Chem chapter 11 study guide

Chapter 11: Liquids, Solids, and Intermolecular Forces

11.2 Solids, Liquids, and Gases: A Molecular Comparison

This section explores the fundamental differences between the three states of matter—solids, liquids, and gases—based on molecular properties and intermolecular forces.

• Key Properties: Density, molar volume, molecular shape, and strength of intermolecular forces distinguish solids, liquids, and gases.

• Phase Changes: Temperature and pressure can induce phase changes between states.

• Example: Water exists as ice (solid), liquid water, and steam (gas) depending on temperature and pressure.

11.3 Intermolecular Forces: The Forces That Hold Condensed States Together

Intermolecular forces are responsible for the physical properties of liquids and solids. They arise from attractions between molecules and determine boiling points, melting points, and solubility.

• Types of Intermolecular Forces:

  • Dispersion (London) Forces: Present in all molecules due to temporary fluctuations in electron distribution.

  • Dipole-Dipole Forces: Occur between polar molecules due to permanent dipoles.

  • Hydrogen Bonding: A strong dipole-dipole interaction when H is bonded to N, O, or F.

  • Ion-Dipole Forces: Occur between ions and polar molecules, important in solutions.

• Polarity and Dipole Moment: The dipole moment quantifies the separation of charge in a molecule and affects physical properties such as boiling point and solubility.

• Example: Water's high boiling point is due to strong hydrogen bonding.

• Additional info: Dispersion forces increase with molecular size and shape.

11.4 Intermolecular Forces and Physical Properties

Physical properties such as boiling point, melting point, and solubility are determined by the strength and type of intermolecular forces present.

• Boiling and Melting Points: Substances with stronger intermolecular forces have higher boiling and melting points.

• Solubility: "Like dissolves like"—polar substances dissolve in polar solvents, nonpolar in nonpolar solvents.

• Example: Sodium chloride dissolves in water due to ion-dipole interactions.

11.5 Surface Tension and Viscosity

Surface tension and viscosity are two important properties of liquids that arise from intermolecular forces.

• Surface Tension: The energy required to increase the surface area of a liquid. Stronger intermolecular forces result in higher surface tension.

• Viscosity: The resistance of a liquid to flow. Liquids with strong intermolecular forces are more viscous.

• Example: Honey is more viscous than water due to stronger intermolecular attractions.

11.6 Vaporization and Vapor Pressure

Vaporization is the process by which molecules escape from the liquid phase into the gas phase. Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid.

• Factors Affecting Vaporization: Temperature, surface area, and strength of intermolecular forces.

• Distribution of Energies: Molecules in a liquid have a range of kinetic energies; those with enough energy can escape into the gas phase.

• Heat of Vaporization ($riangle H_{vap}$): The energy required to vaporize one mole of a liquid at constant temperature.

• Equation: $q = n riangle H_{vap}$

• Clausius-Clapeyron Equation: Relates vapor pressure and temperature: $\ln P = -\frac{\triangle H_{vap}}{R} \left( \frac{1}{T} \right) + C$

• Example: Water has a high heat of vaporization due to strong hydrogen bonding.

11.6 Sublimation and Deposition

Sublimation is the direct transition from solid to gas, while deposition is the reverse process.

• Sublimation: Occurs when molecules escape directly from the solid phase to the gas phase.

• Deposition: Gas molecules transition directly to the solid phase.

• Example: Dry ice (solid CO2) sublimes at room temperature.

11.7 Heating Curve for Water

The heating curve for water illustrates the temperature changes as heat is added, showing phase transitions from solid to liquid to gas.

• Segments: Each segment represents a phase or phase change (e.g., melting, boiling).

• Calculations: Use $q = m c riangle T$ for temperature changes and $q = n riangle H$ for phase changes.

• Example: Heating ice from -10°C to steam at 110°C involves multiple steps and energy calculations.

11.8 Phase Diagrams

Phase diagrams graphically represent the states of matter of a substance as a function of temperature and pressure.

• Regions: Solid, liquid, and gas regions are separated by lines indicating phase boundaries.

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

• Critical Point: The temperature and pressure above which the liquid and gas phases are indistinguishable.

• Example: Water's phase diagram shows the triple point at 0.01°C and 0.006 atm.

• Additional info: Phase diagrams can be used to predict the relative density and phase transitions under different conditions.