonmetals.

• Comparison Table:

Property	Ionic Bond	Covalent Bond
Electron Movement	Transfer	Sharing
Types of Elements	Metal + Nonmetal	Nonmetal + Nonmetal
Physical State	Solid (usually)	Solid, liquid, or gas
Melting Point	High	Low to moderate

IV. Bond Polarity (Types of Bonds)

Bond polarity refers to the distribution of electron density between two atoms in a bond.

• Nonpolar Covalent: Electrons are shared equally (e.g., H2).

• Polar Covalent: Electrons are shared unequally due to differences in electronegativity (e.g., HCl).

• Ionic: Electrons are transferred (e.g., NaCl).

V. Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.

• Trend: Increases across a period (left to right), decreases down a group.

• Most Electronegative Element: Fluorine (F).

• Difference in Electronegativity: Determines bond type (nonpolar, polar, ionic).

VI. Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

• Steps to Draw:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to satisfy the octet rule.

  4. Use double/triple bonds if necessary.

• Example: Lewis structure of CO2 shows double bonds between C and O.

VII. Formal Charge

Formal charge is a bookkeeping tool to determine the most stable Lewis structure.

• Formula: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

• Lowest formal charges are preferred for stability.

VIII. Octet Rule: Incomplete Octet, Expanded, and Odd Number of Electrons

The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons. There are exceptions:

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

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons (e.g., SF6).

• Odd Number of Electrons: Molecules with an odd number of electrons (e.g., NO) cannot satisfy the octet rule for all atoms.

IX. Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule, determined by the number of bonding and lone pairs around the central atom.

• Common Geometries:

  • Linear (180°)

  • Trigonal planar (120°)

  • Tetrahedral (109.5°)

  • Trigonal bipyramidal (90°, 120°)

  • Octahedral (90°)

• Example: CH4 is tetrahedral.

X. Electron Domain & Hybridization

Electron domain refers to regions of electron density (bonds or lone pairs) around a central atom. Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Electron Domain: Each bond (single, double, triple) and lone pair counts as one domain.

• Hybridization Types:

  • 2 domains: sp

  • 3 domains: sp2

  • 4 domains: sp3

  • 5 domains: sp3d

  • 6 domains: sp3d2

• Example: In NH3, the central N atom is sp3 hybridized.

Additional info: Students should also review key math skills for chemistry and definitions/concepts from all chapters as recommended in the study guide.