4A+Types+of+Solids+and+Lattice+Structures+S25

Inter-Particle Forces and Melting Points

• List of substances and their inter-particle forces:

  • 1. He atoms

    • Forces: Dispersion (weak due to small size)

  • 2. CH3Cl

    • Forces: Dispersion, Dipole-dipole

  • 3. CH3OH

    • Forces: Dispersion, Hydrogen bonding

  • 4. CH3COOH

    • Forces: Dispersion, Hydrogen bonding (moderate due to pi electrons in C=O bond)

  • 5. AlCl3

    • Forces: Dispersion, Ion-ion

• Ranked Order of Melting Points (Lowest to Highest):

  1. He atoms (1)

  2. CH3Cl (2)

  3. CH3OH (3)

  4. CH3COOH (4)

  5. AlCl3 (5)

Properties of Solids

• Particles are in close contact and primarily vibrate, without translational motion.

• Intermolecular forces play a significant role in solid properties.

• Many solids exhibit a structured arrangement, known as a crystal lattice.

Crystalline vs Amorphous Solids

• Crystalline Solids

  • Defined by a regular, repeating pattern.

  • Exhibit three-dimensional structure (e.g., silicon dioxide (SiO2)).

• Amorphous Solids

  • Lack a defined structure and have irregular patterns.

Examples of Molecular Solids

• Carbon Dioxide (CO2)

  • Structure: Small nonpolar molecules

  • Melting Point: −78 °C

• Iodine (I2)

  • Structure: Larger nonpolar molecules

  • Melting Point: 114 °C

• General Properties of Molecular Solids:

  • Held together by intermolecular forces (e.g., dispersion, dipole-dipole, hydrogen bonding).

  • Typically have low melting points and do not conduct electricity.

Examples of Covalent Network Solids

• Composed of atoms bound by covalent bonds.

• Properties:

  • Extremely hard

  • High melting points (due to breaking covalent bonds)

  • Non-conductive

Ionic Solids

• Consist of ions bound by ionic bonds (attraction between cations and anions).

• Properties:

  • High melting points (due to breaking ionic bonds)

  • Very hard and brittle

  • Non-conductive

Properties of Solid Types

• Molecular Solids: Localized electrons, discrete molecules, held together by intermolecular forces.

• Ionic Solids: Localized electrons, held together by ionic bonds.

• Covalent Network Solids: All atoms covalently bonded, extensive networks.

• Metallic Solids: Delocalized valence electrons, conductive.

Unit Cells and Crystal Lattice Structures

• Unit Cells: Smallest repeating unit in a crystal structure, analogous to bricks in a building.

• Types of Unit Cells:

  • Simple cubic (SCC)

  • Body centered cubic (BCC)

  • Face centered cubic (FCC)

• Total of 14 possible crystal geometries.

Coordination Number and Packing in Cubic Cells

• SCC (Simple Cubic)

  • Pack density: 1 atom, coordination number = 6.

• BCC (Body Centered Cubic)

  • Pack density: 2 atoms, coordination number = 8.

• FCC (Face Centered Cubic)

  • Pack density: 4 atoms, coordination number = 12.

Calculating Radius and Density

• Atomic Radius in SCC:

  • Example: Alpha Polonium - edge length = 336 pm

  • Radius = Edge Length / 2 -> 168 pm

• Density:

  • D = Mass of atoms per unit cell / Volume of unit cell.

Summary of Density Calculation in Different Cubes

• SCC Density: Calculated using mass of 1 atom, edge length.

• FCC and BCC Density: Consider multiple atoms per unit cell.

Other Crystal Lattice Shapes

• Types:

  • Cubic

  • Tetragonal

  • Orthorhombic

  • Monoclinic

  • Triclinic

  • Hexagonal

  • Rhombohedral

Closest Packing Arrangements

• Efficiency: Face-centered packing is the most efficient.

• Hexagonal Close Packing (HCP): Alternative packing arrangement.

Structure of Ionic Crystals

• Ions arrange in a manner maximizing the attraction while minimizing repulsion.

• Cations placed in specific holes (e.g., tetrahedral and octahedral) in a closely packed anion array.

• Example: NaCl

  • Arrangement of Na+ and Cl- ions in the crystal lattice.

Stoichiometry in Ionic Compounds

• Different stoichiometries affect packing. Example: NaCl vs CsCl; differences arise from ionic sizes.

• Stoichiometric Ratios in CaF2:

  • 8 Ca2+ exchanges and 6 F- yields a stoichiometry of 1:2.

Crystal Growth and Habits

• Crystal habit reflects unit cell shape; varies based on conditions during growth phases.

• Examples:

  • Octahedral growth patterns.

  • Mixed habits often observed in natural crystals.