CHAPTER 11 (MOORE)
STATES OF MATTER AND INTERMOLECULAR FORCES
This chapter deals the forces of intermolecular attraction in pure liquids and solids. These forces
are responsible for the physical state of compounds at specified temperature and pressures. We will
consider vapor pressure, boiling point and melting point as surrogates for intermolecular forces, and
we also will look at changes of state and the energies that accompany them. Phase diagrams will be
important representations of these changes.
“The properties of liquids and solids … can be understood on the molecular, nanoscale level”
…their “atoms, molecules or ions are close enough to have strong interactions with each other”
(Moore, p. 489)
“We can use knowledge of intermolecular forces that exist in various substances to make
predictions about trends in such properties as melting points, boiling points and enthalpies of
vaporization.” (Hill, p. 453)
Molecular Forces
Intramolecular forces are forces between atoms in the same molecule (bonds) that determine
individual molecular properties, e.g. molecular shapes and dipole moments.
Intermolecular forces are forces between atoms in different molecules that determine the
macroscopic physical properties of liquids and solids, e.g. melting points and boiling points.
Intermolecular forces are relatively weak.
Intramolecular forces are relatively strong …
States of Matter
Gases - least ordered, intermolecular forces are not very important
Liquids - somewhat ordered, intermolecular forces are very important
Solids - most ordered, intermolecular forces are very important
Intermolecular Forces … are forces between molecules.
• They determine melting points, freezing points, and other physical properties.
• Types of intermolecular forces include:
• dispersion forces
• Dipole–dipole forces
• hydrogen bonding
Dispersion Forces
• exist between any two particles.
• Are also called London forces (after Fritz London, who offered a theoretical explanation of
these forces in 1928).
• Dispersion forces arise because the electron cloud is not perfectly uniform.
• Tiny, momentary dipole moments can exist even in nonpolar molecules.
Dispersion Forces Illustrated
At a given instant, electron density, even in a nonpolar molecule, is not perfectly uniform. Then a
region of momentary higher electron density attains a small (–) charge … and the other end of the
molecule is slightly (+). When another nonpolar molecule approaches … the first molecule induces
a tiny dipole moment … in this molecule. The opposite charges (+ ) (−) cause a momentary
attraction between the two molecules.
Strength of Dispersion Forces
• Dispersion force strength depends on polarizability: the ease with which the electron cloud
is distorted by an external electrical field.
• The greater the polarizability of molecules, the stronger the dispersion forces between them.
• Polarizability in turn depends on molecular size and shape.
• Heavier molecule => more electrons => a more- polarizable molecule.
• As to molecular shape …
Molecular Shape and Polarizability
Polarizability – a measure of how easily the shape of the electron cloud in an atom or a molecule is
distorted by an external electric field, i.e. how easily a dipole can be produced in the atom or
molecule.
Long skinny molecules … can have greater separation of charge along their length and stronger
forces of attraction, meaning … higher boiling points.
Compact molecules … have less possible separation of charge … with giving weaker dispersion
forces and lower boiling points.
Dipole–Dipole Forces
Polar molecules have positively charged “ends” (δ+) and negatively charged “ends” (δ–).
• When molecules come close to one another, repulsions occur between like-charged regions
of dipoles. Opposite charges tend to attract one another.
• The more polar a molecule, the more pronounced is the effect of dipole–dipole forces on
physical properties.
•
In dipole-dipole interactions, opposites attract!
Predicting Physical Properties of Molecular Substances
• Dispersion forces become stronger with increasing molar mass and elongation of molecules.
In comparing nonpolar substances, molar mass and molecular shape are the essential factors.
• Dipole–dipole and dipole-induced dipole forces are found in polar substances. The more
polar the substance, the greater the intermolecular force is expected to be.
• Because they occur in all substances, dispersion forces must always be considered. Often
they predominate!
See Examples 11.1 and 11.2, Moore, p. 491
Strategy. Draw molecular structures for compounds listed in the table and find which one is most
“like” the structure for the compound in question.
Homology: a series of compounds whose formulas and structures vary in a regular manner also
have properties that vary in a predictable manner. This principle is called homology.
Example: both densities and boiling points of the straight-chain alkanes increase in a continuous
and regular fashion with increasing numbers of carbon atoms in the chain.
Trends result from the regular increase in molar mass, which produces a fairly regular increase in
the strength of dispersion forces.
Hydrogen Bonds: a hydrogen bond is an intermolecular force in which:
(1) a hydrogen atom that is covalently bonded to a small, electronegative, nonmetal atom (Y) in one
molecule … and (2) is simultaneously attracted to a small, electronegative, nonmetal atom (Z) of a
neighboring molecule.
Y ––– H - - - Z ~~~~
When Y and Z are small and highly electronegative (N, O, F) … this force is called a hydrogen
bond; a special, strong type of dipole–dipole force.
Hydrogen Bonds in Water
Hydrogen bonding usually occurs between
molecules, but, sometimes, hydrogen bonding occurs
within the molecule (see Fig. 11.22, Hill, p. 456)
Intermolecular hydrogen bonds give proteins their secondary shape, forcing the protein molecules
into particular orientations, like a folded sheet.
Liquids and Intermolecular Forces
The behavior and properties of liquids are related to intermolecular forces.
Adhesive forces are intermolecular forces between unlike molecules.
Cohesive forces are intermolecular forces between like molecules.
Viscosity is a measure of the resistance to flow of a liquid, e.g. honey.
Surface tension (γ) is work required to extend a liquid surface and is usually expressed in J/m2.
The molecules at the surface must be separated from one another to create more surface.
Liquids “spread” because adhesive forces are exceed cohesive forces.
When liquids “bead up,” which forces are stronger, adhesive or cohesive? See Moore, p. 490
Changes of State for Gases: Vaporization and Condensation
Vaporization is the conversion of a liquid to a gas; condensation is the conversion of a gas to a
liquid.
The molar enthalpy of vaporization (ΔHvap) is the quantity of heat energy that must be absorbed
to vaporize exactly 1 mol of liquid at a constant temperature.
The molar enthalpy of condensation (ΔHcon) is the quantity of heat energy that is given off when
exactly 1 mol gas condenses to a liquid a constant temperature.
Enthalpy of Vaporization Problems
See Problem-Solving Example 11.2 and Concept Exercise 11.7, Moore, pp. 497-498
Vapor Pressure
The vapor pressure of a liquid is the partial pressure exerted by the vapor above the liquid when it
is in ‘dynamic equilibrium” with the liquid at a constant temperature.
Liquid ' Vapor ( is vaporization and is condensation)
Liquid-Vapor Equilibrium and Vapor Pressure
Equilibrium occurs when vaporization/condensation rates become the same, and, more generally,
when the rates of any two opposing processes are equal.
Vapor Pressure of Water
Note the temperature at which the vapor pressure just equals 1.00 atm (760 mmHg) … this is the
“normal boiling point”
Boiling Point and Critical Point: Terminology
Boiling point: the temperature at which vapor pressure of a liquid equals the external pressure.
Normal boiling point: boiling point at 1 atm external pressure
Critical temperature (Tc): the highest temperature at which a pure substance can exist as a liquid
The critical pressure, Pc, is the vapor pressure at the critical temperature.
The condition corresponding to a temperature of Tc and a pressure of Pc is called the critical point.
Phase Changes Involving Solids
Melting (fusion): transition of solid liquid.
Melting point: temperature at which melting occurs (same as the freezing point)
Enthalpy of fusion, ΔH(fusion), quantity of heat required to melt a set amount (e.g. one mole) of
solid.
Sublimation: transition of solid vapor. (Ex. why ice cubes slowly “disappear” in your freezer)
Enthalpy of sublimation, ΔH(subln), the sum of the enthalpies of fusion and vaporization.
Triple point: all three phases — solid, liquid, vapor — are in simultaneous equilibrium.
Heating Curve for Water (see Moore, Fig. 11.13, p. 502)
Phase Diagrams (Moore, p. 503): graphical representations that show the temperatures and
pressures under which a substance exists as a solid, liquid, a gas, or some combination of these in
equilibrium.
See Example 11.5 Moore, pp. 505 and Problem-Solving Practice, p. 506.
Types of Solids
Amorphous solids have no significant long-range order, while crystalline solids have
atoms/ions/molecules arranged in a regular pattern.
Types of crystalline solids:
1. Molecular solids, containing molecules held to one another by dispersion/dipole–dipole/
hydrogen bonding forces.
2. Network solids, have a network of covalent bonds that extend throughout the solid, holding it
firmly together. The allotropes of carbon (Graphite and diamonds) are good examples of
network solids.
3. Ionic solids … here, ionic (+ −) attractions are the source of “intermolecular” forces
Since there are no molecules in an ionic solid, there can’t be any intermolecular forces – so the
attractions are electrostatic inter-ionic attractions, where the lattice energy (see Chapter 9) is a
measure of the strength of inter-ionic attraction.
Rule: attractive forces between pairs of oppositely charged ions increases: (1) as the charges on
the ions increase and (2) as the ionic radii decrease, and the lattice energy increases accordingly.
Toolkit for Ionic Compounds: Monatomic Ions
Recall the charges on monatomic Group A ions …
Interionic Forces of Attraction (increase with increased charge)
Melting point of NaCl (Na 1+ and Cl1- is about 801 oC, while Mg2+ and O2– have much stronger
forces of attraction for one another … melting point of MgO is about 2800 oC.
Crystal Lattices
Three-dimensional views are used to describe crystals, and the “repeating unit” of the lattice is
called the unit cell. Types of unit cell include hexagonal, rhombic, and cubic. The three types of
cubic unit cells: simple cubic, body-centered cubic (bcc), face-centered cubic (fcc).
4. Metallic solids (metals).