Liquids, Solids, and Intermolecular Forces: Structured Study Notes

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 substances like 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 and strong intermolecular interactions.

• Example: Gecko adhesion is a biological application of intermolecular forces.

Properties of the Three Phases of Matter

Matter exists in three primary phases: gas, liquid, and solid, each with distinct properties.

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 in a container (solids).

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

Three Phases of Water

• Density and molar volume: Ice and liquid water have much higher densities than steam.

• Ice has a lower density than liquid water, which is unusual and crucial for life.

• Example: Ice floats on water due to its lower density.

Phase	Temperature (°C)	Density (g/mL)	Molar Volume (mL)
Gas (steam)	100	0.0006	30.5
Liquid (water)	20	0.998	18.0
Solid (ice)	0	0.917	19.6

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.

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 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; some kinetic energy overcomes attractive forces, but not enough to escape.

Kinetic Energy

• Increasing kinetic energy increases particle motion and freedom.

• Average kinetic energy is proportional to temperature: $KE_{avg} = 1.5 kT$

Attractive Forces

• Particles are attracted by electrostatic forces.

• Strength varies by particle type; stronger forces resist motion.

Kinetic-Molecular Theory of the Phases

• Gases: No attractions; kinetic energy dominates.

• Solids: No translational/rotational motion; strong attractive forces.

• Liquids: Limited translational/rotational freedom; intermediate attractive forces.

Phase Changes

• Changing state requires altering kinetic energy or limiting freedom.

• Melting: Particles gain enough energy to partially overcome attractions.

• Boiling: Particles gain enough energy to completely overcome attractions.

• Condensation: Achieved by decreasing temperature or increasing pressure.

Intermolecular Attractions

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

• Stronger attractions lead to higher boiling and melting points.

Why Are Molecules Attracted to Each Other?

• Attractions arise from opposite charges: ion-ion, polar-polar, and hydrogen bonding.

• Even nonpolar molecules can have temporary charges.

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

Trends in the Strength of Intermolecular Attraction

• Stronger attractions require more energy to separate molecules.

• Boiling point is a measure of intermolecular force strength.

Kinds of Attractive Forces

• Dispersion forces: Temporary polarity due to electron distribution.

• 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

• Result from temporary dipoles due to electron fluctuations.

• 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 and more surface contact increase dispersion force strength.

Effect of Molecular Size and Shape on Dispersion Forces

Noble Gas	Molar Mass (g/mol)	Boiling Point (K)
He	4.00	4.2
Ne	20.18	27
Ar	39.95	87
Kr	83.80	120
Xe	131.30	165

• Boiling point increases with molar mass and surface area.

• Example: n-Pentane (large area) has stronger dispersion forces than neopentane (small area).

Dipole-Dipole Attractions

• Polar molecules have permanent dipoles, increasing boiling and melting points.

• Stronger than dispersion forces.

Effect of Dipole-Dipole Attraction on Boiling and Melting Points

Name	Formula	Molar Mass (g/mol)	Boiling Point (°C)	Melting Point (°C)
Formaldehyde	CH2O	30.03	-19.5	-92
Ethane	C2H6	30.07	-88	-172

Dipole Moment and Boiling Point

• Boiling point increases with dipole moment for molecules of similar size.

Example: Dipole-Dipole Forces

• CO2: No dipole forces (linear, nonpolar).

• CH2Cl2: Dipole forces present (polar, tetrahedral).

• CH4: No dipole forces (nonpolar, tetrahedral).

Attractive Forces and Solubility

• Solubility depends on the attractive forces between solute and solvent.

• Like dissolves like: Polar dissolves in polar; nonpolar in nonpolar.

• Hydrophilic groups: OH, CHO, C=O, COOH, NH2, Cl.

• Hydrophobic groups: C-H, C-C.

Immiscible Liquids

• Pentane (nonpolar) and water (polar) do not mix due to stronger water-water attractions.

Hydrogen Bonding

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

• Exposes the H proton, creating a strong center of positive charge.

• Hydrogen bonds are much stronger than other intermolecular forces, but weaker than covalent bonds.

H-Bonding in Water and Ethanol

• Hydrogen bonds in water and ethanol lead to higher boiling points and unique properties.

• Example: Water's high boiling point and solvent capabilities are due to hydrogen bonding.