Chapter 11 – Intermolecular Forces and Liquids and Solids
Notes based on Expectations for Learning:
1. List the states of matter and contrast differences in their kinetic energy, shape, volume, density,
compressibility, ability to flow, and molecular motion.
State
Submicroscopic Model
Plasma
+
+
-
+ -
-
Changes from
one state to
another
From other states:
Ionization
(endothermic)
Kinetic
Energy
Shape
Volume
Molecular
Motion
Compress
-ibility
Ability
to flow
Density
Very
High
Indefinite
Indefinite
Particles move
very rapidly
and randomly
in straight-line
paths. They
collide with
sufficient
energy to
separate
electrons for
the atoms.
Compressible
Fluid
Extremely
low
To liquid:
Condensation
(exothermic)
High
Indefinite
Indefinite
Particles feel
little attraction
for each other
and move
randomly in
straight-line
paths
Compressible
Fluid
Very low
To solid:
Freezing
(exothermic)
To gas:
Boiling or
Evaporation
(endothermic)
Moderate
Indefinite
Definite
The particles
are held close
together by
attractive
forces but
vibrate about a
moving point
and can slip
and slide past
one another.
Incompressible
Fluid
High
To Liquid:
Melting
(exothermic)
To gas:
Sublimation
(endothermic)
Low
Definite
Particles are
rigidly held in
an ordered
arrangement
and.vibrate
about a fixed
point. They
are locked in
position
relative to one
another.
Incompressible
Not
fluid
Very high
+
+
Plasma
Gas
To solid:
Deposition
(endothermic)
Liquid
Solid
Have zero
dimensions of
organizati
on
Definite
Crystals
have 3-D
arrangement
2. Define interparticle forces and intraparticle forces and differentiate between them.
Interparticle are the forces of attraction between different molecules.
Examples: Ionic bonds, metallic bonds, Van der Waals forces
Intraparticle forces are the forces within a molecule. In other words, they are the forces that hold
atoms together to form a molecule.
Example: Covalent bonds
Note: The terms intermolecular and intramolecular forces are used exclusively for molecules whereas the terms interparticle forces and
intraparticle forces are more general and may include the forces between ions and metal atoms.
3. Define, compare, and contrast following sets of forces between particles; ion-dipole, hydrogen bonds,
dipole-dipole forces, ion-induced dipole, dipole-induce dipole, and London dispersion forces.
AKA: Van der Waals Forces
FORCES OF
ATTRACTION
STRONG
IONIC
COVALENT
WEAK
METALLIC
ION-DIPOLE
DIPOLE-DIPOLE
DIPOLEINDUCED
DIPOLE
LONDON
DISPERSION
FORCES
ION-INDUCED DIPOLE
HYDROGEN BONDS
ion-dipole
hydrogen
bonds
dipole-dipole
(polar
molecule- polar
molecule)
The attraction
between an ion
and the
oppositely
charged end of a
permanent
dipole.
(See the
definition
below)
These are a
strong type of
dipole-dipole
forces
The force of
attraction
between the
oppositely
charges ends of
adjacent polar
molecules.
ion-induced
dipole*
(ion-nonpolar)
dipole-induce
dipole*
(polarnonpolar)
London
dispersion
forces*
(Found in all
molecules)
The attraction
The attraction
The fleeting
between an ion
between a
force of
and a temporary permanent
attraction
dipole created in dipole (polar
between
a nonpolar
molecule) and a momentary
molecule due to temporary
aggregations of
the presence of
dipole created in electrons in one
the said ion.
a nonpolar
molecule and
molecule due to the protons in
the presence of
another.
the said dipole.
Decreasing Strength (Debye units)
4. Define and describe the formation of dispersion forces and how the number of electrons in a
molecule affects its polarizability.
Dispersion forces are created when electrons in one molecule temporarily “clump” together due to their
random motion. This concentration of negative charge causes electrons in a second molecule to be repelled. As
a result, the protons of the second molecule have a thinner “layer” of electrons covering them. Therefore, for a
brief moment in time, the electrons in the first molecule feel a stronger pull from protons in the second
molecule. Dispersion forces are sometimes called temporary dipoles because of their brief existence.
*Note: The tendency for electrons to migrate within a particle is called polarizability and increases with the
number of electrons in the particle.
5. Define hydrogen bonding and compare its strength to other dipole-dipole forces.
The force of attraction between a hydrogen atom, bonded to a
highly electronegative element (N, O, F) in one molecule, and a
highly electronegative element in another molecule. Hydrogen
bonds are stronger than ordinary dipole-dipole forces of
attraction due to the nature of a hydrogen atom. Hydrogen
contains only one proton and one electron. When hydrogen is
bonded to a highly electronegative atom, its single electron is
pulled toward the other atom, leaving practically a bare proton.
The strength of the attraction of this nearly bare proton for
electrons is less than that of a cation with a +1 charge but greater than that of the positive end of an ordinary
polar molecule. Substances that contain hydrogen bonds have boiling points that are higher than what is
expected. Below and right: Trends in the boiling points of covalent binary hydrides show the anomalous
behavior of NH3, H2O, and HF due to their hydrogen-bonding intermolecular forces.
In the diagrams to the right and above, liquid water contains a vast three-dimensional network of hydrogen
bonds resulting from the attraction between positively polarized hydrogen atoms and electron pairs on
negatively polarized oxygen atoms. Each oxygen can form two hydrogen bonds, represented by dotted lines.
The two long molecular strands that are coiled
around each other in DNA are held together by
hydrogen bonding between
A-T and C-G nitrogen base pairs.
6. Describe the arrangement of the particles in the liquid state.
Particles in the liquid sate are haphazardly arranged, and freeflowing.
7. Define surface tension and describe how intermolecular forces affect it.
Surface tension is a cohesive force present in like particles such as water
and mercury. Surface tension is the apparent skin-like quality on the
surface of a liquid due to the presence of unbalanced forces on surface
particles. There is a net force pulling surface particles toward the center
of the liquid because surface particles are attracted to neighbors located
only to the side and toward the center, whereas center particles are completely surrounded by neighbors.
As interparticle forces increase, so does surface tension.
Mercury’s particles are attracted to each other with the relatively
strong metallic bonds. It has one of the highest surface tensions of
all liquids. Hydrogen bonding in water gives it a fairly high surface
tension as well. This allows spiders and insects to walk on water as
in the photo above.
8. Define adhesion and capillary rise. Then describe the forces of
attraction that affect them.
Adhesion is the attraction between unlike substances. Capillary
rise (a.k.a. capillarity) is the upward motion of a liquid in a small
tube. This behavior is similar to the concave meniscus of water as
seen in a graduated cylinder. Attraction between the liquid and its
container cause the liquid particles to creep up the sides. When the
container is very small, such as in a capillary tube with a 2-4 mm
diameter, the liquid creeps up the tube, resulting in the
phenomenon known as capillary rise.
9. Define viscosity and describe how temperature affects it.
Viscosity is the resistance of a substance to flow. Liquids with high viscosity flow slowly, such as
molasses. Liquids with low viscosity flow quickly, such as water.
Factors that affect viscosity
Temperature
As temperature increases, viscosity decreases. Additional kinetic energy
breaks intermolecular forces of attraction, allowing more movement of
the molecules.
Type of forces of attraction
As the relative strength of intermolecular forces of attraction increases,
viscosity increases. Substances with stronger intermolecular forces flow
more slowly than weaker ones.
Size of the molecules
As the size of the molecules increases , viscosity increases. Larger
molecules have more intermolecular forces of attraction per molecule
than smaller molecules. This allows more contact between molecules,
and they stick together more than smaller molecules with the same type
of intermolecular forces of attraction.
10. Describe the arrangement of water molecules in the solid state and why it is less dense than liquid
water.
In the solid state, each water molecule is hydrogen-bonded to four other water molecules in a highly
ordered three-dimensional crystal. During melting, energy added to
the system is used to break the hydrogen bonds. As the hydrogen
bonds break, the 3-D structure collapses, causing the particles to move
in closer to one another, thus increasing water’s density. Once a
liquid, additional energy increases the kinetic energy of the molecules
(and the temperature from 0°C to 3.98°C) AND breaks the many more
of the hydrogen bonds present, resulting in water’s highest density at
3.98°C. Above this temperature, additional energy creates more
forceful and frequent collisions between the particles, causing the
molecules to spread out. The result is a lowering of the density of
water from 3.98°C to 100°C when it begins to boil (at 1atm).
11. Differentiate between crystalline solids, smectic liquid crystals, nematic liquid crystals, amorphous
solids, and liquids in terms of their arrangement, their behavior during melting, and ability to flow.
Crystalline solids
Smectic liquid crystals
Nematic liquid Amorphous solids
(Super cooled Liquids)
crystals
Liquids
3-D Order
Made up of repeating units
called UNIT CELLS
(See below)
Melt completely at one temp.
2-D Order
1-D Order
0-D Order
Haphazard Arrangement
(Considerable disorder)
0-D Order
Haphazard
Arrangement
Melt over a temp. range
Melt over a temp.
range
Melt over a temp. range
Melt all at once
Do not flow.
Flow slowly (like jelly)
Flow slowly
Flow very slowly
(sometimes, over a
period of years)
Viscosity varies
w/ IMF’s of
attraction.
12. Define unit cell and crystal lattice and describe their relationship to one another.
A unit cell is the simplest repeating unit in the arrangement of a crystal. (a brick)
The arrangement of the units in a crystal is determined by the bonds between particles. Bonding determines
properties of crystals.
A crystal (or space) lattice is a regular arrangement of unit cells. (a brick wall)
Brief History of Solids:
Nicholas Steno began the study of crystals in 1669. He
noticed the corresponding angles between faces on
different crystals of the same substance were always
the same. For Example: Table Salt, sodium chloride,
is made up of tiny cube-shaped crystals. The angles
between surfaces (faces) are always 90°.
STENO'S LAW: All crystals of the same substance
have the same corresponding angles between faces.
This is true for all crystals, regardless of the source.
For example: A 40 ton crystal of beryl (unearthed in
New Hampshire!) has the same intensive properties as
all beryl crystals.
The extensive properties of crystals vary while the
intensive properties always remain the same.
There are seven different crystal systems;
cubic, tetragonal, orthorhombic, trigonal
(rhombohedral), monoclinic, triclinic, and
hexagonal.
13. Differentiate between a hygroscopic ionic compound and a deliquescent compound.
A hygroscopic substance is an ionic compound that is strongly attracted to water molecules.
A dessciant is a hygroscopic substance that is used to remove water from the air, such as silica gel.
A deliquescent substance is a hygroscopic compound that is so strongly attracted to water that it is able to form
a liquid solution.
14. Define and calculate the internuclear distance between ions in a
crystal.
The internuclear distance is the distance between the nuclei of adjacent
atoms. This distance can be determined if the angle of incidence (θ) and
wavelength (λ) of a ray striking a crystal are known.
The Bragg Equation relates the angle,
and wavelength.
nλ = 2d sinθ
“n” is the number of wavelengths that fit into the distance BC + CD in the
diagram to the right. It is referred to as the order.
14. Calculate the number of particles per unit cell, given the edge length and density.
Particle radius (pm)
Edge length (cm)
Volume of cube
(cm3) x Density of crystal (g/cm3)
Mass (g)
Mass(amu)
Atoms
15. Calculate the length of the side of a unit cell, given the
crystal type and density of the crystal.
Crystal Type
Atoms
(g/cm3)
(volume)1/3
(pm)
Mass (amu)
Mass (g)/density
side length (cm)
particle radius
Simple Cubic
16. Differentiate between simple
cubic, body-centered cubic, and
face-centered cubic unit cells in
terms of the location
and number of the particles in the
unit cell, the number of nearest
neighbors, the percent of space
occupied by the particles, the relative
occurrence in nature, the pattern of
the layers in the crystal, and the
relationship between the edge length
and the radius of the particles.
17. Define liquid crystals and
differentiate between the different
types in terms of their spatial
arrangement.
Body-centered
cubic
Corners and center
Face-centered
cubic
Corners and faces
8 x 1/8 = 1
(8 x 1/8) + (1
center) = 2
(8 x 1/8) + (6 x 1/2
)=4
6
8
12
52 - 54%
68%
74%
Location of
Particles
number of
particles
per unit cell
Coordination
number
(# nearest
neighbors)
% of space
occupied by
particles
Occurrence
Corners only
Rare in nature
Found in metals
(especially alkali)
Pattern of layers
a, a, a
a, b, a,
Found in many
elements,
Covalent, and ionic
compounds
a, b, c, a,
Edge(a) - radius(r)
relationship
a = 2r
a = 4r3-1/2
A = 81/2 r
The information on liquid crystals that follows is an excerpt from a Web page prepared by the Chemistry Department at McGill University, Montreal,
Canada. For more information go directly to the site. The URL for the web site is:
http://barrett-group.mcgill.ca/teaching/liquid_crystal/LC02.htm
The common states of matter, namely solid, liquid, and gas differ mainly by the types of and degree of order present in the
phase. These states of matter, however, are not sufficient to characterize the structures found in all systems. As most
substances are heated, they go from a solid (usually crystalline, possessing high order) to an isotropic liquid (highly
disordered). Some substances, however, exhibit intermediate states lacking some of the order found in solids, but possessing
more order than found in liquids. These ordered fluids are called liquid crystals. Crystalline solids have positional and
orientational order. Conventional liquids have neither. Liquid crystals, on the other hand, might have no positional order, but
some orientational order (with correlations between the directions of neighboring molecules). This ordering usually persists
only for a fairly narrow temperature range, and is related to the intrinsic molecular shape. Liquid crystalline molecules, called mesogens, are usually highly anisotropic
in shape, which gives rise to the preferred orientations of nearby molecules. For illustrative purposes, let us begin by looking at some of the phases observed for rod-like
mesogens. (continued on next page)
The nematic phase has no positional order, but has orientational order. That is, the mesogens all point in the same direction, essentially expressing the molecular
anisotropy as a phase anisotropy. A given molecule's orientation is not constant, and all the mesogens do not point in exactly the same direction. For an isotropic liquid,
averaging molecular orientations gives no result, since there are as many molecules lying along one axis as another. In the nematic phase, averaging molecular
orientations gives a definite preferred direction, which is referred to as the director. It is important to remember that liquid crystals are liquids, meaning that although
there is an average order, molecules are constantly flowing, and moving, changing position and orientation.
Smectic phases have orientational order, and some degree of positional order. These phases are distinguished by the presence of layers perpendicular to the director.
The Smectic A phase has layers oriented at 90 degrees to the director. The Smectic C phase has a director tilted with respect to the layers. Below,
schematics for these phases are shown.
The diagram to the right shows how liquid crystals are used in modern
technology in devices such as calculators. Liquid crystals change in color or
opacity under varying conditions of temperature, pressure, or electric current.
18. Demonstrate an understanding of the position of cations and anions in
crystals of ionic compounds.
In unit cells of ionic crystals, the smaller cations fit in the spaces between anions.
The side of a face-centered unit cell, such as NaCl, equals the diameter of the cation
plus the diameter of the anion.
19. Define the types of crystalline solids and describe differences in the forces of attraction holding their
particles together.
KEY:
P = particles that make
up the lattice points
B = bonding present
S = substance example
C = conductivity
MP = melting point
SOLIDS
CRYSTALLINE
AMORPHOUS
P = nonmetal atoms
B = covalent
S = SiO2(glass)
MP = melt over a range (no set temp)
IONIC
MOLECULAR
P = ions
B = ionic
S = NaCl. CaF2
C = molten or in water
MP = very high
P = discrete molecules
B = dipole-dipole or
London dispersion
S = H2O, CO2
MP = low
ATOMIC
(3 kinds)
NONMETALS
METALS
P = metal atoms
B = delocalized
covalent (metallic)
S = Cu, Fe, Na
C = solid state
MP = high
ALLOYS
(Metallic solutions)
Substitutional alloy:
Some metal atoms are
replaced by other metal
atoms
Interstitial alloy: Smaller
metal atoms fit inside the
spaces between larger metal
ions
(Network Crystals
or
Macromolecules)
P = nonmetal atoms
B = covalent
S = diamond, quartz)
MP = extremely high
NOBLE GASES
P = group 8A atoms
(noble gas atoms)
B = London dispersion
S = Ne, Ar
MP = very low
20. Define vapor pressure and demonstrate how to calculate changes in it at different temperatures.
Vapor pressure is the upward force per unit area of liquid particle during evaporation. When
comparing different liquids, vapor pressure increases as the strength of the intermolecular forces
between molecules decreases. For example, acetone has weak intermolecular forces, evaporates readily
(highly volatile), and has a high vapor pressure.
The vapor pressure of a liquid increases as its temperature increases due to the increases kinetic energy
of the molecules. The Clausius-Clapeyron equation (at right) is used to
calculate changes in vapor pressure at different temperatures.
Liquids after sitting for a length of time in (a)
an open container and (b) a closed container.
The liquid in the open container has evaporated, but the liquid in the closed
container has brought about a rise in pressure. Evaporation of a liquid results in
more gas phase molecules which exert a pressure in a closed container
The distribution of molecular kinetic energies in a liquid at two temperatures. Only
the faster-moving molecules have sufficient kinetic energy to escape from the liquid
and enter the vapor. The higher the temperature, the larger the number of molecules
with enough energy to escape.
21. Define molar heat of fusion and vaporization and demonstrate how to
calculate the amount of energy required for phase changes with them.
molar heat of fusion is the amount of energy (q) required to
change one mole of a solid to a liquid at its melting point.
q = moles x Hfus
molar heat of vaporization is the amount of energy (q) required
to change one mole of a liquid to a gas at its boiling point.
q = moles x Hvap
A heating curve for H2O, shows the temperature changes and
phase transitions that occur when heat is added. The plateau at
0°C represents the energy required to melt solid ice (Hfus), and the
plateau at 100°C represents the energy required to boil liquid
water(Hvap). Plateau regions in a heating curve indicate a change
in phase of the substance where the temperature remains constant. This is because all of the energy of the
system is used to break the forces of attraction between particles. Enthalpy and entropy increase as matter
changes from solid to liquid to gas.
22. Define liquefaction, critical temperature, and critical pressure and describe how they relate to phase
diagrams.
When a liquid is heated in a closed container, the vapor pressure increases until no clear distinction
between liquid and gas is visible and the meniscus disappears. The highest temperature reached before the
meniscus disappears is called the critical temperature. The critical pressure is the vapor pressure of the liquid at
this temperature. Liquefaction is the process by which a gas is transformed into a liquid. This is achieved by
lowering the temperature, increasing the pressure, or both. Critical temperature and pressure can be defined
with respect to liquefaction. The critical temperature is the temperature above which no amount of pressure will
liquefy a gas. Critical pressure is the vapor pressure of a liquid at the critical temperature. When comparing
different liquids, those with high critical temperatures and critical pressures have strong intermolecular forces.
23. Define phase diagram and triple point and interpret information given on a phase diagram.
A phase diagram is a graph that shows the states of a
substance under varying conditions of temperature and
pressure. The triple point is the point at which all three
phases are in equilibrium with one another. Water is
the only substance whose solid-liquid equilibrium has a
negative slope. In accordance with LeChatelier’s
Principle, increasing the pressure favors the denser,
liquid state. For all other substances, increasing the
pressure favors the denser, solid state.