Gen Chem 2-- CH 12

Chapter Twelve: Liquids, Solids, and Intermolecular Forces

• There are three phases of matter:
  1. Gases
  2. Liquids
  3. Solids
• Molecules/atoms are furthest apart in the gas phase and closest together in the solid phase.

12.1 Intermolecular Forces (IMF) vs. Intramolecular Forces

• The state of matter depends on the relative magnitudes of:
  1. Intermolecular forces (IMF)
  2. Thermal energy
• Thermal energy is the energy associated with the random motion of molecules and atoms in matter.
• Gases: Have weak IMF relative to thermal energy.
• Liquids & Solids: Have strong IMF relative to thermal energy.

12.2 Solids, Liquids, and Gases: A Molecular Comparison

12.3 Phase Changes and Intermolecular Forces

• To change a solid to a liquid, add heat (increase temperature).
• To change a gas to a liquid, remove heat (decrease temperature).
• SOLID → LIQUID → GAS (by adding heat)
• GAS → LIQUID (by increasing pressure)
• Example: Propane tanks - Gas becomes liquid under pressure.
• IMF – Intermolecular Forces:
  • Weak IMF
  • Moderate IMF
  • Strong IMF
• IMF result from the attraction between molecules.
• IMF are MUCH weaker than bonds.
• Bonds result from sharing or transferring electrons.

Intermolecular Forces (IMF) Types:

• There are four main types of IMF:
  1. Dispersion Forces
  2. Dipole-Dipole Forces
  3. Hydrogen Bonding
  4. Ion-Dipole Force

1. Dispersion Forces

• All molecules have dispersion forces.
• Arise from instantaneous, temporary dipoles due to electron fluctuations.
• Beryllium example: Beryllium (Be) has 4 electrons.

Magnitude of Dispersion Forces:

• Noble Gas Boiling Point (K):
  • He
  • Ne
  • Ar
  • Kr
  • Xe
• Alkane Boiling Point (K):
  • C5H12
  • C6H14
  • C7H16
  • C8H18
• The larger the molecule (more electrons), the stronger the dispersion forces, and the higher the boiling point.
• I2 has a higher melting point than Cl2 because it is larger and has more electrons.
• The shape of a molecule also influences the strength of IMF. "n-pentane" versus "neopentane"

2. Dipole-Dipole Forces

• Occur in polar molecules.
• Polar molecule has a positive (δ+) end and a negative (δ-) end.
• If a molecule is polar, dipole-dipole forces will be present.
• Examples:
  • CO2
  • CH2Br2
  • CH4
  • NH3

Polarity

• The presence of dipole-dipole forces affects both boiling and melting points.
• If you have two compounds with similar molecular weights, and one has dipole-dipole forces (is polar) while the other does not (is nonpolar), the polar molecule will have a higher boiling point.
• Propane (C3H8): MW = 44.10 g/mole, bp = -42.1 ºC
• Acetaldehyde (CH3CHO): MW = 44.05 g/mole, bp = 20.2ºC

Miscibility

• Miscibility: "likes dissolve likes".
• Polar compounds dissolve polar compounds.
• Nonpolar compounds dissolve nonpolar compounds.

Solubility Example

• CH3Cl would dissolve in water because both are polar.
• CI4 would not dissolve in water because CI4 is nonpolar, but water is polar.

3. Hydrogen Bonding

• Involves polar molecules.
• Occurs when hydrogen is bonded to highly electronegative atoms (N, O, F). *Example molecules with potential hydrogen bonding:
  • N2
  • HF
  • HI
  • CH3CHO
  • CH3OH
• Molecules with H-bonding have higher boiling points. *Comparing boiling points:
  • CH4
  • HCl
  • HF

4. Ion-Dipole Force

• Occurs between ions and polar molecules.
• Ionic compound: formed when a metal bonds with a nonmetal.
• Molecular compound: formed when two nonmetals form a covalent bond. *Examples:
  • Na2O
  • CO2
  • MgO
  • KCl
  • NO2
  • PI5

12.4 Intermolecular Forces in Action

• Influence three different properties:
  1. Surface Tension
  2. Viscosity
  3. Capillary Action

1. Surface Tension

• Refers to the tendency of liquids to minimize surface area.
• Surface molecules are attracted inward, creating a "skin" on the surface.
• The stronger the IMF, the stronger the surface tension.

2. Viscosity

• Measure in poise (P); g/cm·s.
• Viscosity depends on three things:
  1. Temperature
  2. IMF
  3. Molecular shape

3. Capillary Action

• Result of two forces:
  1. Cohesive forces: IMF between liquid molecules.
  2. Adhesive forces: IMF between liquid and surface of the container.
• Two possible scenarios:
  1. Cohesive < Adhesive
  2. Cohesive > Adhesive

12.5 Vaporization and Vapor Pressure

• Molecules in a liquid are in constant motion due to thermal energy.
• Molecules at the surface have fewer "neighbors", so less energy is required for them to become gas molecules.
• Vaporization: Liquid to gas.
• Condensation: Gas to liquid.
• Three things that affect the rate of vaporization:
  1. Temperature
  2. Surface area
  3. Strength of IMF
• Vaporization is endothermic (requires energy).
• Condensation is exothermic (releases heat).
• Heat of Vaporization (ΔHvap):
  • Definition: Energy required to vaporize one mole of a liquid at a given temperature.
  • For water:
    • ΔH_{vap} = 40.7 \frac{kJ}{mole} at 100ºC
    • ΔH_{vap} = 44.0 \frac{kJ}{mole} at 25ºC

Heat Required to Vaporize Water

*To calculate the heat required to vaporize 575 g of water at 100ºC, use the following:

• Dynamic Equilibrium: Rate of vaporization equals rate of condensation.
• Vapor Pressure: Pressure exerted by a gas in equilibrium with its liquid.

Vapor Pressure

• Flask #1: Water (polar, strong IMF)
• Flask #2: n-pentane (nonpolar, weak IMF)
• Higher IMF = Lower vapor pressure.

Boiling Point

• Normal Boiling Point: Temperature at which the vapor pressure equals atmospheric pressure.
• When water boils, it remains at that temperature until ALL the liquid is gone. Once ALL the liquid has vaporized, the temperature of the vapor will start rising.

Temperature and Vapor Pressure

• If you increase the temperature of a liquid, the vapor pressure also increases.
• For water:
  • T = 25ºC, P_{vap} = 23.3 Torr
  • T = 60ºC, P_{vap} = 149.4 Torr

Clausius-Clapeyron Equation

• The relationship between temperature and vapor pressure is not linear.
• The equation is given for exams/quizzes. \ln(P{vap}) = \frac{-\Delta H{vap}}{RT} + \Beta
• ß = constant for each different gas
• T = temperature (Kelvin)
• R = 8.314 J/mol·K
• P_{vap} = Vapor Pressure (atm)
• Rearranged: \ln(\frac{P1}{P2}) = \frac{-\Delta H{vap}}{R} (\frac{1}{T2} - \frac{1}{T1}) Example: If you know ΔH{vap} of acetone is 29.1 kJ/mole and its normal boiling point is 56.1 ºC, what is the vapor pressure of acetone at 15.5ºC?

12.6 Sublimation and Fusion

• Sublimation: Solid to gas.
• Deposition: Gas to solid.
• Fusion: Solid to liquid (melting).
• Freezing: Liquid to solid.
• \Delta H_{fus}: Heat of fusion (energy required to melt one mole of a solid).
• The heat of vaporization is normally a lot larger than the heat of fusion because you must overcome all IMF to go from liquid to gas, while you only need to disrupt some IMF to go from solid to liquid.

12.7 Heating Curve for Water

How much energy is required to heat a 235 g sample of water from -45.5°C to 144°C?

• C_{ice} = 2.09 J/g∙°C
• ΔH_{vap} = 40.7 kJ/mol (at 100°C)
• C_{water} = 4.184 J/g∙°C
• ΔH_{fus} = 6.02 kJ/mol (at 0°C)
• C_{vapor} = 2.01 J/g∙°C

12.8 Phase Diagrams

• A map of the phases that exist at varying temperatures and pressures, shown as a graph.
• Triple Point: The temperature and pressure at which all three phases (solid, liquid, gas) are in equilibrium.
• Critical Point: The temperature and pressure above which a distinct liquid phase does not exist.
• Supercritical Fluid: A state where neither a gas nor a liquid is present; occurs at very high temperatures and pressures. Possesses properties of both gas and liquid. It can effuse through solids like a gas, and dissolve materials like a liquid. Supercritical fluids are used in a variety of applications, including as a solvent in chemical reactions and extractions. The most common supercritical fluid is supercritical CO2. A key application of supercritical CO2 is in decaffeinating coffee. It also has the advantage of being environmentally friendly.