Intermolecular Forces: LDF and Dipole-Dipole

• Kinetic molecular model explains solid and liquid properties using intermolecular forces and kinetic energy.
• Intermolecular forces attract neighboring particles and kinetic energy keeps particles moving and apart.
• Kinetic energy depends on temperature.

Intermolecular Forces

• Attractive forces between neighboring particles in one or more substances.

London Dispersion Forces (LDF)

• Weakest intermolecular forces present in all molecules.
• Arise from temporary fluctuations in electron distribution, creating temporary dipoles.
• Strength increases with molecule size and shape due to more electrons and larger surface area.
• Methane (CH4) is gas at room temperature due to weak LDF.
• Octane (C8H18) is liquid at room temperature due to stronger LDF compared to methane.
• Iodine (I2) is solid at room temperature due to strong LDF.

Dipole-Dipole Forces

• Occur between polar molecules with permanent dipole moments.
• Positive end of one molecule attracts to the negative end of another.
• Stronger than LDF for molecules of similar size.
• Acetone (CH3COCH3) has relatively high boiling point compared to similarly sized nonpolar molecules due to dipole-dipole forces from the carbonyl group (C=O).
• Acetaldehyde (CH3CHO) experiences dipole-dipole interactions due to the carbonyl group.
• Hydrogen sulfide (H2S) experiences dipole-dipole interactions due to bent geometry and electronegativity difference.

Hydrogen Bonding

• Strong type of dipole-dipole interaction between hydrogen bonded to nitrogen, oxygen, or fluorine.
• Small size of hydrogen and high bond polarity lead to strong attractions.
• Water (H2O) hydrogen bonding causes unusually high boiling point and surface tension.
• Ethanol (CH3CH2OH) can form hydrogen bonds due to the -OH group, contributing to a relatively high boiling point.
• Ammonia (NH3) can form hydrogen bonds, affecting properties like water solubility.

Ion-Dipole Forces

• Occurs between an ion and a polar molecule and is stronger than dipole-dipole forces.
• Sodium Chloride in Water (NaCl(aq)) - Na+ cations are attracted to partially negative oxygen atoms of water, and Cl− anions are attracted to partially positive hydrogen atoms of water molecules, stabilizing the ions in solution.
• Potassium Chloride in Water (KCl(aq)) - K+ cations and Cl− anions are solvated by water molecules.
• Lithium Chloride in Water (LiCl(aq)) - Li+ cations and Cl− anions interact with water molecules.

Properties of Liquids

• Liquids are more ordered than gases but less ordered than solids.

Viscosity

• A liquid's resistance to flow.
• High viscosity means slow flow and low viscosity means easy flow.
• Motor oil viscosity is selected based on engine and climate; low viscosity for cold weather and high viscosity for hot weather.

Surface Tension

• Tendency of liquid's surface to minimize its area to form a "skin" on the surface.
• Water striders can walk on water due to surface tension supporting their weight.

Capillary Action

• Ability of liquid to flow in narrow spaces against gravity, relying on cohesion and adhesion.
• Cohesion is attraction between liquid molecules.
• Adhesion is attraction between liquid molecules and surface.
• Plants use capillary action to transport water from roots to leaves through xylem via cohesion and adhesion.

Vapor Pressure

• Pressure exerted by vapor of a liquid when liquid and vapor are in equilibrium.

• Liquids with higher vapor pressures evaporate more easily.

• Perfumes utilize volatile liquids with high vapor pressures for quick scent evaporation.

• Stronger IMFs generally lead to higher viscosity, higher surface tension, and lower vapor pressure.

• Higher temperatures generally decrease viscosity and surface tension, and increase vapor pressure.

• Meniscus is the curved surface of a liquid in a narrow tube, resulting from interplay between adhesion and cohesion.

• Concave meniscus indicates adhesive forces are stronger than cohesive forces.

• Liquid substances with weak intermolecular forces have high vapor pressure, low viscosity, and low surface tension.

Phase Changes and Phase Diagrams

• Phase changes are physical processes where a substance transitions between solid, liquid, and gas states that involve energy absorption or release.
• Melting/Freezing: Solid to liquid or liquid to solid.
• Boiling/Condensation: Liquid to gas or gas to liquid.
• Sublimation/Deposition: Solid directly to gas or gas directly to solid.
• Endothermic processes absorb energy and include melting, boiling, and sublimation.
• Exothermic processes release energy and include freezing, condensation, and deposition.

Phase Diagrams

• Graphical representation of physical states of a substance under different conditions.
  • Axes: Temperature (x-axis) and pressure (y-axis).
  • Regions: Represent specific phases (solid, liquid, or gas).
  • Lines: Represent equilibrium conditions where two phases coexist.
  • Triple Point: Point where all three phases coexist in equilibrium.
  • Critical Point: Point beyond which a distinct liquid phase doesn't exist; substance becomes a supercritical fluid.

Example: Water (H2O)

• Melting Point Line: Crosses temperature axis at 0°C (32°F) at 1 atm.

• Boiling Point Line: Crosses temperature axis at 100°C (212°F) at 1 atm.

• Triple Point: Approximately 0.01°C (32.018°F) and 0.006 atm (611.73 Pa).

• Critical Point: Approximately 374°C (705°F) and 218 atm.

• Phase changes involve energy transfer.

• Phase diagrams summarize stable phases under various conditions.

• The triple point and critical point are important.

Solution Concentrations and Colligative Properties

• Solution concentration describes amount of solute dissolved in a solvent.
  • Molarity (M): Moles of solute per liter of solution (mol/L).
  • Molality (m): Moles of solute per kilogram of solvent (mol/kg).
  • Mass Percent (%): (Mass of solute / Mass of solution) x 100%
  • Mole Fraction (X): (Moles of solute / Total moles in solution)

Colligative Properties

• Properties of solutions that depend on number of solute particles, regardless of solute type, primarily observed in dilute solutions.
• Vapor Pressure Lowering happens when the vapor pressure of a solution is lower than that of the pure solvent.
• Boiling Point Elevation happens when the boiling point of a solution is higher than that of the pure solvent.
• Freezing Point Depression happens when the freezing point of a solution is lower than that of the pure solvent.
• Osmotic Pressure happens when pressure is required to prevent solvent flow across a semipermeable membrane.
• Solute is the substance being dissolved.
• Solvent is the substance doing the dissolving.
• Solutions are homogeneous mixtures of solute and solvent.
• Colligative properties depend ONLY on number of solute particles, not their identity.
• Ionic compounds dissociate into multiple ions in solution, increasing their effect on colligative properties.
• NaCl dissociates into Na+ and Cl-, effectively doubling the number of particles.