Intermolecular Forces and Phase Diagrams: Study Notes

Intermolecular Forces and Phase Diagrams

1. Intermolecular Forces (IMFs)

Intermolecular forces are the forces of attraction or repulsion between neighboring particles (atoms, molecules, or ions). They are responsible for the physical properties of substances, such as boiling and melting points, vapor pressure, and solubility.

• Types of Movement:

  • Solid: Particles vibrate in fixed positions.

  • Liquid: Particles move past each other but remain close together.

  • Gas: Particles move freely and are far apart.

• Limiting Movement: IMFs limit the movement of particles, especially in solids and liquids. In gases, IMFs are much weaker, but still present (as described by the van der Waals equation): $[P + a(n/V)^2](V - nb) = nRT$

2. Nature of Intermolecular Forces

• IMFs have a range of strengths, called a spectrum.

• IMFs are short-range attractions (significant only when particles are close).

• Average IMF strength varies from weak (London dispersion) to strong (ion-dipole).

• Contrast with Intramolecular Forces: Intramolecular forces (e.g., covalent bonds) hold atoms together within a molecule and are much stronger than IMFs.

Example: Water molecules are held together by hydrogen bonds (IMFs), but the atoms within each water molecule are held by covalent bonds (intramolecular).

3. Review of Molecular Shapes and Polarity

To assess IMFs, you must:

1. Draw correct Lewis dot structures for molecules.

2. Draw the molecule in 3D (using VSEPR theory).

3. Determine the molecule's polarity (or lack thereof).

Example: CH4 is nonpolar; H2O is polar.

4. Ranking the IMFs

All IMFs arise from attractions of some form of positive and negative charges (full or partial).

• Ion-Dipole: Strongest IMF; occurs between an ion and a polar molecule. Important in solutions (e.g., salt in water).

• Hydrogen Bonding: Not a true bond, but a strong dipole-dipole attraction. Occurs when H is bonded to N, O, or F. Example: H2O, NH3, HF.

• Dipole-Dipole: Occurs between polar molecules. Strength depends on the magnitude of the dipole moment and molecular orientation.

• London Dispersion Forces (LDF): Weakest IMF; present in all molecules, especially nonpolar ones. Arise from temporary (instantaneous) dipoles due to electron movement.

5. Hydrogen Bonding

• Special case of dipole-dipole interaction.

• Requires H bonded to N, O, or F.

• Responsible for unique properties of water and biological molecules (e.g., DNA base pairing).

Examples: H2O, NH3, HF, and hydrogen bonding in DNA.

6. Dipole-Dipole and London Dispersion Forces

• Dipole-Dipole: Strength depends on the polarity of the molecules. Molecules align to maximize attraction between opposite charges.

• London Dispersion: All molecules have LDFs. Strength increases with molecular size and polarizability.

Polarizability: The ease with which the electron cloud of a molecule can be distorted, increasing with size and number of electrons.

Induced Dipole: Nonpolar molecules can become temporarily polar due to the presence of nearby charges.

7. Properties of Liquids Resulting from IMFs

IMFs influence several key properties of liquids:

• Surface Tension: The energy required to increase the surface area of a liquid. Caused by IMFs pulling molecules at the surface inward.

• Capillary Action: The ability of a liquid to flow up a narrow tube against gravity, due to adhesive and cohesive forces.

• Viscosity: Resistance to flow. Increases with stronger IMFs and decreases with higher temperature.

• Boiling Point: The temperature at which vapor pressure equals atmospheric pressure. Higher IMFs lead to higher boiling points.

8. Phase Changes

Phase changes involve energy changes and are classified as endothermic (absorbing heat) or exothermic (releasing heat).

• Melting (fusion): Solid to liquid ($\Delta H_{fus}$)

• Vaporization: Liquid to gas ($\Delta H_{vap}$)

• Sublimation: Solid to gas ($\Delta H_{sub}$)

• Freezing: Liquid to solid ($-\Delta H_{fus}$)

• Condensation: Gas to liquid ($-\Delta H_{vap}$)

• Deposition: Gas to solid ($-\Delta H_{sub}$)

Relationship: $\Delta H_{sub} = \Delta H_{fus} + \Delta H_{vap}$

9. Heating and Cooling Curves

Heating/cooling curves show temperature changes as heat is added or removed. Plateaus represent phase changes where temperature remains constant as energy is used to break IMFs.

10. Clausius-Clapeyron Equation

Describes the relationship between vapor pressure and temperature:

$\ln P = -\frac{\Delta H_{vap}}{R}\left(\frac{1}{T}\right) + C$

• Used to calculate boiling points at different pressures.

• Higher $\Delta H_{vap}$ means lower volatility and higher boiling point.

11. Special Properties of Water

• High specific heat capacity

• High heat of vaporization

• Density anomaly: ice is less dense than liquid water

• Excellent solvent due to polarity and hydrogen bonding

12. Phase Diagrams

Phase diagrams show the state of a substance at various temperatures and pressures.

• Critical Point: The end point of the liquid-gas boundary; above this, the substance is a supercritical fluid.

• Triple Point: The unique set of conditions where all three phases coexist.

• Supercritical Fluid: Has properties of both liquids and gases; used in applications like decaffeination.

Example Table: Phase Change Enthalpies

Phase Change	Symbol	Sign of ΔH
Melting (fusion)	ΔHfus	+
Freezing	ΔHfus	-
Vaporization	ΔHvap	+
Condensation	ΔHvap	-
Sublimation	ΔHsub	+
Deposition	ΔHsub	-

13. Practice Problems and Applications

• Predicting phase changes and energy changes for various substances.

• Interpreting and drawing phase diagrams for water, CO2, and ammonia.

• Calculating energy changes using enthalpy values and mass.

• Understanding the relationship between IMFs and physical properties (e.g., boiling point, solubility).

Summary Table: Types of Intermolecular Forces

Type of IMF	Relative Strength	Occurs Between	Example
Ion-Dipole	Strongest	Ion and polar molecule	Na+ in H2O
Hydrogen Bonding	Strong	H bonded to N, O, or F	H2O, NH3
Dipole-Dipole	Moderate	Polar molecules	HCl, SO2
London Dispersion	Weakest	All molecules	I2, CH4

Additional info: These notes expand on the provided class notes by including definitions, examples, and equations for clarity and completeness. The tables summarize key comparisons and classifications relevant to IMFs and phase changes.