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Chapter 10: Liquids and Solids
STATES OF MATTER: At any temperature above absolute zero, the atoms, molecules,
or ions that make up a substance are moving. In the kinetic theory of gases, this
motion is the only thing that has to be known about a gas in order to describe its
behavior. A gas molecule moves until it hits something, then it bounces off and moves
in a different direction. The molecules of a gas act as if there were absolutely no
attractive forces acting between them.
In the liquid and solid states, the particles of matter act as if they were stuck together.
In liquids, this attractive force is strong enough to keep the particles close together, but
still free to flow. In solids, this force is strong enough to lock the particles together in a
rigid crystal lattice, so that the particles are free only to vibrate.
The transition from solid to liquid to gas is accomplished by giving the particles of
matter enough kinetic energy (heat) to overcome whatever attractive forces hold them
together.
10.1 Intermolecular Forces
Actually, there are several different types of attractive forces, all of which can be
ultimately described in terms of electrostatic attractions. For ionic compounds, the
attractive force is the ionic bond, a strong electrostatic attractive force between
oppositely charged ions. This force is so strong that all ionic compounds are solid at
room temperature, with melting points usually in the hundreds of degrees. In
molecular substances there are no ions, so subtler interactions must be involved. These
are of three main types: dipole forces, dispersion forces, and hydrogen bonds.
DIPOLE FORCE - electrostatic attractive force between oppositely charged ends of
polar molecules. This force is not as strong as it would seem. The magnitude of the
dipole force depends on the dipole moment of the molecule. Since there is an upper
limit to bond polarity, there is also a limit to how strong the dipole force can be. Even
very polar molecules may be gases at room temperature in the absence of any other
forces.
∂+ ∂- ∂+ ∂- Example: CO
HYDROGEN BOND - an extremely strong special
case of the dipole force which occurs when H is
bonded to F, O, or N. Since H has no inner shell
electrons, a sufficiently polar bond will leave the H
nucleus almost stripped of its electron density,
which allows the partially positive H atom to
approach the negative end of a molecule much
closer than it normally could.
Since the magnitude of an electrostatic attraction
depends inversely on the square of the charge
separation, the close approach of the H atom
results in a much stronger attractive force than the
normal dipole force. As a result, molecules which
exhibit hydrogen bonding have melting and boiling
points that are quite higher than similar molecules
which cannot hydrogen bond. Compare ethyl
alcohol, CH 3
CH
OH, b.p.= 78°C , with dimethyl
ether. CH 3
OCH
, b.p.=-25°C. Both molecules are
the same size and have comparable dipole forces,
but because ethyl alcohol can hydrogen bond, it has
a boiling point 103° higher than dimethyl ether.
Figure 11.6, pg. 449 in the text shows how
hydrogen bonding affects the boiling points of the
hydrides of the group IV, V, VI, and VII elements.
H F
1.9 difference in electronegativity
H F H F
H
2
O
CH
4
SnH 4
GeH 4
SiH 4
H
2
S
H
2
Se
H
2
Te
DISPERSION FORCE - electrostatic attractive force between oppositely charged ends
of temporary dipoles.
I I I I
A temporary, or induced, dipole is an unstable state
which occurs when the electron density of a molecule
is distorted away from its equilibrium position,
causing a temporary overabundance of electron
density at one end of the molecule (resulting in a
temporary partially negative charge), and leaving a
deficit of electron density at the other end of the
molecule (resulting in a temporary partially positive
charge).
Since the electrons are not rigidly fastened to the nucleus, such a distortion could occur
simply as a result of thermal motion. The degree to which a molecule may become
polarized by this mechanism depends on how tightly the electrons are held, which
depends on how large the molecule is. Since there is no upper limit to how large a
molecule can be, and since the dispersion force can act over the entire surface of the
Cohesion: intermolecular attraction between molecules.
Adhesion: attraction between unlike molecules.
If adhesion > cohesion, liquid will rise until countered by gravity.
If cohesion > adhesion, liquid will be depressed until countered by gravity.
water mercury
The viscosity, or resistance to flow, of a liquid, depends both on the intermolecular
forces and on the molecular structure of the liquid. Stronger forces result in higher
viscosity, as does a structure which allows for tangling of molecules. See table 11.4 on
pg. 452.
In most solids, molecules are arranged so as to maximize dispersion forces, which
results in most surface contact and densest packing
WATER
O
O: :O
H H
H H
H-bonding in ice causes an open 3-
dimensional tetrahedral structure
with a density less than that of liquid
water. From 0°C to 4°C the density of
liquid water actually increases with
temperature as more H-bonds break
and structure becomes more compact.
Above 4° (or below freezing point)
water behaves normally.
temperature
normal substance
water
10.3 Structures and Types of Solids
In the solid state the atoms, ions, or molecules of the substance lack sufficient kinetic
energy to do anything but vibrate in place in a rigid three dimensional crystalline
lattice. Crystals are classified according to the unit occupying each lattice point and by
the forces holding the lattice units together. Substances which appear to be solid but
lack a crystalline structure are called amorphous solids.
Ionic solids - crystalline solids in which positive and negative ions occupy lattice points,
held together by ionic bonds.
Molecular solids - crystalline solids in which molecules occupy lattice points, held
together by intermolecular forces (van der Waals forces and hydrogen bonds). In
general (but not for ice) molecules are packed as close as their size and shape will
allow, to maximize dispersion forces. They are usually soft, with low melting points,
and are nonconductors.
Atomic solids - crystalline solids in which individual atoms occupy lattice points. The
atoms can be metals, C, B, Si, or others. The entire crystal is a single molecule.
Metallic crystals - crystalline solids in which individual metal atoms occupy lattice
points, held together by a delocalized network of molecular orbitals extending over the
entire crystal.
10. Structure and Bonding in Metals
Metals have few valence electrons and low electronegativity, which favors forming
decentralized bonds over localized bonds. This mode of bonding is often described as a
sea of valence electrons in which metal nucleii (and inner shell electrons) float. There
is no rigid network, which accounts for malleability and ductility. Metals are hard and
usually have high melting points and are conductors.
The MO treatment of metal bonding results in the band model. The many valence
orbitals in a metal crystal form a near-continuum of energy levels, in the form of MOs
which extend over the entire crystal. Since there is no energy gap between filled levels
and empty levels, electrons are easily promoted and can move.
Metal alloys: substitutional alloys have metal atoms of similar size replacing atoms in
the lattice. Interstitial alloys have small atoms in the interstices. Steel is an
interstitial alloy of carbon and iron. The strong directional Fe-C bonds add strength to
the alloy.
10.8 Vapor Pressure and Changes of State
Although the average
kinetic energy of a
molecule depends on
temperature, the actual
kinetic energy of any
individual molecule may
be much lower or much
higher. The dependence
of energy distribution
with temperature is
given by a curve such as
shown on the next page,
with many slow
molecules at low
temperatures, but more
fast molecules as the
temperature increases.
The significant point to
note is that even at low
temperatures some
molecules have high
kinetic energy.
Kinetic Energy
low T
high T
med T
Evaporation occurs when molecules with sufficient
energy break free from the liquid phase and enter the
gas phase. If the liquid is in good thermal contact with
its surroundings the temperature of the liquid will
remain constant so the rate of evaporation will be
constant. The pressure exerted by the vapor molecules
as they collide with other molecules in the gas phase
and with the walls of the container is called the vapor
pressure. As molecules of vapor accumulate (as the
vapor pressure increases) the rate at which vapor
molecules collide into the surface of the liquid
(condensation) increases. At some point the rates of
evaporation and condensation will be equal and the
process apparently stops: this condition is called
dynamic equilibrium. Both evaporation and
condensation are still occuring but with no net change
in vapor pressure. The equilibrium vapor pressure of a
pure liquid is a function only of temperature.
molecules moving fast
enough can overcome
attractive forces and enter
gas phase. Gas molecules
moving slow enough can be
recaptured by liquid.
liquid → vapor rate depends on T
vapor → liquid rate depends on P
equilibrium is reached when vapor pressure reaches certain value, depending on T:
Equilibrium vapor pressure. Increase volume, more vapor will form to maintain P,
until all liquid vaporizes. Large volume → total evaporation.
At low temperature, evaporation is limited to the surface of the liquid. Evaporation is
speeded up (with no effect on equilibrium vapor pressure) by increasing the surface
area of the liquid. Evaporation cannot take place within the volume of the liquid until
the vapor pressure is great enough to hold back the liquid around it. Once this occurs,
bubbles of vapor can form throughout the volume of the liquid, and the liquid is in a
state of boiling.
boiling point - temperature at which the vapor pressure of the liquid is equal to the
external pressure
normal boiling point - temperature at which the vapor pressure of the liquid is equal to
1 atmosphere (760 mm Hg) of pressure.
as external P increases, so does T b (pressure cooker, high alt. cooking)
temperature
liquid
gas
critical point
Since the boiling point of a liquid is
pressure dependant, a gas can be
liquified by either a decrease in
temperature or an increase in pressure:
however, there is a temperature beyond
which the gas cannot be liquified
regardless of pressure. This
temperature is called the critical
temperature , and the vapor pressure
at the critical temperature is called the
critical pressure. Gases with critical
temperatures below room temperature
can only be liquified cryogenically.
Heat of Vaporization ∆Hvap: energy required to convert 1 mole of liquid at 1 atm to
vapor
liquid(1 mol) → vapor(1 mol) ∆H = ∆Hvap
10.9 Phase Diagrams
A phase diagram is a plot which shows the phase of a substance as a function of its
pressure and temperature. There are sections for the three regions solid, liquid, and
gas, separated by lines indicating the equilibria between these phases. A point of
interest is the triple point, where solid, liquid, and gas all coexist at equilibrium.
temperature
solid
liquid
gas triple point
critical point
melting/freezing point ∆H fus
sublimation and deposition ∆H sub
∆H
sub
= ∆H
fus
+ ∆H
vap
WATER
Because of hydrogen bonding, the behavior of water is very unusual compared to other
liquids. First, water has a boiling point about 200° higher than expected by comparison
with H 2
S, H
Se, or H 2
Te. Without hydrogen bonding, water would be a gas; there
would be no liquid water anywhere on the planet. Another unusual property of water is
that ice floats. For most substances, the solid phase is denser than the liquid phase,
but water actually expands as it freezes. This effect is also due to hydrogen bonding.
The solid state of any substance is a crystalline lattice in which the particles are
arranged so as to maximize the attractive forces between them. For most substances,
this results in an efficient packing, especially if the dispersion force is strongest. For
water, hydrogen bonding is the strongest force, and this force is maximized when the H
atoms on one water are pointed at the lone pairs on neighboring water molecules (see
fig. 10.12, pg. 269 in the text). This structure leaves a lot of empty space, which is
partially occupied when the ice melts. As a result, ice is less dense than water and will
float in water.
The phase diagram of water reflects this behavior. For most substances, the line
separating the solid and liquid phases slopes to the right, so that an increase in
pressure favors the solid phase (the denser phase). For water, the solid-liquid line
slopes to the left (shown much exaggerated in the following phase diagram).
temperature
solid
liquid
gas triple point
critical point
triple point:
4.56 mm Hg
(0.0060 atm)
0.01° C
critical point:
219.5 atm
374.4° C
