Chapter 11: Liquids, Solids, and Intermolecular Forces – Study Notes

Liquids, Solids, and Intermolecular Forces

Learning Objectives

• Describe how various intermolecular forces hold molecules or atoms together in liquids and solids.

• Understand how intermolecular forces arise, their strength, and their effects on properties such as melting point, boiling point, and state of matter.

• Explain vaporization, including enthalpy change ($\Delta_{vap}H$), and dynamic equilibrium between a liquid and its vapor.

• Determine vapor pressure experimentally and estimate values from tables or graphs.

• Describe the significance of the critical point and how pressure and temperature relate to boiling.

• Interpret phase diagrams and predict changes during heating, cooling, or pressure changes.

• Describe aspects of network covalent bonding in solids.

Solids, Liquids, and Gases: A Molecular Comparison

States of Matter

The three common states of matter—solid, liquid, and gas—differ in the arrangement and movement of their constituent particles.

• Solids: Particles are closely packed in a fixed arrangement and vibrate in place. Solids have definite shape and volume.

• Liquids: Particles are close together but can move past one another, allowing liquids to flow and take the shape of their container while maintaining a definite volume.

• Gases: Particles are far apart and move freely, filling the entire volume of their container and having neither definite shape nor volume.

State	Density	Shape	Volume	Strength of Intermolecular Forces
Gas	Low	Indefinite	Indefinite	Weak
Liquid	High	Indefinite	Definite	Moderate
Solid	High	Definite	Definite	Strong

Phase Changes

Changing temperature or pressure can transform matter between solid, liquid, and gas states. For example, heating ice (solid) turns it into water (liquid), and further heating produces water vapor (gas).

Intermolecular Forces: The Forces That Hold Condensed States Together

Nature of Intermolecular Forces

• Intermolecular forces arise from interactions between charges, partial charges, and temporary charges on molecules, atoms, or ions.

• They are generally much weaker than covalent or ionic (bonding) forces.

• Bonding forces involve large charges at close distances; intermolecular forces involve smaller charges at greater distances.

Types of Intermolecular Forces

Type	Present In	Molecular Perspective
Dispersion (London) Forces	All molecules and atoms	Temporary dipoles
Dipole-Dipole Forces	Polar molecules	Permanent dipoles
Hydrogen Bonding	Molecules with H bonded to F, O, or N	Strong dipole-dipole interaction
Ion-Dipole Forces	Mixtures of ionic compounds and polar molecules	Ions interacting with dipoles
Ion-Induced Dipole	Mixtures of ionic compounds and nonpolar molecules	Ion induces dipole in nonpolar molecule
Dipole-Induced Dipole	Mixtures of polar and nonpolar molecules	Dipole induces dipole in nonpolar molecule

Dispersion Forces (London Forces)

Dispersion forces are present in all atoms and molecules due to fluctuations in electron distribution, which create temporary dipoles.

• Even in nonpolar or symmetrical molecules, instantaneous dipoles can arise, inducing dipoles in neighboring particles and resulting in attraction.

• Dispersion forces are stronger in molecules with more electrons and in elongated (less compact) molecules.

• The ability of a molecule to form temporary dipoles is called polarizability.

Example: Noble gases (e.g., Xe) have higher boiling points than lighter noble gases (e.g., Ne) due to stronger dispersion forces.

Dipole-Dipole Forces

Dipole-dipole forces exist in polar molecules, where permanent dipoles interact with each other.

• All polar molecules have both dispersion and dipole-dipole forces.

• Boiling points increase with increasing dipole moment, all else being equal.

• Polarity also affects miscibility: polar liquids mix with other polar liquids.

Example: Methanal (CH2O) has a higher boiling point than ethene (C2H4) due to dipole-dipole interactions.

Hydrogen Bonding

Hydrogen bonding is a special, strong type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms (F, O, or N).

• Hydrogen bonds are much stronger than regular dipole-dipole or dispersion forces (8–40 kJ/mol), but weaker than covalent bonds.

• Hydrogen bonds are highly directional, leading to unique properties (e.g., high boiling point of water).

Example: Water (H2O) and ammonia (NH3) exhibit hydrogen bonding, resulting in higher boiling points compared to similar molecules without hydrogen bonding.

Ion-Induced Dipole and Dipole-Induced Dipole Forces

• Ion-Induced Dipole: An ion distorts the electron cloud of a nonpolar molecule, inducing a dipole.

• Dipole-Induced Dipole: A polar molecule induces a dipole in a nonpolar molecule.

• The magnitude of these forces depends on the charge of the ion or the size of the dipole and the polarizability of the nonpolar molecule.

Ion-Dipole Forces

Ion-dipole forces occur when ionic compounds are mixed with polar molecules, such as in aqueous solutions. These are important for dissolving salts in water.

Relative Strengths of Intermolecular Forces

• Hydrogen bonding > Dipole-Dipole > Dispersion Force

• All intermolecular forces are weaker than covalent or ionic bonds, except in network covalent solids.

Intermolecular Forces in Action

Surface Tension

Surface tension is the energy required to increase the surface area of a liquid by a unit amount. It arises because molecules at the surface have fewer neighbors and are less stable.

• Liquids minimize surface area to reduce potential energy.

• Surface tension decreases as intermolecular forces decrease.

Example: Water has higher surface tension than benzene due to stronger hydrogen bonding.

Viscosity

Viscosity is the resistance of a liquid to flow. It is measured in pascal-seconds (Pa·s) or poise (P).

• Viscosity increases with stronger intermolecular forces and with longer, more entangled molecules.

• Viscosity decreases with increasing temperature.

Example: Water at room temperature has a viscosity of 1 centipoise (cP).

Capillary Action

Capillary action is the ability of a liquid to flow against gravity up a narrow tube, resulting from cohesive (liquid-liquid) and adhesive (liquid-surface) forces.

• If adhesive forces > cohesive forces, the liquid rises in the tube.

• The rise continues until balanced by gravity.

The Process of Vaporization

Kinetic Energy and Vaporization

• At any temperature, some molecules in a liquid have enough kinetic energy to escape into the gas phase.

• As temperature increases, more molecules can vaporize.

• The rate of vaporization increases with temperature, surface area, and decreasing intermolecular forces.

Enthalpy of Vaporization ($\Delta_{vap}H$)

The enthalpy of vaporization is the heat required to vaporize one mole of a liquid at constant pressure.

• $\Delta_{vap}H^\circ = H_{vapour} - H_{liquid}$

• For water: $\Delta_{vap}H^\circ = 40.7$ kJ/mol at 100°C

Example Calculation: To find the mass of water vaporized by a given amount of heat, use:

$\text{mass} = \frac{q}{\Delta_{vap}H} \times \text{molar mass}$

Vapor Pressure and Dynamic Equilibrium

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

• At equilibrium, the rate of vaporization equals the rate of condensation.

• Disturbing the equilibrium (e.g., changing volume) shifts the system to restore equilibrium.

Temperature Dependence of Vapor Pressure

• Vapor pressure increases with temperature because more molecules have enough energy to escape into the vapor phase.

Clausius-Clapeyron Equation

The Clausius-Clapeyron equation relates vapor pressure and temperature:

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

or, for two temperatures:

$\ln \left( \frac{P_2}{P_1} \right) = -\frac{\Delta_{vap}H}{R} \left( \frac{1}{T_2} - \frac{1}{T_1} \right)$

where $P$ is vapor pressure, $T$ is temperature (in K), $R$ is the gas constant, and $C$ is a constant.

Critical Point

The critical point is the temperature and pressure above which the liquid and vapor phases become indistinguishable. The substance exists as a supercritical fluid.

Heats of Sublimation and Fusion

• Heat of Sublimation ($\Delta_{sub}H$): Energy required to convert a solid directly to a gas.

• Heat of Fusion ($\Delta_{fus}H$): Energy required to melt a solid to a liquid.

• The reverse processes are deposition (gas to solid) and freezing (liquid to solid).

Heating Curves

Heating curves show temperature changes as a substance is heated, including plateaus where phase changes occur (e.g., melting, boiling).

Phase Diagrams

Phase diagrams plot pressure versus temperature and show the conditions under which a substance exists as a solid, liquid, or gas.

• Lines (boundaries) separate different phases.

• The triple point is where all three phases coexist.

• The critical point marks the end of the liquid-gas boundary.

Example: Water and carbon dioxide have distinct phase diagrams with different triple and critical points.

Crystalline Solids: Types and Properties

Molecular Solids

• Composed of molecules held together by intermolecular forces (dispersion, dipole-dipole, hydrogen bonding).

• Generally have low to moderately low melting points.

Ionic Solids

• Composed of cations and anions held together by strong electrostatic (Coulombic) interactions.

• High melting points and are often hard and brittle.

Atomic Solids

• Composed of individual atoms.

• Nonbonding atomic solids (e.g., solid noble gases) are held together by weak dispersion forces and have very low melting points.

• Network covalent solids (e.g., diamond, graphite) are held together by covalent bonds throughout the structure, resulting in high melting points and hardness.

Summary Table: Types of Crystalline Solids

Type	Constituent Particles	Forces	Examples	Melting Point
Molecular	Molecules	Dispersion, dipole-dipole, H-bonding	Ice, CO2	Low to moderate
Ionic	Cations and anions	Ionic (Coulombic)	NaCl, CaF2	High
Atomic (Nonbonding)	Atoms	Dispersion	Xe (solid)	Very low
Atomic (Network Covalent)	Atoms	Covalent bonds	Diamond, graphite	Very high

Additional info: These notes are based on lecture slides and textbook content for a General Chemistry course, focusing on the properties and behavior of liquids, solids, and the intermolecular forces that govern them.