Liquids, Solids, and Intermolecular Forces

Liquids, Solids, and Intermolecular Forces

Introduction

This chapter explores the properties of liquids and solids, focusing on the intermolecular forces that govern their behavior. Understanding these forces is essential for explaining phenomena such as boiling, melting, solubility, and the unique properties of water.

Climbing Geckos: An Application of Intermolecular Forces

• Geckos can adhere to almost any surface due to intermolecular attractive forces.

• Millions of tiny hairs (setae) on their feet branch out and flatten (spatulae), allowing close contact with surfaces and maximizing intermolecular attractions.

• Example: Gecko adhesion is a practical demonstration of the strength of intermolecular forces.

Properties of the Three Phases of Matter

Comparison of States

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

• Definite shape: Keeps shape when placed in a container (solids).

• Indefinite shape: Takes the shape of the container (liquids and gases).

Three Phases of Water

Phase	Temperature (°C)	Density (g/mL)	Molar Volume (mL)
Gas (steam)	100	0.590 × 10−3	30.5
Liquid (water)	20	0.998	18.0
Solid (ice)	0	0.917	19.6

• Density and molar volume of ice and liquid water are much closer to each other than to steam.

• Ice is less dense than liquid water, which is unusual and vital for life.

Degrees of Freedom

• Translational freedom: Ability to move from one position to another.

• Rotational freedom: Ability to reorient direction in space.

• Vibrational freedom: Ability to oscillate about a point in space.

Kinetic-Molecular Theory

• The state of a material depends on:

  1. The amount of kinetic energy the particles possess.

  2. The strength of attraction between the particles.

• These factors compete to determine the phase.

States and Degrees of Freedom

• Gases: Complete freedom of motion; kinetic energy overcomes attractive forces.

• Solids: Locked in place; only vibrational motion; attractive forces dominate.

• Liquids: Limited freedom; enough kinetic energy to partially overcome attractions.

Kinetic Energy

• Increasing kinetic energy increases particle motion and freedom.

• Average kinetic energy is directly proportional to temperature:

$KE_{avg} = 1.5kT$

Attractive Forces

• Particles are attracted by electrostatic forces.

• Strength varies with particle type and affects resistance to movement.

• Kinetic-Molecular Theory of the Phases

• Gases: No attractions; kinetic energy dominates.

• Solids: No translational/rotational motion; attractions dominate.

• Liquids: Limited motion; attractions partially overcome by kinetic energy.

Phase Changes

• Changing state requires altering kinetic energy or limiting freedom.

• Solids melt and liquids boil as particles gain enough kinetic energy to overcome attractions.

• Gases condense by lowering temperature or increasing pressure.

Intermolecular Attractions

• Moderate to strong attractions at room temperature result in solids or liquids.

• Stronger attractions lead to higher boiling and melting points.

Why Are Molecules Attracted to Each Other?

• Attractions arise from opposite charges (ion-ion, dipole-dipole, hydrogen bonding, temporary charges).

• Larger charge = stronger attraction; longer distance = weaker attraction.

• Intermolecular forces are weaker than covalent bonds.

Trends in the Strength of Intermolecular Attraction

• Stronger attractions require more energy to separate molecules.

• Boiling a liquid requires overcoming intermolecular (not covalent) forces.

• Higher boiling point indicates stronger intermolecular forces.

Kinds of Attractive Forces

• Dispersion forces: Temporary polarity due to electron distribution fluctuations.

• Dipole-dipole attractions: Permanent polarity due to molecular structure.

• Hydrogen bonds: Strong dipole-dipole attraction when H is bonded to O, N, or F.

Dispersion Forces (London Forces)

• Result from temporary dipoles in atoms/molecules.

• Present in all molecules and atoms; weakest intermolecular force.

Size of the Induced Dipole

• Magnitude depends on electron polarizability and molecular shape.

• Larger molar mass = more electrons = stronger dispersion forces.

• More surface contact = stronger dispersion forces.

Effect of Molecular Size and Shape on Dispersion Forces

• Noble gases: Higher molar mass leads to higher boiling points due to stronger dispersion forces.

• Linear molecules (e.g., n-pentane) have stronger dispersion forces than spherical ones (e.g., neopentane).

Dipole-Dipole Attractions

• Polar molecules have permanent dipoles, increasing boiling and melting points compared to nonpolar molecules of similar size.

• Stronger than dispersion forces.

Hydrogen Bonding

• Occurs when H is bonded to highly electronegative atoms (O, N, F).

• Hydrogen bonds are much stronger than other dipole-dipole interactions but weaker than covalent bonds.

• Responsible for high boiling points of substances like water and ethanol.

Attractive Forces and Solubility

• "Like dissolves like": Polar substances dissolve in polar solvents; nonpolar in nonpolar solvents.

• Hydrophilic groups (e.g., OH, COOH) increase solubility in water; hydrophobic groups (e.g., C–H, C–C) decrease it.

Immiscible Liquids

• Nonpolar and polar liquids (e.g., pentane and water) do not mix due to stronger attractions within each liquid than between them.

Summary Table: Types of Intermolecular Forces

Type	Relative Strength	Occurs Between
Dispersion (London)	Weakest	All molecules/atoms
Dipole-Dipole	Intermediate	Polar molecules
Hydrogen Bonding	Strongest (in pure substances)	H bonded to O, N, or F

Additional info: Ion-dipole forces (not detailed here) are even stronger and occur in mixtures of ionic compounds and polar molecules.