States of matter
Chapter 11
Intermolecular Forces –
Liquids and Solids
• By changing the T and P, any matter
can exist as solid, liquid or gas.
• Forces of attraction determine physical
state
• Phase – homogeneous part of system
in contact with other parts of system,
separated by well-defined boundary
• e.g., ice in water, subliming dry ice,
evaporating isopropanol
Kinetic Molecular Theory of
Liquids and Solids
• Liquids and solids = “condensed states”
• Liquids
– Molecules are close together w/ little empty
space – difficult to compress
– Molecules held together by attractive
forces
– Liquid has definite volume
– Molecules can move past each other
freely, flow into shape of container
Kinetic Molecular Theory of
Solids
• Solids
– Molecules are held rigidly in place
• vibrate about a fixed point
– Less compressible than liquids.
– Usually solid is denser than liquid
• important exception = water
– Definite shape and volume
11.2 Intermolecular Forces
• Attractive forces between molecules.
• As T of gas drops, intermolecular forces
overcome thermal motion
– Condensation – gas molecules slow until
attraction pulls them together into liquid
• Intermolecular forces < intramolecular
forces (bonds)
• Stronger intermolecular force d
higher boiling and melting points
Types of Intermolecular
Forces
•
•
•
•
•
Ion - dipole
Dipole - dipole *
Dipole - induced dipole*
Dispersion (London) forces*
Hydrogen bond
Ion-dipole interactions
Cations generally have stronger interactions
because they are smaller, with more concentrated
charge than anions.
* van der Waals force
Dipole-Dipole
Pairs of cation-anion attractions align polar molecules
Ion-Induced Dipole &
Dipole-Induced Dipole
Non-polarized
atom
Polarized
atoms
• Induced dipole =
separation of + and
- charges in atom
or nonpolar
molecule caused
by proximity of ion
or polar molecule
Induced Dipoles
Instantaneous dipole –
induced dipole attraction
• Polarizability - ease with which edistribution of atom can be distorted
• More e-s d greater polarizability
• Often anions, with more diffuse charge,
are more polarizable than cations
• Molecules w/ dipoles can be polarized,
causing them to become more polar
• Instantaneous dipole – any atom or
non-polar molecule is briefly polar
because of separation between e- and +
nucleus
• Instantaneous dipole can induce dipole
in adjacent molecule
• Attraction between instantaneous and
induced dipole is very weak
• London or dispersion forces
Dispersion forces
Melting Points of Nonpolar CX 4
COMPOUND
CH4
CF4
CCl4
CBr4
CI4
MELTING POINT
(OC)
-182.5
-150.0
- 23.0
90.0
171.0
•Larger molecule d higher melting point
•Why?
• Occur in all substances (polar and non-polar)
• Allow non-polar molecules to condense
• Increase w/ polarizability, which increases with
molar mass
• Polar molecules have dipole-dipole attractions
as well as dispersion forces
• For polar molecules, larger dipole moment d
stronger attraction.
• For any molecule, larger molar mass (more e-)
d greater intermolecular attraction
• For a large molecule, dispersion force can
exceed dipole-dipole attraction
Hydrogen Bond
Approximate energies in kJ/mol
Compare to covalent single bonds, 200-600 kJ/mol
• Special dipole-dipole interaction between
H on N, O or F and a nearby N, O or F
– weak H-bonds with other elements such as
S and P
25
• H-bonds are strong intermolecular
attractions
• Large effects on intermolcular
organization
5
5-35
– higher b.p. and m.p.
– alignment of protein molecules
15-20
up to 45
weak
F
H
F
-
HO
169
H
OH
-
85
Identify “intermolecular” forces
Strong hydrogen bonds
Weak or no hydrogen bonds
•
•
•
•
•
•
•
H2O
CH2 Cl2
KBr
F- + H2O
I2
CH3OH
PCl3
•
•
•
•
•
•
•
C6 H6
SiH4
Fe
N(CH3) 3
CS2
BCl3
Na+ + NH3
11.3 Properties of Liquids
Anisotropic forces –
• Surface Tension – differ with direction
amount of energy
required to increase
surface area of
liquid by a unit area
• Interior molecules
are attracted in all
directions
• Surface molecules
are attracted
selectively into liquid
Isotropic forces –
same in all directions
• Measure of a fluid’s resistance to flow
• High viscosity (“thick”) = slow flow
• Strong intermolecular forces d high
viscosity
– glucose (typical sugar) is viscous because
of intermolecular hydrogen bonding
HC
OH
H
OH
OH
C
C
C
C
H
OH
H
H
CH2 O H
Viscosity
Liquid
Water (H2 O)
Ethanol (C2 H5OH)
Adhesion > cohesion
H 2O in glass
Concave meniscus
Cohesion > adhesion
Hg in glass
Convex meniscus
Viscosity
Viscosity
O
Capillary action
• Cohesion – intermolecular attraction between like
molecules (pure liquid)
• Adhesion – intermolecular attraction between unlike
molecules (liquid and its container)
• Newtonian Fluids –
viscosity is
independent of shear
rate
• Non- Newtonian
Fluids (examples)
– Thixotropic – viscosity
decreases with time
under constant shear
(gel-flow paint)
– Dilatant – viscocity
increases with shear
rate (Silly Putty)
Liquid
Water (H2O)
Ethanol (C2H5 OH)
Viscosity
(N s/m 2)
1.01 x 10-3
1.20 x 10-3
Glycerol
(HOCH2CHOHCH2 OH)
1.49
Blood
4 x 10-3
More hydrogen bonds
d more viscous liquid
Water
Viscosity
(N s/m 2)
1.01 x 10-3
1.20 x 10-3
Glycerol
(HOCH2CHOHCH2OH)
1.49
Blood
4 x 10- 3
More hydrogen bonds d more viscous liquid
• Large specific heat – bodies of water
moderate climate by absorbing and
releasing heat
• Liquid is denser than solid (ice floats)
• H2O can form 2 H-bonds per molecule,
leading to a very open solid structure.
Ice
Large open cavities in
solid ice are filled in
liquid water,
increasing density
11.4 Crystal Structure
• Crystalline solid – has rigid, longrange order
• Amorphous solid – lacks welldefined arrangement and long range
order
• Unit cell – basic repeating structural
unit of a crystalline solid
Lattice point – atom, molecule, or fixed
arrangement of atoms
•Coordination number (CN) – number of
atoms (ions) surrounding an atom (ion) in a
crystal lattice
•CN of an atom is a simple cubic lattice is 6
•1/8 of each corner atom is inside the unit cell
How many atoms are in a body -centered
cubic unit cell?
Corner atom is
shared by 8
cells.
Face-centered atom
is shared by two
cells.
Edge atom is shared
by 4 cells
Closest packing
• (a) Most efficient
arrangement of
spheres
• (b) 2nd layer
• (c) Hexagonal
closest packing
– ABABA...
• (d) Cubic closest
packing
– ABCABC...
X-Ray Diffraction
• Scattering of X-rays by the atoms of a
crystalline solid
• Atoms in a crystal absorb then re-emit
X-radiation
• When scattered X-rays from adjacent
layers of atoms are in phase,
constructive interference occurs at
certain point in space
8 corner atoms: 8 x 1/8 = 1
1 central atom: 1 x 1 = 1
d 1+1 = 2 atoms in body -centered cubic
cell
Hexagonal vs. cubic
close packing
• (a) Hexagonal close
packing
– ABABA...
• (b) Cubic close
packing
– ABCABC...
• Cubic close packing
is identical to facecentered cubic
packing
A
B
A
A
B
C
A
X-Ray diffraction by a crystal
X-rays of wavelength 0.0900 nm are used
to analyze a metallic crystal. 1st-order
diffraction (n = 1) occurs at 15.2°. What
is the layer spacing? What is the 2ndorder angle (n = 2)?
nλ = 2d sinθ
When additional distance traveled by an X-ray
(BC + CD = 2d sinθ) equals integral number of
wavelengths (nλ), waves are in phase
nλ = 2d sinθ (n = 1, 2, 3...)
11.6 Types of Crystals
• Ionic Crystals – array of anions
and cations
– high-melting
– brittle
•Covalent crystals – 3-D network of covalent
bonds
–high-melting, often hard
–diamond, graphite and quartz (SiO2)
154 pm
• NaCl, Li2O, CaF2, MgO
• Ionic bonds become stronger but
more covalent as charges increase
142 pm
Diamond: sp3 carbon
Types of Crystals
• Molecular crystals – lattice points
occupied by molecules, held together by
van der Waals forces and/or H-bonding
– low -melting (<100 oC)
• Metallic crystals – metal atoms at
lattice points
– dense, conductive, shiny etc.
– “metal ions in sea of electrons”
Graphite: sp2 carbon
Amorphous solids
• Lack a regular, three-dimensional
arrangement of atoms
• Glass – (optically transparent) fusion
product of (inorganic) materials cooled
to non-crystalline rigid state
• Common glass is mainly SiO2 (quartz),
w/ B 2O3 (Pyrex) or CaO, Na2O (lime
glass)
• Transition metal oxides color glass
11.8 Phase Changes
• Transformation between phases
• Energy (usually heat) is added or
removed
– Liquid-vapor equilibrium
– Liquid-solid equilibrium
– Solid-vapor equilibrium
Higher T d more KE d faster evaporation
Liquid - Vapor Equilibrium
• Evaporation (vaporization) –
process in which a liquid is
transformed into a gas
• Occurs when molecules have
enough energy to escape from the
liquid’s surface
(Equilibrium) Vapor Pressure
Vacuum
Low T d
High T d
fewer energetic molecules more energetic molecules
Heat of Vaporization, ∆Hvap
• E required to vaporize 1 mole of a liquid
X(l) d X(g)
∆Hvap
• Stronger intermolecular forces d larger
∆Hvap
• Table 11.6
• Everyday experience
– evaporating water or alcohol makes you feel
cold
– heat required to vaporize liquid comes from
your skin
• Pressure exerted by
evaporated gas molecules
above a liquid
• When rate of evaporation =
rate of condensation,
dynamic equilibrium is
reached
• P at this point =
(equilibrium) vapor
pressure
Hg
Quantitative relationship between vapor
pressure (P) and temperature (T)
• Clausius-Clapeyron Equation
ln P = -
∆Hvap
+C
RT
–at a single T; C is a constant
–plot log P vs. 1/T to find ∆Hvap
ln
P1
P2
=
∆Hvap
R
1
T2
-
1
T1
•Form to compare two temperatures
Critical Points
Boiling Point
• T at which vapor pressure of a
liquid = applied pressure
• Stronger intermolecular forces d
higher boiling point
• Normal boiling point – defined at 1
atm
• Critical Temperature (Tc) – T above
which gas cannot be liquefied at any
pressure
• Critical Pressure (P c) – minimum P
needed to liquefy a gas at Tc
• Supercritical fluids (phase above Tc) are
industrially important
– supercritical CO2 is used to decaffeinate
coffee, extract oils from grain, dry-clean
clothing, etc.
(Molar) Heat of Fusion, ∆Hfus
Liquid-Solid Equililbrium
• Freezing – phase change from liquid to
solid
• Melting or Fusion – phase change from
solid to liquid
• Melting point (freezing point) – T at
which the solid and liquid phases coexist
in equilibrium
• Normal m.p. (f.p.) – defined at 1 atm
• E required to melt one mole of a
solid
• Table 11.8
• ∆Hfus < ∆Hvap
– molecules are closely packed in the
liquid and solid states, but widely
separated in the gas state
How much heat is required to convert 425 g
of ice at -15°C to steam at 125°C? (specific
heats: ice = 2.03 J/g°C, water = 4.18 J/g°C,
steam = 1.99 J/g°C; ∆Hfus = 6.01 kJ/mol; ∆Hvap
= 40.79 kJ/mol)
Σ heats of
segments
q = ms∆t
= total
heat
Compare slopes:
s (solid > liquid > gas)
126 oC
5
100 oC
q4 = n∆ Hvap
Compare lengths:
∆H fus < ∆H vap
0 oC
-10 oC
q1 = ms∆t
q3 = ms∆t
q2 = n∆ Hfus
Solid-Vapor Equilibrium
• Sublimation – conversion of solid
directly to vapor
• Deposition – vapor to solid
• (Molar) heat of sublimation, ∆Hsub =
E required to sublime one mole of
solid
∆Hsub = ∆Hfus + ∆Hvap
11.9 Phase Diagrams
H2O
•Triple point – condition
in which all 3 phases
are in equilibrium
CO2
•Slope of curve (line) between phases
d P dependence of transition
•mp ice decreases at higher P
•mp CO2 increases at higher P
•Evaporations (up)
are endothermic
•Condensations
(down) are
exothermic