Inter and Intramolecular forces

This document explores intermolecular forces, phase changes, and the properties of liquids and solids, emphasizing how molecular interactions govern physical behavior.

1. Phases of Matter (Physical States)

• Condensed Phases: Liquids and solids are denser than gases.

• Phases:

  • Gas: Molecules move freely with little interaction.

  • Liquid: Molecules held together by intermolecular forces.

  • Solid: Molecules tightly packed in a structured form.

2. Phase Changes and Phase Diagrams

Phase Changes

• Condensation: Gas → Liquid.

• Vaporization: Liquid → Gas.

• Freezing: Liquid → Solid.

• Melting: Solid → Liquid.

• Sublimation: Solid → Gas.

• Deposition: Gas → Solid.

Phase Diagrams

• Depict phases at different pressures and temperatures.

• Supercritical Fluid: When temperature exceeds the critical temperature and pressure exceeds the critical pressure, the liquid and gas phases become indistinguishable.

  • Example: Supercritical CO₂ is used to decaffeinate coffee.

3. Relationship Between Energy, Phase, and Intermolecular Forces

• Kinetic Energy (KE): Energy of motion.

• Boiling Point:

  • Temperature at which vaporization and condensation reach equilibrium.

  • Determined by the strength of intermolecular forces.

• Energy Changes:

  • As temperature decreases, KE drops, and molecules move slower, increasing intermolecular forces.

4. Intermolecular Forces (IMFs)

Types of Intermolecular Forces

1. Ion-Dipole Forces:

   • Between ions and polar molecules.

   • Stronger with higher charge and greater dipole moment.

   • Example: NaCl dissolving in water.

   • Strength: >50 kJ/mol.

2. Dipole-Dipole Forces:

   • Between polar molecules.

   • Stronger for molecules with higher polarity.

   • Example: Acetone (CH₃COCH₃).

   • Strength: 2–15 kJ/mol.

3. Hydrogen Bonding:

   • Special dipole-dipole force where H bonds with N, O, or F.

   • Strong due to high electronegativity and small size.

   • Example: H₂O, NH₃, DNA hydrogen bonding (zipper effect).

   • Strength: 10–40 kJ/mol.

4. London Dispersion Forces (LDFs):

   • Present in all molecules, strongest in large, polarizable atoms/molecules.

   • Example: Cl₂, noble gases, hydrocarbons.

   • Strength: 0.1–30 kJ/mol.

   • Increases with molecular weight and shape.

Ranking of Intermolecular Forces (Strength)

London Dispersion < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole

5. Select Properties of Liquids

1. Viscosity (Resistance to Flow):

   • Higher molecular weight = Higher viscosity.

   • Stronger IMFs = Higher viscosity.

   • Example: Honey (high viscosity) vs. water (low viscosity).

2. Surface Tension (Resistance to increasing surface area):

   • Higher IMFs = Higher surface tension.

   • Example: Water droplets forming beads.

3. Critical Temperature and Pressure:

   • Critical Temperature: Highest temperature at which a liquid can exist.

   • Critical Pressure: Pressure needed to liquefy a gas at the critical temperature.

6. Vapor Pressure and Boiling Point

• Vapor Pressure: Pressure exerted by a liquid’s vapor when at equilibrium.

  • Higher vapor pressure = Easier to evaporate (more volatile).

  • Stronger IMFs = Lower vapor pressure.

• Boiling Point:

  • When vapor pressure equals external pressure.

  • Lower atmospheric pressure (e.g., on Mount Everest) = Lower boiling point.

7. Solids

Types of Solids

1. Molecular Solids:

   • Held by IMFs.

   • Example: Ice (H₂O).

2. Ionic Solids:

   • Held by ionic bonds.

   • High melting points, brittle, conduct electricity in solution.

   • Example: NaCl (table salt).

3. Covalent Network Solids:

   • Atoms linked by covalent bonds.

   • Very hard, high melting points.

   • Example: Diamond (C), Quartz (SiO₂).

4. Metallic Solids:

   • Metal atoms held by metallic bonds (delocalized electrons).

   • Good conductors of electricity and heat.

   • Example: Gold (Au), Bismuth (Bi).

Ranking of Solid Strength

Molecular < Metallic/Ionic < Covalent Network

8. Summary and Trends

Intermolecular Forces and Boiling Point Trends

• Boiling Point Increases With:

  1. Stronger IMFs (London < Dipole-Dipole < H-Bonding < Ion-Dipole).

  2. Higher Molecular Weight.

  3. Linear Shape (More Surface Area = More Dispersion Forces).

Intermolecular Forces in Different Substances

Comparing Boiling Points

Comparison of Solid Types

Key Takeaways

1. Intermolecular forces (IMFs) determine phase changes, boiling points, and solubility.

2. London Dispersion Forces exist in all molecules but are weak.

3. Hydrogen bonding occurs only in molecules with N-H, O-H, or F-H bonds.

4. Stronger IMFs = Higher boiling point, higher viscosity, and lower vapor pressure.

5. Solid types vary in bonding, with molecular solids being the weakest and covalent networks being the strongest.

This comprehensive summary covers all topics from the document, including concepts, equations, trends, and example problems. Let me know if you need further clarification! 🚀