Chapter 11: Liquids, Solids and Intermolecular Forces

Structure Determines Properties

• Chemical composition and molecular structure dictate intermolecular forces.

• Intermolecular forces are attractive forces between molecules and atoms, holding liquids and solids together.

• Phases of matter (solid, liquid, gas) depend on the strength of intermolecular forces relative to thermal energy:

  • High thermal energy = gas

  • Low thermal energy = liquid or solid.

Properties of Phases of Matter

• Definite Shape: Solids maintain shape; Indefinite Shape: Liquids take the shape of the container.

Comparison of Phases

• Density: Ice has lower density than liquid water, which is atypical for solids.

• Ice and liquid water have densities close to each other, while steam's density is significantly lower.

Liquids

• Particles in liquids are closely packed with limited movement, making them incompressible.

• Liquids take the shape of their container but do not fill it completely.

Explaining Liquids' Properties

• Higher densities than gases, definite volume, and indefinite shape due to particle proximity.

Gases

• Particles have complete freedom of motion and negligible intermolecular forces, allowing gases to expand and fill their container.

• Gases are compressible due to large spaces between particles.

Solids

• Particles are tightly packed and fixed, resulting in definitive shape and volume.

• Crystalline Solids: Ordered structure (e.g., salt, diamonds).

  • Amorphous Solids: Lack regular geometric patterns (e.g., plastic, glass).

Phase Changes

• Changes in state can occur through heating, cooling, or pressure changes.

Intermolecular Forces Overview

• Attractive forces keeping matter in condensed states; all are electrostatic.

• Stronger intermolecular forces result in higher boiling and melting points.

Strength Trends

• Stronger attractions = more energy needed to separate particles and higher boiling points.

Causes of Molecular Attraction

• Intermolecular attractions arise from oppositely charged interactions (e.g., ions, dipoles).

Types of Intermolecular Forces

• Dispersion Forces: Temporary polarity from electron distribution.

• Dipole-Dipole Attractions: Permanent polarity in polar molecules; strong when H is bonded to electronegative atoms (hydrogen bonding).

Dispersion Forces

• Result from temporary dipoles; present in all molecules; also called London forces.

Factors Affecting Dispersion Forces

• Polarizability: Greater electron cloud leads to stronger dispersion forces.

• Molecular Shape: More surface contact leads to stronger induced dipoles.

Dipole-Dipole Attractions

• Interactions between positive end of one polar molecule and the negative end of another; stronger due to permanent dipoles.

Vapor Pressure Relations

• Vapor pressure is affected by strength of attractive forces: weaker attractions lead to higher vapor pressure and evaporation is more rapid

Hydrogen Bonding

• Strong type of dipole-dipole interaction; crucial for water's properties.

• affects their structure and behavior in various contexts

Conclusion on Intermolecular Forces

• Dispersion Forces: Weakest type of Intermolecular Force, present in all molecules; strength increases with molar mass.

• Dipole–Dipole: Found in polar molecules; stronger than dispersion, important in molecular interactions

• Hydrogen Bonds: Strongest of intermolecular forces involving H bonded with F, O, or N.

• Ion-Dipole: Strongest in mixtures of ionic and polar compounds; important in solubility.

11.5 Vaporization

• Process where High-energy molecules at the surface escape into vapor.

• The Larger surface area of the liquid, the faster the molecules can escape (rate of evaporation is impacted by surface area)

Distribution of Thermal Energy

• In any liquid, there’s a distribution of thermal energy among molecules

• Only a fraction of liquid molecules escape into the vapor because only molecules that have enough kinetic energy to overcome the attractive forces can escape

• Temperature of the liquid is directly proportional to the fraction of molecules that can escape: as temperature increases, the kinetic energy of the molecules also increases, leading to a higher number of molecules achieving the necessary energy to enter the vapor phase.

Vaporization Process

• Rate increases with temperature, surface area, and weaker intermolecular forces.

• Requires energy to overcome molecular attractions.

Condensation

• Vapor molecules lose energy and return to liquid when colliding with each other.

• leads to formation of droplets on surfaces.

Vaporization vs. Condensation

• Open container: vapor spreads faster than it condenses.

• Closed container: eventual equilibrium between vaporization and condensation.

Effect of Intermolecular Attraction

• Weaker attractive forces lead to faster evaporation and higher vapor pressure.

• Volatile liquids evaporate easily (e.g., gasoline).

• Nonvolatile liquids evaporate slowly (e.g., motor oil).

Heat of Vaporization

• Amount of heat needed to vaporize 1 mole of liquid (Hvap).

• Always endothermic and temperature dependent.

Dynamic Equilibrium

• Defined once vaporization and condensation rates equalize in a closed container.

• System can respond to changes in conditions while maintaining equilibrium.

Vapor Pressure

• Pressure exerted by vapor in dynamic equilibrium with its liquid form.

• Weaker attractive forces result in higher vapor pressure and volatility. (strength of intermolecular force = vapor pressure)

Vapor Pressure vs. Temperature

• Increased temperature raises vapor pressure significantly.

Boiling Point of a Liquid

• Temperature where vapor pressure equals external pressure allows bubble formation.

Normal boiling point is when vapor pressure equals 1 atm.

Clausius–Clapeyron Equation

• Linear relationship between ln(Pvap) and 1/T.

• Slope: -Hvap/R, applicable for calculating vapor pressure changes.

Heating Curve of a Liquid

• Temperature rises linearly until boiling point.

• All added heat after boiling point goes into vaporization, maintaining constant temperature until phase change is done.

Sublimation and Fusion

• Sublimation occurs when solid molecules escape to gas phase; deposition is when vapor forms solid.

• Dynamic equilibrium exists in closed containers below melting point.

Heating Curve of a Solid

• Heating causes temperature increase until melting point, after which added heat is for melting.

Energetics of Melting

• Melting draws energy (endothermic); freezing releases energy (exothermic).

Heat of Fusion

• Heat needed to melt one mole of solid (ΔHfus); endothermic process.

Heating Curve for Water

• Segments for heating ice, melting, heating water, boiling, and heating steam.

• Specific heat values for ice, liquid water, and steam influence calculations.

These values are crucial in determining the amount of energy required for each phase change.

1.8 Phase Diagrams

• Definition: Phase diagrams illustrate different states and transitions based on temperature and pressure.

  • provides info about stability of each state in terms of the conditions they exist under and transition from one form to another

• Regions: Represent distinct states of matter. (divided by lines)

• Lines: Indicate state changes (e.g., vapor pressure curve for liquid/gas).

• Critical Point: Farthest point on the vapor pressure curve. (max pressure and temp at which a substance can exist as a liquid and gas in equilibrium)

• Triple Point: Condition where solid, liquid, and gas coexist.

• Freezing Point: Generally increases with pressure.

Phase Diagram Navigation

• Components:

  • Fusion curve: melting point

  • Sublimation curve: transition from solid to gas

  • Vaporization curve: transition from liquid to gas

• Temperature Ranges: Example points given for CO2 and H2O.

Phase Diagram of CO2

• Temperature at -78°C represents state changes from solid to liquid to vapor as pressure increases and decreases

Phase Diagram of H₂O

• Displays solid, liquid, and vapor states at different pressures and temperatures.

Phase Diagram of Iodine and CO2

• Critical Point: At 535°C and 72.9 atm for CO2; relevant points for iodine included.

The Critical Point

• Critical Temperature: Temperature above which substance can’t exist as liquid and needed to create a supercritical fluid. (intermediate state that blends properties of gases and liquids)

• Critical Pressure: Pressure necessary to liquefy a gas at critical temperature; considered un-condensable.

Supercritical Fluid

• Process: Heating liquid in a sealed container increases pressure and temperature, rising pressure increases the density of vapor and decreases for liquid.

• Phenomenon: Meniscus (boundary surface between gas and liquid) disappears; liquid and gas intermingle, forming supercritical fluid with properties of both states.

Water: An Extraordinary Substance

• Water remains liquid at room temperature, whereas similar substances like NH3 or CH4 are gases.

• Hydrogen Bonding: Contributes to water's unique properties.

• Solvent Capabilities: Effective in dissolving ionic/polar substances due to its polarity; some nonpolar molecules like O2 and CO2 also show solubility.

• High Specific Heat: Allows water to absorb and retain large amounts of heat without a big change in temp —> affects climate regulation.

• Density Change on Freezing: Ice is less dense than liquid water due to expansion upon freezing —> floats

Key Concepts in Intermolecular Forces

• Types of Forces:

  • Dipole-dipole interactions: Attractions between polar molecules, where the positive end of one molecule is attracted to the negative end of another, playing a significant role in determining the physical properties of substances.

  • Hydrogen bonding: specific case of dipole-dipole interaction that occurs when hydrogen is covalently bonded to high electronegative atoms resulting in strong attractions

  • Dispersion forces: weak intermolecular forces arising from temporary shifts in electron density within molecules, which induce temporary dipoles that can attract neighboring molecules, significantly impacting nonpolar substances.

  • Ion-dipole interactions: Forces that occur between an ion and a polar molecule, where the charged ion interacts with the dipole of the polar molecule, often seen in solutions where ionic compounds are dissolved in polar solvents.

• Properties of Liquids: Includes surface tension, viscosity (resistance to flow), and vapor pressure.

• Phase Changes: Expounds on transitions among liquid-vapor, solid-liquid, and solid-vapor states, as well as applicable phase diagrams.