OME General Chemistry
Lecture 5: Intermolecular Forces, Solids, Liquids
Dr. Hartwig Pohl
Office: Beyer-Bau 122e
Email: [email protected]
Phone: +49 351 463 42576
1
Intermolecular Interactions
We have studied the intramolecular forces among atoms or ions
within a molecule…
What holds molecules together to form matter?
Intermolecular Forces (IMFs): the forces that hold one molecules
to another molecule
• Electrostatic attraction between oppositely charged portions of
molecules
• Cause changes in the state of a material (solid, liquid, gas)
• Changes in physical properties (melting point, boiling point,
vapor pressure, etc.) based on strength of the IMFs
Main Types of IMFs: Dipole-Dipole, Hydrogen Bonding, IonInduced Dipole, Dipole-Induced Dipole, Induced Dipole-Induced
Dipole (London Dispersion Forces)
2
LI QUI DS & SOLI DS
w it is time to consider the forces that condense matter. The forces that hold one molecule to
lecule are referred to as intermolecular forces (IMFs). These forces arise from unequal dis
he electrons in the molecule and the electrostatic attraction between oppositely charged po
lecules. We briefly visited the IMFs earlier when discussing the nonideal behavior of gase
ces cause changes of state by causing changes among the molecules NOT within them.
ysical properties such as melting points, boiling points, vapor pressures, etc. can be attributed
Intermolecular Interactions
Gas (g)
Liquid (l)
Solid (s)
ngth of the intermolecular
between
It works
like this:
Oxygen, O2 attractions present
Water,
H2O molecules.
Sodium
Chloride,
NaClthe lowe
ling point (or vapor
pressure
weaker the intermolecular
attractions; the
very low
boiling or melting point),
lowthe
boiling
very high boiling
boiling point, the
stronger
attractions.
gasoline
weakest
IMFsthe intermolecular
stronger
IMFs For example,
strongest
IMFsevaporates m
re quickly than water. Therefore, the intermolecular attractive forces that hold one gasoline m
another are much weaker than the forces of attraction that hold one water
(ionic molecule
bonding) to another
lecule. In fact, water molecules are held together by the strongest of the intermolecular attrac
ces, hydrogen bonds. Hydrogen bonds are not true bonds—they are just forces of attraction t
ween a hydrogen atom on one molecule and the unshared electron pair on fluorine, oxygen
o
3
ogen atoms of a neighboring molecule. The strands of DNA that make up our genetic code a
Dipole-Dipole Interactions
Dipole-Dipole: the force of
attraction that enables two
polar molecules to attract
one another.
Polar molecules are those
which have an uneven
charge distribution since
their dipole moments do not
cancel.
Strongest of the IMFs!
4
Hydrogen Bonding
Hydrogen bonding: the force of
attraction
the H-atom
of one
ding—the
force between
of attraction
between
the
molecule and an unshared electron pair
of one molecule
and an unshared electron
on F, O, or N of a neighboring molecule
N of a neighboring
molecule
(a special
(a special case of
dipole-dipole).
bonded H
H “bond”
dipole). This is the strongest IMF. Never
en bonding
a bondedIMF.
hydrogen. The
This iswith
the strongest
l properties of water are due to the fact
Note:
Don’t confuse
H-bonding
with a
hydrogen
bonding.
As a result
of these
hydrogen.
er has a bonded
high boiling
point, high specific
other unusual properties.
The unique physical properties of water are due to the fact that it exhibits
hydrogen bonding. As a result of these attractions, water has a high b.p., high
specific heat, and many other unusual properties.
5
ydrogen atom of one molecule and an unshared electron
air on F, O, or N of a neighboring molecule (a special
ase of dipole-dipole). This is the strongest IMF. Never
onfuse hydrogen bonding with a bonded hydrogen. The
nique physical properties of water are due to the fact
hat it exhibits hydrogen bonding. As a result of these
ttractions,
water
a high
boilinginpoint,
high specific
Why
so has
much
variation
b.p. among
the covalent hydrides of groups IV through VII?
eat, and many other unusual properties.
Hydrogen Bonding
Expectation: increase b.p. with
molecular mass  more e– in a
molecule, more polarizable e–
cloud, stronger IMFs  more
energy needed to overcome
those increased attractions
WHY is there such variation in the boiling point among the covalent hydrides of groups
lighter
hydrides
have
higher
EN values
which
to more
polar H−X bonds.
V through1.VII?
One would
expect
that
BP would
increase
withleads
increasing
molecular
small
size of
dipolethe
allows
a closer approach
of theabout
dipoles,
mass [since2.theThe
more
electrons
in each
a molecule,
morefor
polarizable
the cloud {more
hat in the nextfurther
section},strengthening
the stronger the
IMFs,
thereforeRemember,
the more E needed
to overcome
the
attractions.
attractive
forces dissipate
hose increasedwith
attractions
and vaporize,
increased
distance.thus the higher the boiling point. That’s how it
s supposed to work!]. Hydrogen bonding, that’s why!
TWO reasons: both enhance the IMF we refer to as hydrogen bonding.
6
Ion/Dipole–Induced Dipole
Ion-Induced Dipole: the force of attraction between a charged ion and a
nonpolar molecule. The ion greatly perturbs the electron cloud of the
nonpolar molecule and polarizes it transforming it into a temporary dipole
which enhances the ion’s attraction for it.
–
–
–
Dipole-Induced Dipole: the force of attraction between a polar molecule and
a nonpolar molecule. The polar molecule induces a temporary dipole in the
nonpolar molecule. Larger molecules are more polarizable than smaller
molecules since they contain more electrons. Larger molecules are more
likely to form induced dipoles.
7
London Dispersion Forces
Induced Dipole-‐Induced Dipole (London Dispersion Forces): the force of
attraction between two non polar molecules due to the fact that they can
form temporary dipoles. Nonpolar molecules have no natural attraction
for each other. Without these forces, we could not liquefy covalent gases
or solidify covalent liquids.
These forces are a function of the number of electrons in a given
molecule and how tightly those electrons are held.
8
London Dispersion Forces
Understanding London Dispersion Forces…
Molecules are constantly colliding with each other. When these collisions
occur, the electron cloud around the molecule is distorted. This produces a
momentary induced dipole within the molecule! The amount of distortion
of the electron cloud is referred to as polarizability.
All molecules have electrons  all molecules have these forces.
Strength ranges from 5—40 kJ/mol and increases as the number of
electrons increase due to increasing polarizability.
Ex: Halogens exist as nonpolar, diatomic molecules at room temperature
and atmospheric conditions. F2 and Cl2 are gases, Br2 is a liquid and I2 is a
solid. Increasing strength of IMFs due to increased polarizibility!
9
Intermolecular Forces Overview
10
Example: IMFs
What are the most important IMFs between the following
molecules and atoms:
a) NaCl (aq)
b) Fe2+ and O2
c) CH3Cl an CCl4
11
id lauqenu morf esira secrof esehT .)sFM I( sec rof ra luce lom retn i sa ot derrefer era eluce
p degrahc yletisoppo neewteb noitcartta citatsortcele eht dna elucelom eht ni snortcele eh
sag fo roivaheb laedinon eht gnissucsid nehw reilrae sFMI eht detisiv ylfeirb eW .seluce
.meht nihtiw TON selucelom eht gnoma segnahc gnisuac yb etats fo segnahc esuac se
etubirtta eb nac .cte ,serusserp ropav ,stniop gniliob ,stniop gnitlem sa hcus seitreporp lacis
Physical States & IMFs
retni eht fo htgn
celomIMFs
aluFlow
cartta rto
eb tneserp snoitAbility
I .selucelom neewtCompressibility
l skrow& tVolume
siht ekiShape
ol eht :State
rekaew eht ,)tniop gnitlem ro erusserp ropav ro( tniop gni
i ehtshape
etnown
alucelomrits
ht ;snoitcartta rMaintains
Solid
ht ,tniop gniliob
gnorts eStrong
i eht renone
celomretnAlmost
tta ralunone
,elpmaxe roF .snoitcarAlmost
sagvolume
setaropave eniloand
f evitcofartta ralucelomretni eht ,eroferehT .retaw naht ylkciuq e
h taht sectoroshape
enilosag eno dloConforms
hcum era rehton
ht naht rekaewModerate
fo secrof eModerate
carttalow
h taht noitVery
eno dlolimited
om retaw volume
t elucelcontainer,
ehtona oLiqui
i eht fo tsegnorts eht yb rehtegot dleh era selucelom retaw ,tcaf nI .eluce
rtta raludcelomrebytnsurface
noitcartta fo secrof tsuj era yeht—sdnob eurt ton era sdnob negordyH .sdnob negordyh ,se
Conforms to shape and
negordyh a nee
m eno no mota Weak
eht dna eluceloHigh
nortcele derahsnuHigh
apcontainer
lf no riof
,enirouvolume
negyxoGas
edoc citeneg ruo pu ekam taht AND fo sdnarts ehT .elucelom gnirobhgien a fo smota neg
.noitcartta ralucelomretni fo epyt siht yb reht
12
Closer Look: Solids
The particles that make up a solid material, whether ionic, molecular,
covalent, or metallic, are held in place by strong attractive forces.
In discussing solids, the positions of the atoms, molecules, or ions, which
are fixed in place, are considered.
The constituents of a solid can be arranged in two general ways:
1. They can form a regular repeating three-dimensional structure called a
crystal lattice, thus producing a crystalline solid.
2. They can aggregate with no particular order, in which they form an
amorphous solid.
13
Crystalline Solids (Crystals)
General Properties of Crystalline Solids:
• Distinctive internal structures that lead to distinctive flat surfaces, or faces
• Faces intersect at angles that are characteristic of the substance and reflect the
regular repeating arrangement of the component atoms, molecules, or ions
• Structure produces a distinctive X-ray pattern that can be used for
identification
• Relatively sharp, well-defined melting points because all component atoms,
molecules, or ions are the same distance from the same number and type of
neighbors: regularity of the crystalline lattice creates local environments that
are the same
– Intermolecular forces holding the solid together are uniform, and the
same amount of thermal energy is needed to break all of the interactions
simultaneously
14
Amorphous Solids
General Properties of Amorphous Solids:
• When cleaved or broken, they produce fragments with irregular surfaces.
• They have poorly defined X-ray patterns
• Almost any substance can solidify in amorphous form if the liquid phase is
cooled rapidly enough, but some solids are intrinsically amorphous
• The local environment, including both the distances to the neighboring units
and the numbers of neighbors, varies throughout the material.
– Different amounts of thermal energy are needed to overcome these
different interactions.
– Tend to soften slowly over a wide temperature range rather than having
a well-defined melting point.
– If maintained at a temperature just below the m.p. for long periods of
time, the component molecules, atoms, or ions can gradually rearrange
into a more highly ordered crystalline form.
15
X-Ray Analysis of Crystalline Solids
X-ray diffraction is a useful tool for obtaining information about the structures
of crystalline substances because the wavelength of X-ray radiation is
comparable to the interatomic distances in most solids.
•
A beam of X-rays is aimed at a sample of a crystalline material, and the Xrays are diffracted by layers of atoms in the crystalline lattice.
•
When the beam strikes photographic film, it produces pattern, which
consists of dark spots on a light background.
16
X-Ray Analysis of Crystalline Solids
Bragg worked out the mathematics that allow X-ray diffraction to be used to
measure interatomic distances in crystals and to provide information about
the structures of crystalline solids…
X-rays
diffracted
fromshows
different
ofwaves
atomsbeing
in a solid
reinforce
one in two different la
The diagram
below
two planes
in-phase
reflected
by atoms
another if they are in phase, which occurs only if the extra distance they travel
a crystal. The
extra
distance
traveled
y the lower wave is the sum of the distances xy and y
corresponds
to an
integral
number
of wavelengths.
the waves will be in phase after reflection if xy + yz = nλ.
Bragg equation: 2d sin = n
17
Arrangement of Atoms in Crystals
A crystalline solid consists of repeating patterns of its components in three
dimensions (a crystal lattice).
An entire crystal can be represented by drawing the structure of the
smallest identical units (called a unit cell) that, when stacked together, form
the crystal.
18
The Cubic Unit Cell
The three types of cubic unit cells:
Simple cubic: unit cell consists of eight component atoms, molecules, or
ions located at the corners of the cube
Body-centered cubic (bcc):
unit cell also contains an
identical component in the
center of the cube
Face-centered cubic (fcc):
unit cell has components in
the center of each face in
addition to those at the
corners of the cube
19
Density Based on Unit Cell
A solid consists of a large number of units cells arrayed in three
dimensions.
Density is the mass of substance per unit volume, and the density of the
bulk material can be calculated from the density of a single unit cell.
For the calculation, it is necessary to know the size of the unit cell (to
obtain its volume), the molar mass of its components, and the number
of components per unit cell.
When counting atoms or ions in a unit cell, those lying on a face, an
edge, or a corner contribute to more than one unit cell.
– Face  two adjacent unit cells  1/2 atom per unit cell
– Edge  four adjacent unit cells  1/4 atom per unit cell
– Corner  eight adjacent unit cells  1/8 atom per unit cell
– Atoms that lie entirely within a unit cell belong to only that unit cell.
20
Simple Binary Compounds
The structures of most binary compounds are dictated by the packing
arrangement of the largest species (the anions), with the smaller
species (the cations) occupying appropriately sized holes in the anion
lattice.
A simple cubic lattice of anions contains
only one kind of hole, located in the
center of the unit cell.
• Equidistant from all eight atoms at the
corners of the unit cell  cubic hole
• Many ionic compounds that contain large
cations and have a 1:1 cation:anion ratio
have this structure.
21
Simple Binary Compounds
The structures of most binary compounds are dictated by the packing
arrangement of the largest species (the anions), with the smaller
species (the cations) occupying appropriately sized holes in the anion
lattice.
A face-centered cubic array of atoms or
anions contains both octahedral holes and
tetrahedral holes
• Octahedral holes: one in the center plus a
shared one in the middle of each edge
• Tetrahedral holes: located between an atom
at a corner and the three atoms at the
centers of the adjacent faces
• Cations of intermediate size occupy
octahedral holes in an fcc anion lattice, and
small cations occupy tetrahedral holes
22
Ex: Simple Binary Compounds
Simple Cubic, CsCl
Face Centered Cubic, NaCl
octahedral holes
Face Centered Cubic, ZnS
tetrahedral holes
23
Defects in Crystals
Real crystals contain large numbers of defects ranging from variable
amounts of impurities to missing or misplaced atoms or ions
Defects occur for three main reasons:
1. It is impossible to obtain any substance in 100% pure form; some
impurities are always present.
2. Forming a crystal requires cooling the liquid phase to allow all
atoms, ions, or molecules to find their proper positions, but cooling
results in one or more components being trapped in the wrong place
in a lattice or in areas where two lattices that grow separately
intersect.
1. Applying an external stress to a crystal can cause microscopic regions
of the lattice to move with respect with the rest; this results in
imperfect alignment.
24
Classification of Solids
Based on the nature of the forces that hold the component atoms, molecules, or
ions together, solids are classified as…
Atomic Solids
Lattice Points
IMFs
Metallic
Network
Group 8A
Molecular
Solids
Metal
Atoms
Nonmetal
Atoms
Group 8A
Atoms
Discrete
Molecules
Ions
London
Dispersion
Forces
Dipoledipole
and/or
London
Dispersion
Forces
Ionic
Delocalized Directional
Covalent
Covalent
Ionic Solids
 Variation in the relative strengths of these four types of interactions correlates
with their wide variation in properties of these four kinds of solids.
25
Metallic Solids
Bonding Models for Metals
Bonding in metallic solids is quite different from the bonding in the other kinds
of solids… the
valence electrons
are delocalized
throughout
crystal,and
providing
Remember,
metals conduct
heat and electricity,
arethe
malleable
ductile, and have
a strong cohesive
that holds
atoms
together
These force
facts indicate
thatthe
the metal
bonding
in most
metals(delocalized
is both strong and nondirection
covalent bonding)
separate atoms, but easy to move them provided they stay in contact with each othe
Electron Sea Mode
metals in a “sea” o
A metals pictured a
Band (Molecular
Electrons assumed
metal crystal in
valence atomic orbi
• Strength of metallic bonds varies dramatically
• Metallic bonds tend to be weakest for elements that have nearly empty or
nearly full valence shells, and are strongest for elements with half-filled
valence shells
Intermolecular Forces, Liquids & Solids
26
Metallic Solids (cont.)
General Properties and Notes:
• Packing efficiency in metallic crystals tends to be high, so metallic solids
are dense, with each atom having as many as 12 nearest neighbors
• Every lattice point in a pure metallic element is occupied by an atom of the
same metal
• Reflect light, called luster
• Have high electrical and thermal conductivity
• Have high heat capacity
• Are malleable and ductile
• High density
• Melting points depend strongly on electron configuration
27
Metallic Solids (cont.)
An alloy is a mixture of metals with metallic properties that differ from those
of its constituent elements.
Most metallic substances are alloys, and the composition of most alloys
varies over wide ranges.
Substitutional alloys are metallic solids
that contain large numbers of
substitutional impurities and can be
formed by substituting one metal atom
for another of similar size in the lattice.
Interstitial alloys are metallic solids that
contain small numbers of interstitial
impurities and can be formed by inserting
smaller atoms into holes in the metal
lattice.
28
Network (Covalent) Atomic Solids
λ
Formed by networks or chains of atoms held together by covalent bonds,
tend to be very hard and have high melting points, not easily deformed,
brittle, poor conductors of heat and electricity, low density, dull surface
co
Pl
in
ho
at
di
iro
Network Atomic
Network
Examples: Diamond vs. Graphite
• Diamond: hard, colorless, insulator, bonds
with neighbors in a tetrahedral 3D
geometry (sp3 hybridization, 109.5º bond
angles)
• Graphite: slippery, black, conductor, layers
of C atoms fused in six-membered rings
(graphene) with weak bonding in the 3rd
dimension (sp2 hybridization, 120º bond
angles, unhybridized p-orbital forms πbonds and accounts for electrical
conductivity)
Composed
Atomic
S
molecule”
Composed
diamondofb
molecule”. B
weak bond
diamond bond
weak bonding
o
o
o
o
o
29
Molecular Solids
Consist of molecules held together by dipole-dipole interactions, London
dispersion forces, or hydrogen bonds
• Intermolecular interactions in a molecular solid are relatively weak
compared with ionic and covalent bonds
• Tend to be soft, low melting, and easily vaporized
• For similar substances, the strength of the London dispersion forces
increases smoothly with increasing molecular mass
• Usually poor conductors of heat and electricity
• Low density
• Dull surface
30
Ionic Solids
Consist of positively and negatively charged ions held together by
electrostatic forces
• Strength of the attractive forces depends on the charge and size of the
ions that make up the lattice and determines many of the physical
properties of the crystal
• Lattice energy, is directly proportional to the product of the ionic
charges and inversely proportional to the sum of the sizes of the ions
• Poor conductors of heat and electricity
• High melting points
• Hard but brittle; shatter under stress
• Dense with a dull surface
31
Metallic Bonding: Band Theory
Band Theory: An approach to metallic bonding based on MO Theory,
assumes that the valence orbitals of the atoms in a solid interact, generating
a set of molecular orbitals that extend throughout the solid
• Imagine a linear arrangement of n metal atoms, with an ns1 orbital (Li)
• Orbital overlap  LCAO  bonding and antibonding molecular orbitals
• Energy separation between adjacent orbitals decreases as the number of
interacting orbitals increases  continuum of energy levels
corresponding to a particular molecular orbital extending throughout the
metal atoms (energy band)
32
Band Theory (cont.)
For a one-dimensional system, band contains n energy levels corresponding
to the LCAO of s orbitals from n metal atoms
• Valence electrons occupy the lowest energy levels, only the lower half of
the band is filled (the bonding molecular orbitals)  valence band
• Unfilled band (antibonding molecular orbitals)  conduction band
In two- or three-dimensional systems with atoms that contain electrons in p
and d orbitals, the resulting energy-level diagrams are essentially the same.
Each band will have a different bandwidth and will be centered at a
different energy, corresponding to the energy of the parent atomic orbital
of an isolated atom.
Energy difference between the valence band and conduction band is called
the band gap  conductor, semiconductor, insulator
33
Examples: Solids
Classify each of the following substances according to the type of
solid it forms:
a) Gold, Au
b) Carbon Dioxide, CO2
c) Lithium Fluoride, LiF
d) Argon, Ar
Atomic Solids
Lattice Points
IMFs
Metallic
Network
Group 8A
Molecular
Solids
Metal
Atoms
Nonmetal
Atoms
Group 8A
Atoms
Discrete
Molecules
Ions
London
Dispersion
Forces
Dipoledipole
and/or
London
Dispersion
Forces
Ionic
Delocalized Directional
Covalent
Covalent
Ionic Solids
34
Examples: Solids (answers)
Classify each of the following substances according to the type of
solid it forms:
a) Gold, Au
b) Carbon Dioxide, CO2
c) Lithium Fluoride, LiF
d) Argon, Ar
atomic solid – metallic
molecular solid
binary ionic solid
atomic solid – group 8, very weak IMFs
Atomic Solids
Lattice Points
IMFs
Metallic
Network
Group 8A
Molecular
Solids
Metal
Atoms
Nonmetal
Atoms
Group 8A
Atoms
Discrete
Molecules
Ions
London
Dispersion
Forces
Dipoledipole
and/or
London
Dispersion
Forces
Ionic
Delocalized Directional
Covalent
Covalent
Ionic Solids
35
id lauqenu morf esira secrof esehT .)sFM I( sec rof ra luce lom retn i sa ot derrefer era eluce
p degrahc yletisoppo neewteb noitcartta citatsortcele eht dna elucelom eht ni snortcele eh
sag fo roivaheb laedinon eht gnissucsid nehw reilrae sFMI eht detisiv ylfeirb eW .seluce
.meht nihtiw TON selucelom eht gnoma segnahc gnisuac yb etats fo segnahc esuac se
etubirtta eb nac .cte ,serusserp ropav ,stniop gniliob ,stniop gnitlem sa hcus seitreporp lacis
Physical States & IMFs
retni eht fo htgn
celomIMFs
aluFlow
artta rto
eb tneserp snoitcAbility
I .selucelom neewtCompressibility
skrow& tVolume
siht ekilShape
ol eht :State
rekaew eht ,)tniop gnitlem ro erusserp ropav ro( tniop gni
i ehtshape
etnown
alucelomrits
ht ;snoitcartta rMaintains
Solid
ht ,tniop gniliob
gnorts eStrong
i eht renone
celomretnAlmost
tta ralunone
,elpmaxe roF .snoitcarAlmost
sagvolume
setaropave eniloand
artta ralucelomretni eht ,eroferehT .retaw naht ylkciuq e
f evitcof
h taht sectoroshape
enilosag eno dloConforms
hcum era rehton
ht naht rekaew Moderate
fo secrof eModerate
arttalow
taht noitcVery
no dlohlimited
om retaw evolume
t elucelcontainer,
ehtona oLiquid
i eht fo tsegnorts eht yb rehtegot dleh era selucelom retaw ,tcaf nI .eluce
tnsurface
rtta ralucelomreby
noitcartta fo secrof tsuj era yeht—sdnob eurt ton era sdnob negordyH .sdnob negordyh ,se
Conforms to shape and
High
negordyh a nee
eno no mota Weak
eht dna elucelom
nortcele derahsnuHigh
lf no riaofp container
,enirouvolume
negyxoGas
edoc citeneg ruo pu ekam taht AND fo sdnarts ehT .elucelom gnirobhgien a fo smota neg
.noitcartta ralucelomretni fo epyt siht yb reht
36
y rtsimehC *PA
SECROFCloser
RA LUC
E
L
O
M
R
E
T
N
I
Look: Liquids
SD I LOS & SD IUQ I L
Notes About Liquids:
ona ot elucelom eno dloh taht secrof ehT .rettam esnednoc taht secrof e
ubirtsid la•uqMolecules
enu moinrfliquids
esiraare
severy
crofclose
esetogether,
hT .)swith
FMessentially
I( sec rofnora luce lom retn
empty space between them.
noitrop degrahc yletisoppo neewteb noitcartta citatsortcele eht dna elu
T .sesag• foMolecules
roivaheinbliquids
laediare
noinn constant
eht gnmotion,
issucsiand
d ntheir
ehwkinetic
reilrae sFMI eht d
energy
their
temperature.
.m
eht n(and
ihtihence
w TO
N sspeed)
elucedepends
lom ehont gtheir
nom
a segnahc gnisuac yb
ht ot detub• irProperties
tta eb naofc liquids
.cte ,scan
erubessexplained
erp ropusing
av ,satnmodified
iop gnversion
iliob ,of
stniop gnitle
the kinetic molecular theory.
37
Characteristic Properties of Liquids
Density
molecules are packed relatively close together,
similar packing to the same material in the solid
state, measured in g cm–3 = g mL–1
Mol. Order
high KE of molecules, but arrangement of
molecules is not completely random  shortrange order due to strong IMFs
Compressibility
cannot be readily compressed because there is
very little empty space between the molecules
Therm. Exp.
IMFs are strong enough to prevent significant
expansion upon heating
Fluidity
liquids adjust to the shape of their container
Diffusion
molecules diffuse, they are in constant motion
38
Unique Properties of Liquids
There are three unique properties of liquids that depends intimately on
the nature of intermolecular interactions…
Surface Tension: the energy required to increase the surface area of a
liquid by a specific amount, measured in energy per unit area (J m–2)
Stronger IMFs  higher the surface tension
Surfactants are molecules such as soap and detergents that reduce the
surface tension of polar liquids like water by disrupting the
intermolecular attractions between adjacent molecules
39
Unique Properties of Liquids
Capillary Action: tendency of a polar liquid to rise against gravity into a
small-diameter glass tube (a capillary tube)
Net result of two opposing sets of forces: cohesive forces, which are the
intermolecular forces that hold the liquid together, and adhesive forces,
which are the attractive forces between the liquid and the substance that
makes up the capillary (Ex: glass).
• adhesive > cohesive forces, the liquid in the capillary rises to the level
where the downward force of gravity balances the upward force.
• adhesive < cohesive forces, the liquid is pulled down into the capillary
below the surface of the bulk liquid.
• meniscus, shape depends on the relative strengths of the cohesive
and adhesive forces.
40
Unique Properties of Liquids
Viscosity: resistance of a liquid to flow, liquids that have strong
intermolecular forces have high viscosities
Evaluation of viscosity  (1) measure the time it takes a quantity of
liquid to flow through a narrow vertical tube, (2) measure the time it
takes steel balls to fall through a given volume of the liquid.
Viscosity and molecular shape: liquids consisting of long, flexible
molecules tend to have higher viscosities than those composed of more
spherical or shorter-chain molecules.
Viscosity and temperature: decreases rapidly with increasing T because
the kinetic energy of the molecules increases, higher kinetic energy
enables the molecules to overcome the attractive forces that prevent the
liquid from flowing.
41
Vapor Pressure of Liquids
When the liquid is heated, its molecules obtain sufficient kinetic
energy to overcome the IMFs holding them in the liquid phase and
they escape into the gaseous phase.
The result of this phenomenon is that the molecules from the liquid
phase generate a population of molecules in the vapor phase above
the liquid that produces a pressure called the vapor pressure of the
liquid.
only molecules with KE > E0
can escape to vapor phase
increasing T  increased KE
# of molecules with KE > E0
increases with increasing T
Note: only mol. at surface
42
Vapor Pressure of Liquids
Only molecules at the surface can undergo evaporation, or vaporization,
where atoms or molecules gain sufficient energy to enter a gaseous state,
thereby creating a vapor pressure.
A fraction of the molecules in the vapor phase will collide with the surface of
the liquid and reenter the liquid phase (condensation)
# mol. in the vapor phase
increases  # collisions increases
Steady state will be reached
where # mol. vaporized = # mol.
condensing
43
Equilibrium Vapor Phase
Volatile liquids have high vapor pressures (higher than water) and tend to
evaporate readily; nonvolatile liquids have low vapor pressures (lower
than water) and evaporate more slowly.
Two opposing processes (such as evaporation and condensation) that
occur at the same rate, thus producing no net change in the system,
constitute a dynamic equilibrium  pressure exerted by a vapor in
dynamic equilibrium with a liquid is the equilibrium vapor pressure
Equilibrium vapor pressure of a substance at a particular temperature is a
characteristic of the material  Does not depend on the amount of liquid
 Depends strongly on the temperature and IMFs
44
Equilibrium Vapor Phase
Temperature dependence is strong because the vapor pressure depends on
the fraction of molecules that have a kinetic energy higher than that needed
to escape from the liquid
Can use logarithms to express the non-linear relationship between vapor
pressure and temperature as a linear relationship  Clausius–Clapeyron
Equation
45
Equilibrium Vapor Phase
Clausius–Clapeyron Equation can be used to calculate the enthalpy of
vaporization of a liquid from its measured vapor pressure at two or
more temperatures
–DH vap æ 1 ö
ln P =
ç ÷+C
R èT ø
linear relationship
P = vapor pressure
ΔHvap = enthalpy of vaporization
R = gas constant (8.314 J mol–1 K–1)
T = temperature in K
C = y-intercept, constant
Calculating the enthalpy of vaporization:
æ P2 ö –DH vap æ 1 1 ö
ln ç ÷ =
ç - ÷
R è T2 T1 ø
è P1 ø
46
Vapor Pressure  Boiling Point
As the temperature of a liquid increases, the vapor pressure of the
liquid increases until it equals the external pressure, or the
atmospheric pressure.
 Bubbles of vapor begin to form throughout the liquid, and the
liquid begins to boil
 Temperature at which a liquid boils at exactly 1 atm pressure is
the normal boiling point of the liquid.
47