Chapter 11: Liquids and Intermolecular Forces

Chapter 11: Liquids and Intermolecular Forces

11.1 Molecular Comparison of States of Gases, Liquids, and Solids

The physical state of a substance is determined by the balance between the kinetic energy of its particles and the energy of attraction between them. These factors influence whether a substance exists as a gas, liquid, or solid under given conditions.

• Kinetic energy keeps particles apart and moving.

• Attractive forces bring molecules close together.

• Temperature is directly related to the average kinetic energy of particles.

• Decreasing temperature brings particles closer together, favoring the liquid or solid state.

Table 11.1: Characteristic Properties of the States of Matter

	Gas	Liquid	Solid
Shape and Volume	Assumes both volume and shape of its container	Assumes shape of portion of container it occupies	Retains own shape and volume
Expansion	Expands to fill its container	Does not expand to fill its container	Does not expand to fill its container
Compressibility	Is compressible	Is virtually incompressible	Is virtually incompressible
Flow	Flows readily	Flows readily	Does not flow
Diffusion	Diffusion within a gas occurs rapidly	Diffusion within a liquid occurs slowly	Diffusion within a solid occurs extremely slowly

11.2 Intermolecular Forces

Intermolecular forces are the attractions between molecules, which are generally much weaker than the intramolecular (covalent) bonds holding atoms together within a molecule. These forces are responsible for many physical properties of substances, such as boiling and melting points, viscosity, surface tension, and capillary action.

• Intermolecular attractions are weaker than chemical bonds.

• Hydrogen bonds are not true chemical bonds but are a type of strong intermolecular force.

Table 11.2: Melting and Boiling Points of Representative Substances

Force Holding Particles Together	Substance	Melting Point (K)	Boiling Point (K)
Ionic bonds	Lithium fluoride (LiF)	1118	1949
Metallic bonds	Beryllium (Be)	1560	2742
Covalent bonds	Diamond (C)	3600	4300
Dispersion forces	Nitrogen (N2)	63	77
Dipole-dipole interactions	Hydrogen chloride (HCl)	158	188
Hydrogen bonding	Hydrogen fluoride (HF)	190	293

Types of Intermolecular Forces Between Neutral Molecules

• Dispersion forces (London dispersion forces or induced dipole-induced dipole interactions) – weakest

• Dipole–dipole forces

• Hydrogen bonding (a special, strong dipole–dipole force)

Note: Dispersion and dipole–dipole forces are collectively called van der Waals forces.

Dispersion Forces

Dispersion forces arise from temporary fluctuations in electron distribution, which induce temporary dipoles even in nonpolar molecules. The ease with which an electron cloud can be distorted is called polarizability.

• The more easily an electron cloud is distorted, the stronger the dispersion forces.

• Factors affecting dispersion forces:

  • Number of electrons (more electrons, stronger forces)

  • Size/molecular weight (larger, stronger forces)

  • Shape (more compact, weaker forces for similar masses)

• Substances that are easy to polarize have lower boiling points, indicating weaker intermolecular forces.

• Boiling points of noble gases and halogens increase with molecular weight due to stronger dispersion forces.

Dipole–Dipole Interactions

Dipole–dipole forces occur between polar molecules, which have regions of partial positive and negative charge (dipoles). The positive end of one molecule is attracted to the negative end of another.

• For molecules of similar mass and size, increased polarity leads to higher boiling points.

• When comparing molecules of similar size and shape, dipole–dipole interactions dominate over dispersion forces.

• For much larger molecules, dispersion forces may dominate.

Hydrogen Bonding

Hydrogen bonding is a particularly strong type of dipole–dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms (N, O, or F). The nearly bare hydrogen nucleus interacts strongly with lone pairs on nearby electronegative atoms.

• Hydrogen bonds are responsible for unusually high boiling points in compounds like H2O, HF, and NH3.

• Hydrogen bonding in water leads to lower density in ice compared to liquid water, causing ice to float.

Ion–Dipole Forces

Ion–dipole interactions occur in solutions containing ions and polar molecules. These forces are crucial for the dissolution of ionic compounds in polar solvents (e.g., NaCl in water).

Determining Intermolecular Forces in a Substance

Type of Interaction	Atoms (e.g., Ne, Ar)	Nonpolar molecules (e.g., Br2, CH4)	Polar molecules without OH, NH, or HF (e.g., HCl, CH2Cl2)	Polar molecules with OH, NH, or HF (e.g., H2O, NH3)	Ionic solids in polar liquids (e.g., NaCl in H2O)
Dispersion forces	✓	✓	✓	✓	✓
Dipole–dipole interactions			✓	✓
Hydrogen bonding				✓
Ion–dipole interactions					✓

Note: All substances exhibit dispersion forces. The strongest force present dictates the extent of intermolecular attractions.

Generalizations About Relative Strengths of Intermolecular Forces

1. For molecules with similar molar masses and shapes, dispersion forces are roughly equal.

2. For molecules with very different molar masses and no hydrogen bonding, dispersion forces determine the stronger attractions.

3. Complex molecules may have a combination of intermolecular forces, and their relative importance must be weighed.

11.3 Select Properties of Liquids

Several properties of liquids are influenced by intermolecular forces, including boiling point, melting point, viscosity, surface tension, capillary action, heats of vaporization, and heats of fusion.

Viscosity

• Viscosity is the resistance of a liquid to flow.

• It depends on the ease with which molecules can move past each other.

• Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Substance	Formula	Viscosity (kg/m·s) at 20°C
Hexane	CH3CH2CH2CH2CH2CH3	3.26 × 10−4
Heptane	CH3CH2CH2CH2CH2CH2CH3	4.09 × 10−4
Octane	CH3CH2CH2CH2CH2CH2CH2CH3	5.42 × 10−4
Nonane	CH3CH2CH2CH2CH2CH2CH2CH2CH3	7.11 × 10−4
Decane	CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3	1.42 × 10−3

Surface Tension

• Surface tension is the energy required to increase the surface area of a liquid due to intermolecular forces.

• It causes water to "bead up" on nonpolar surfaces and allows small insects to walk on water.

Cohesion and Adhesion

• Cohesive forces are intermolecular forces that bind similar molecules together (e.g., water molecules).

• Adhesive forces are intermolecular forces that bind a substance to a surface (e.g., water to glass).

• These forces are important in capillary action, which enables liquids to rise in narrow tubes and is essential for processes like water transport in plants and absorption in paper towels.

Capillary Action

• The rise of liquids up narrow tubes is called capillary action.

• Water has stronger adhesive forces with glass, resulting in a concave meniscus.

• Mercury has stronger cohesive forces with itself, resulting in a convex meniscus.

11.4 Phase Changes

A phase change is the conversion from one state of matter to another, involving the addition (endothermic) or release (exothermic) of energy.

• Common phase changes: melting/freezing, vaporizing/condensing, subliming/depositing.

Energy Change Accompanies Phase Change

• Heat of fusion ($q_{fus}$): Energy required to change a solid at its melting point to a liquid.

• Heat of vaporization ($q_{vap}$): Energy required to change a liquid at its boiling point to a gas.

• Heat of sublimation ($q_{sub}$): Energy required to change a solid directly to a gas.

Heating Curves

• A heating curve is a graph of temperature versus heat added.

• During a phase change, temperature remains constant while the substance absorbs or releases heat.

• The enthalpy change (heat, $q$) is calculated as:

  • For temperature change: $q = m c \, \Delta T$

  • For phase change: $q = n \Delta H_{fus}$ or $q = n \Delta H_{vap}$

Supercritical Fluids

When a gas is compressed above its critical temperature and critical pressure, it forms a supercritical fluid, which has properties of both gases and liquids. Supercritical CO2 is widely used as a solvent.

Substance	Critical Temperature (K)	Critical Pressure (atm)
Nitrogen, N2	126.1	33.5
Argon, Ar	150.9	48.0
Oxygen, O2	154.4	49.7
Methane, CH4	190.0	45.4
Carbon dioxide, CO2	304.3	73.0
Phosphine, PH3	324.4	64.5
Propane, CH3CH2CH3	370.0	42.5
Hydrogen sulfide, H2S	373.5	88.9
Ammonia, NH3	406.5	111.5
Water, H2O	647.6	217.7

11.5 Vapor Pressure

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid at a given temperature. As temperature increases, more molecules have enough energy to escape the liquid phase, increasing vapor pressure.

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

• The boiling point is the temperature at which vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which vapor pressure is 1 atm (760 torr).

• The relationship between vapor pressure and temperature is given by the Clausius-Clapeyron equation:

  • $\ln P = -\frac{\Delta H_{vap}}{R} \left( \frac{1}{T} \right) + C$

Summary Table: Types of Intermolecular Forces

Type of Force	Relative Strength	Examples
Dispersion (London)	Weakest	N2, noble gases
Dipole–dipole	Intermediate	HCl, CH2Cl2
Hydrogen bonding	Strongest (of neutral molecules)	H2O, HF, NH3
Ion–dipole	Very strong	Na+ in H2O

Example: Water (H2O) exhibits all types of intermolecular forces except ion–dipole (unless ions are present). Its high boiling point and surface tension are due to hydrogen bonding.