Intermolecular Forces, Properties of Liquids, Phase Changes, and Solubility – General Chemistry II Study Guide

IMFs and Properties of Liquids

Types of Intermolecular Forces (IMFs)

Intermolecular forces (IMFs) are the forces of attraction or repulsion between neighboring molecules. They play a crucial role in determining the physical properties of substances, such as boiling point, melting point, viscosity, and surface tension.

• London Dispersion Forces: Weak, temporary forces present in all molecules, especially significant in nonpolar molecules. They arise due to momentary fluctuations in electron distribution, creating temporary dipoles.

• Dipole–Dipole Forces: Attractive forces between polar molecules, where the positive end of one molecule is attracted to the negative end of another.

• Hydrogen Bonding: A special, stronger type of dipole–dipole interaction occurring when hydrogen is bonded to highly electronegative atoms (N, O, or F). Responsible for many unique properties of water.

• Ion–Dipole Interactions: Occur between an ion and a polar molecule. Important in solutions where ionic compounds are dissolved in polar solvents like water.

Order of Strength: Ion–dipole > Hydrogen bonding > Dipole–dipole > London dispersion forces

Vapor Pressure

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid at a given temperature. It increases with temperature.

• Effect of IMFs: Stronger IMFs result in lower vapor pressure because molecules are held together more tightly.

• Boiling Point: The temperature at which the vapor pressure of a liquid equals the external pressure. For water at 1 atm, the normal boiling point is 100°C.

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

Surface Tension and Viscosity

• Surface Tension: The energy required to increase the surface area of a liquid. Caused by cohesive forces between molecules at the surface. Increases with stronger IMFs.

• Viscosity: A measure of a liquid's resistance to flow. Higher viscosity means a liquid flows more slowly. Increases with stronger IMFs and decreases with higher temperature.

Key Comparisons and Applications

• Compare surface tension and viscosity across different liquids.

• Explain why stronger IMFs lower vapor pressure but raise boiling point.

• Relate boiling point to IMF strength.

• Predict which molecules can hydrogen bond.

• Order compounds by IMF strength.

• Explain boiling point trends using IMFs.

Phase Changes, Energy, and Phase Diagrams

Phase Changes and Energy

Phase changes involve the transformation of a substance from one state of matter to another (solid, liquid, gas). Energy is absorbed or released during these changes.

• ΔHfus: Enthalpy of fusion (melting)

• ΔHvap: Enthalpy of vaporization (boiling)

• ΔHsub: Enthalpy of sublimation (solid to gas)

• Equation for energy change: $q = mC\Delta T$ (for temperature change within a phase) $q = n\Delta H$ (for phase change at constant temperature)

Phase Diagrams

A phase diagram shows the state of a substance at various temperatures and pressures. Key points include:

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

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

• Melting Point: Temperature at which solid and liquid phases are in equilibrium at 1 atm.

• Boiling Point: Temperature at which liquid and gas phases are in equilibrium at 1 atm.

Phase diagrams help interpret the physical properties of molecules and predict phase changes under different conditions.

Vapor Pressure

Calculating Vapor Pressure

Vapor pressure can be calculated at a given temperature using the Clausius–Clapeyron equation:

• $\ln \left( \frac{P_2}{P_1} \right) = -\frac{\Delta H_{vap}}{R} \left( \frac{1}{T_2} - \frac{1}{T_1} \right)$

• Normal boiling point: The temperature at which vapor pressure equals 1 atm.

Solids and Crystal Structures

Types of Solids

Solids can be classified based on the nature of their bonding and structure:

• Ionic Solids: Composed of ions held together by electrostatic forces (e.g., NaCl).

• Molecular Solids: Composed of molecules held together by IMFs (e.g., ice, CO2).

• Covalent Network Solids: Atoms connected by covalent bonds in a continuous network (e.g., diamond, quartz).

• Metallic Solids: Metal atoms held together by a 'sea' of delocalized electrons (e.g., copper, iron).

Crystal Structures

Crystalline solids have highly ordered structures. The main types of crystal lattices are:

• Simple Cubic (SC): Atoms at the corners of a cube.

• Body-Centered Cubic (BCC): Atoms at the corners and one in the center of the cube.

• Face-Centered Cubic (FCC): Atoms at the corners and centers of each face of the cube.

Radius and Edge Length: The radius of atoms/ions relates to the edge length of the unit cell, depending on the lattice type.

Crystal Structure	Atoms per Unit Cell	Edge Length (a) in terms of radius (r)
Simple Cubic (SC)	1	$a = 2r$
Body-Centered Cubic (BCC)	2	$a = \frac{4r}{\sqrt{3}}$
Face-Centered Cubic (FCC)	4	$a = \frac{4r}{\sqrt{2}}$

• Be able to calculate radii from density or vice versa.

• Determine molecular formula from a unit cell and ion ratio.

Solubility

Solubility Principles

Solubility is the ability of a substance to dissolve in a solvent. It depends on the nature of solute and solvent, temperature, and pressure.

• Like Dissolves Like: Polar substances dissolve in polar solvents; nonpolar in nonpolar.

• When a nonpolar compound is mixed with a polar one, they generally do not mix.

• When two polar or two nonpolar compounds are mixed, they are more likely to be miscible.

• Hydrogen Bonding: Increases solubility in water for compounds capable of hydrogen bonding.

• Ionic Solids: Dissolve in water due to ion–dipole interactions.

Solubility Graphs and Saturation

• Solubility graphs show how solubility changes with temperature.

• Saturated Solution: Contains the maximum amount of solute at a given temperature.

• Unsaturated Solution: Contains less than the maximum amount of solute.

• Supersaturated Solution: Contains more than the maximum amount of solute; unstable.

• Heat generally increases solubility of solids in liquids but can decrease solubility of gases.

Key Skills

• Explain why ionic solids dissolve in water.

• Interpret solubility graphs.

• Identify if a solution is saturated, unsaturated, or supersaturated.

• Understand how heat affects solubility.