020325 CHEM 102 SP25 Unit 03 pre

Preparation Before Class

• Make sure you can:

  • Calculate edge length from volume of a cube.

  • Calculate volume of a cube from edge length.

  • Convert between mass and volume using density.

  • Convert between grams and moles.

  • Convert between atoms and moles.

Unit Overview

Unit 3: Liquids and Solids

Classifying Liquids and Solids

• Types of substances include:

  • Ionic Compounds

  • Covalent Network

  • Metals

  • Polymers

  • Molecular Compounds

Ionic Bonds and Lattice Energy

• Lattice Energy (ΔHlattice):

  • Definition: The energy required to separate one mole of a solid ionic compound into its gaseous ions.

  • Formula Example:

    • NaCl(s) → Na+(g) + Cl-(g)

  • Factors affecting lattice energy:

    1. Magnitude of charge on the ions (Higher charge = higher lattice energy)

    2. Ionic radius (Smaller radius = higher lattice energy)

• Implications:

  • Stronger attractions increase ΔHlattice and result in higher boiling points.

Predicting Lattice Energy

• Comparison of lattice energies:

  • LiF < MgO < ScN (Increasing charge = increasing lattice energy)

  • LiI > NaI > KI (Decreasing size = increasing lattice energy)

• Lattice energy values:

  • LiF: 1017 kJ/mol (smallest)

  • MgO: 3890 kJ/mol (intermediate)

  • ScN: 7547 kJ/mol (largest)

  • LiI: 732 kJ/mol (largest among compounds)

  • NaI: 686 kJ/mol (intermediate)

  • KI: 632 kJ/mol (smallest)

Classification of Substances

• Identify each listed substance:

  • Nitrogen: Molecular

  • Nickel: Metal

  • Carbon: Covalent Network

  • Lithium Fluoride: Ionic

  • Fluorine: Molecular

  • Carbon Dioxide: Molecular

  • Silicon Dioxide: Covalent Network

  • Magnesium Nitrate: Ionic

  • Iron: Metal

Crystalline vs Amorphous Solids

• Crystalline Solid:

  • Contains ordered, repeating units of atoms, molecules, or ions.

• Amorphous Solid:

  • Lacks the order found in crystalline solids.

Content Exclusion

• Note: X-ray crystallography and Bragg’s Law will not be covered or tested.

Unit Cells

• Definition: The simplest representation of the repeating pattern within a crystal.

• Focus: Cubic unit cells in this course.

Types of Cubic Unit Cells

Simple Cubic Cell

• Atoms: 1 (1/8 at each corner)

• Calculation: 8 corners × 1/8 = 1 atom per unit cell.

Face-Centered Cubic Cell

• Atoms: 4 (1 at each corner and 1 at each face)

• Calculation: 8 corners × 1/8 + 6 faces × 1/2 = 4 atoms per unit cell.

Body-Centered Cubic Cell

• Atoms: 2 (1 at each corner and 1 in the center)

• Calculation: 8 corners × 1/8 + 1 center = 2 atoms per unit cell.

Summary of Cubic Unit Cells

• Unit Cell Overview:

  • Simple Cubic: 1 atom, edge length = 2r, volume = 873, packing efficiency = 52.4%

  • Face-Centered Cubic: 4 atoms, edge length = 2r√2, volume = 16r³√2, packing efficiency = 74.0%

  • Body-Centered Cubic: 2 atoms, edge length = 4r/√3, volume = 64r³, packing efficiency = 68.0%

Characteristics of a Unit Cell

• Key properties include:

  • Unit cell type (impacts # of atoms)

  • Edge length

  • Atomic radius

  • Density

  • Mass

  • Volume

  • Molar mass

• Problem-solving steps:

  1. Identify the quantity to solve for.

  2. Identify given information.

  3. List known information (e.g., molar mass).

  4. Use problem-solving skills and equations.

  • Example formulas:

    • V = a³

    • d = m/V

    • 1 mol = 6.022 × 10²³ atoms

Close-Packed Structures

• Definition: A structure in which metal atoms pack together in the closest arrangement.

• Characteristics:

  • Each central sphere is surrounded by six identical spheres.

  • Two arrangements for stacking layers of atoms.