Transcript Ch 11: Intermolecular Forces

Chapter 11 Intermolecular Forces
11.1: Intermolecular Forces (IMF)
 IMF < intramolecular forces (covalent,
metallic, ionic bonds)
 IMF strength: solids > liquids > gases
 Boiling points and melting points are good
indicators of relative IMF strength.
2
11.2: Types of IMF
3
1. Electrostatic forces: act over larger distances in
accordance with Coulomb’s law
Q Q
F
d2
a. Ion-ion forces: strongest; found in ionic crystals
(i.e. lattice energy)
b. Ion-dipole: between an ion and a dipole (a
neutral, polar molecule/has separated
partial charges)
 Increase with increasing polarity of molecule
and increasing ion charge.
Q Q
F
d2
Ex: Compare IMF in Cl- (aq) and S2- (aq).
d
d
d
d
Cl-
d
d
d
d
d
<
d
d
d
d
S2-
d
d
d
d
d
4
c. Dipole-dipole: weakest
electrostatic force; exist
between neutral polar
molecules
 Increase with increasing
polarity (dipole moment) of
molecule
Ex: What IMF exist in NaCl (aq)?
5
d. Hydrogen bonds (or H-bonds):

H is unique among the elements because it has a single
e- that is also a valence e-.
– When this e- is “hogged” by a highly EN atom (a very
polar covalent bond), the H nucleus is partially
exposed and becomes attracted to an e--rich atom
nearby.
6
7
 H-bonds form with H-X•••X', where X and X'
have high EN and X' possesses a lone pair
of e X = F, O, N (since most EN elements) on
two molecules:
F-H
:F
O-H
:O
N-H
:N
8
 * There is no strict cutoff for the ability to
form H-bonds (S forms a biologically
important hydrogen bond in proteins).
 * Hold DNA strands together in double-helix
Nucleotide pairs form Hbonds
DNA double helix
 H-bonds explain why ice is less dense than
water.
9
Ex: Boiling points of nonmetal hydrides
 Polar molecules have
higher BP than nonpolar
molecules
∴ Polar molecules have
stronger IMF
Boiling Points (ºC)
Conclusions:
 BP increases with
increasing MW
∴ Heavier molecules
have stronger IMF
 NH3, H2O, and HF have unusually high BP.
∴ H-bonds are stronger than dipole-dipole IMF
10
11
2. Inductive forces:
 Arise from distortion of the e- cloud induced
by the electrical field produced by another
particle or molecule nearby.
 London dispersion: between polar or
nonpolar molecules or atoms
– * Proposed by Fritz London in 1930
– Must exist because nonpolar molecules form
liquids
Fritz London
(1900-1954)
How they form:
12
1. Motion of e- creates an instantaneous dipole
moment, making it “temporarily polar”.
2. Instantaneous dipole moment induces a dipole
in an adjacent atom
•
* Persist for about 10-14 or 10-15 second
Ex: two He atoms
* Geckos!
13
 Geckos’ feet make use of
London dispersion forces to
climb almost anything.
 A gecko can hang on a glass
surface using only one toe.
 Researchers at Stanford
University recently developed a
gecko-like robot which uses
synthetic setae to climb walls
http://en.wikipedia.org/wiki/Van_der_Waals%27_force
London dispersion forces increase with:

Increasing MW, # of e-, and # of atoms (increasing # of eorbitals to be distorted)
Boiling points:
Effect of MW:
Effect of # atoms:
pentane 36ºC
Ne
–246°C
hexane
69ºC
CH4 –162°C
heptane 98ºC
??? effect:
H2O
D2O

14
100°C
101.4°C
“Longer” shapes (more likely to interact with other molecules)
C5H12 isomers: 2,2-dimethylpropane 10°C
pentane
36°C
Summary of IMF
Van der Waals forces
Ex: Identify all IMF present in a pure sample of each
substance, then explain the boiling points.
BP(⁰C)
HF
HCl
HBr
HI
20
-85
-67
-35
16
IMF
Explanation
London, dipoledipole, H-bonds
Lowest MW/weakest London,
but most polar/strongest
dipole-dipole and has H-bonds
London, dipoledipole
Low MW/weak London,
moderate polarity/dipole-dipole
and no H-bonds
London, dipoledipole
Medium MW/medium London,
moderate polarity/dipole-dipole
and no H-bonds
London, dipoledipole
Highest MW/strongest
London, but least polar
bond/weakest dipole-dipole
and no H-bonds
11.3: Properties resulting from IMF
1. Viscosity: resistance of a liquid to flow
Viscosity depends on:
-the attractive forces between molecules
-the tendency of molecules to become
entangled
-the temperature
17
11.3: Properties resulting from IMF
•
Surface tension: energy required to
increase the surface area of a liquid
18
3. Cohesion: attraction of molecules for other
molecules of the same compound
4. Adhesion: attraction of molecules for a
surface
19
20
5. Meniscus: curved upper surface of a liquid
in a container; a relative measure of
adhesive and cohesive forces
Ex:
Hg
(cohesion rules)
H2O
(adhesion rules)
Phase Changes
• Surface molecules are only attracted inwards towards
the bulk molecules.
• Sublimation: solid  gas.
• Vaporization: liquid  gas.
• Melting or fusion: solid  liquid.
• Deposition: gas  solid.
• Condensation: gas  liquid.
• Freezing: liquid  solid.
Energy Changes Accompanying Phase Changes
• Energy change of the system for the above processes
are:
21
Intermolecular Forces Bulk and Surface
22
Phase Changes
Energy Changes Accompanying Phase Changes
–
–
–
–
–
–
Sublimation: Hsub > 0 (endothermic).
Vaporization: Hvap > 0 (endothermic).
Melting or Fusion: Hfus > 0 (endothermic).
Deposition: Hdep < 0 (exothermic).
Condensation: Hcon < 0 (exothermic).
Freezing: Hfre < 0 (exothermic).
• Generally heat of fusion (enthalpy of fusion) is less
than heat of vaporization:
– it takes more energy to completely separate molecules, than
partially separate them.
23
Phase Changes
Energy Changes Accompanying Phase Changes
• All phase changes are possible under the right
conditions (e.g. water sublimes when snow disappears
without forming puddles).
• The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
• The sequence
cool gas  condense  cool liquid  freeze  cool
solid
is exothermic.
24
Phase Changes
Energy Changes Accompanying Phase Changes
25
Phase Changes
Heating Curves
• Plot of temperature change versus heat added is a
heating curve.
• During a phase change, adding heat causes no
temperature change.
– These points are used to calculate Hfus and Hvap.
• Supercooling: When a liquid is cooled below its
melting point and it still remains a liquid.
• Achieved by keeping the temperature low and
increasing kinetic energy to break intermolecular
forces.
26
Phase Changes
Heating Curves
27
Heating Curve Illustrated
28
Phase Changes
Critical Temperature and Pressure
• Gases liquefied by increasing pressure at some
temperature.
• Critical temperature: the minimum temperature for
liquefaction of a gas using pressure.
• Critical pressure: pressure required for liquefaction.
29
Critical Temperature, Tc
30
Transition to Supercritical CO2
31
Supercritical CO2 Used to Decaffeinate Coffee
32
Vapor Pressure
Explaining Vapor Pressure on the Molecular
Level
• Some of the molecules on the surface of a liquid have
enough energy to escape the attraction of the bulk
liquid.
• These molecules move into the gas phase.
• As the number of molecules in the gas phase
increases, some of the gas phase molecules strike the
surface and return to the liquid.
• After some time the pressure of the gas will be
constant at the vapor pressure.
33
Gas-Liquid Equilibration
34
Vapor Pressure
Explaining Vapor Pressure on
the Molecular Level
• Dynamic Equilibrium: the
point when as many molecules
escape the surface as strike the
surface.
• Vapor pressure is the pressure
exerted when the liquid and
vapor are in dynamic
equilibrium.
35
Vapor Pressure
Volatility, Vapor Pressure, and Temperature
• If equilibrium is never established then the liquid
evaporates.
• Volatile substances evaporate rapidly.
• The higher the temperature, the higher the average
kinetic energy, the faster the liquid evaporates.
36
Liquid Evaporates when no Equilibrium is Established
37
Vapor Pressure
Volatility, Vapor Pressure, and Temperature
38
Vapor Pressure
Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals the
vapor pressure.
• Temperature of boiling point increases as pressure
increases.
• Two ways to get a liquid to boil: increase temperature
or decrease pressure.
– Pressure cookers operate at high pressure. At high pressure
the boiling point of water is higher than at 1 atm.
Therefore, there is a higher temperature at which the food is
cooked, reducing the cooking time required.
• Normal boiling point is the boiling point at 760 mmHg
(1 atm).
39
Phase Diagrams
• Phase diagram: plot of pressure vs. Temperature
summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams
tell us which phase will exist.
• Features of a phase diagram:
– Triple point: temperature and pressure at which all three
phases are in equilibrium.
– Vapor-pressure curve: generally as pressure increases,
temperature increases.
– Critical point: critical temperature and pressure for the gas.
– Melting point curve: as pressure increases, the solid phase is
favored if the solid is more dense than the liquid.
– Normal melting point: melting point at 1 atm.
40
Phase Diagrams
• Any temperature and pressure combination not on a
curve represents a single phase.
41
Phase Diagrams
The Phase Diagrams of H2O and CO2
• Water:
– The melting point curve slopes to the left because ice is less
dense than water.
– Triple point occurs at 0.0098C and 4.58 mmHg.
– Normal melting (freezing) point is 0C.
– Normal boiling point is 100C.
– Critical point is 374C and 218 atm.
• Carbon Dioxide:
– Triple point occurs at -56.4C and 5.11 atm.
– Normal sublimation point is -78.5C. (At 1 atm CO2
sublimes it does not melt.)
– Critical point occurs at 31.1C and 73 atm.
42
Phase Diagrams
The Phase Diagrams of H2O and CO2
43
11.4: Phase Changes
Processes:
 Endothermic: melting,
vaporization, sublimation
 Exothermic: condensation,
freezing, deposition
I2 (s) and (g)
44
Microchip
Water: Enthalpy diagram or heating curve
Q  mH
Q  m cT
45
Q  m(1.87 J/gC)T
Q  m(2260J/g)
Q  m(4.18J/gC)T
Q  m(334 J/g)
Q  m(2.06 J/gC)T
11.5: Vapor pressure
46
Pressure cooker ≈ 2 atm
Normal BP = 1 atm
10,000’ elev ≈ 0.7 atm
29,029’ elev (Mt. Everest)
≈ 0.3 atm
 A liquid will boil when the vapor pressure equals
the atmospheric pressure, at any T above the
triple point.
11.6: Phase diagrams: CO2
47
 Lines: 2 phases exist in
equilibrium
 Triple point: all 3 phases exist
together in equilibrium (X on
graph)
 Critical point, or critical
temperature & pressure:
highest T and P at which a
liquid can exist (Z on graph)
Temp (ºC)
 For most substances, inc P will cause a gas to condense
(or deposit), a liquid to freeze, and a solid to become more
dense (to a limit.)
Phase diagrams: H2O
• For H2O, inc
P will cause
ice to melt.
48
49
*
*
50
11.7-8: Structures of solids
51
 Amorphous: without orderly structure
Ex: rubber, glass
 Crystalline: repeating structure; have many
different stacking patterns based on
chemical formula, atomic or ionic sizes, and
bonding
Cubic Unit Cells in Crystalline Solids
• Primitive-cubic shared atoms are located only at
each of the corners. 1 atom per unit cell.
• Body-centered cubic 1 atom in center and the
corner atoms give a net of 2 atoms per unit cell.
• Face-centered cubic corner atoms plus halfatoms in each face give 4 atoms per unit cell.
52
Common Lattice Structures
53
Types of Crystalline Solids
Type
Atomic
Particles
Atoms
Forces
London
dispersion
Notable
properties
Poor
conductors
Very low
MP
Examples
Ar (s),
Kr (s)
A small (~2 cm long) piece of rapidly
melting argon ice (the liquid is flowing off at
the bottom) which has been frozen by
allowing a slow stream of the gas to flow
into a small graduated cylinder which was
immersed into a cup of liquid nitrogen
Molecular
crystals
Poor
London
Molecules
conductors
dispersion,
(polar or dipole-dipole, Low to
non-polar)
moderate
H-bonds
MP
Carbon dioxide, dry ice
(g at room T)
Ice
(liq at room T)
SO2(s)
CO2 (s),
C12H22O11,
H2O (s)
Sucrose
(liq at room T)
Covalent
(a.k.a.
covalent
network)
Atoms
bonded in a
covalent
network
Graphite
Very hard
Very high
MP
Covalent
bonds Generally
insoluble
Variable
conductivity
Diamond
C (diamond
& graphite)
SiO2
(quartz)
Ge, Si, SiC,
BN
SiO2
Ionic
Anions
and
cations
Crystals
shatter!
Ion-ion (ionic
bonding)
High Lattice
Energy
Hard & brittle
High MP,BP
Poor
conductors
Some
solubility in
H2O
NaCl,
Ca(NO3)2
Metallic
Excellent
Metallic
conductors
Metal
bonds
Malleable
cations in a
Usually
diffuse,
faceDuctile
delocalized centered or High but
body
e- cloud
wide range
centered
of MP
Cu, Al, Fe
(hard)
Alloys
Pb, Au, Na
(soft)
Overall
59
• Physical properties depend on these forces. The stronger
the forces between the particles,
• (a) the higher the melting point.
• (b) the higher the boiling point.
• (c) the lower the vapor pressure (partial pressure of vapor
in equilibrium with liquid or solid in a closed container at a
fixed temperature).
• (d) the higher the viscosity (resistance to flow).
• (e) the greater the surface tension (resistance to an
increase in surface area).
Practice
60
• Determine the type of solid and the forces
holding the particles together
•
•
•
•
•
•
•
•
SiO2
NaNO3
C2H6
CH3OH
C(diamond)
Al
Kr
H2O
Covalent Network
Ionic
Molecular
Molecular
Covalent Network
Metallic
Atomic (Molecular)
Molecular
Covalent Bonds
Electrostatic Att.
Dispersion
Dispersion, Dipole-Dipole, H-Bond
Covalent Bonds
Metallic
Dispersion
Dispersion, Dipole-Dipole, H-Bond
Extra Material
61
• The following pages contain some additional
material and review items
Examples
62
Ionic Solids
stable, high melting points
 held together by strong electrostatic forces
between oppositely charged ions
 larger ions are arranged in closest packing
arrangement
 smaller ions fit in the holes created by the
larger ions

63
Cubic Unit Cells in Crystalline Solids
• Primitive-cubic shared atoms are located only at each of the
corners. 1 atom per unit cell.
• Body-centered cubic 1 atom in center and the corner atoms
give a net of 2 atoms per unit cell.
• Face-centered cubic corner atoms plus half-atoms in each
face give 4 atoms per unit cell.
8–64
Chapter 11-64
Common Lattice Structures
65
Calculations involving the Unit Cell
•
•
The density of a metal can be calculated if we know the length of the side of a unit
cell.
The radius of an metal atom can be determined if the unit cell type and the density
of the metal known
–
Relationship between length of side and radius of atom:
• Primitive
2r = l; FCC:
r
2
l
4
BCC
r
3
l
4
E.g. Polonium crystallizes according to the primitive cubic structure. Determine its
density if the atomic radius is 167 pm.
E.g.2 Calculate the radius of potassium if its density is 0.8560 g/cm3 and it has a BCC
crystal structure.
8–66
Chapter 11-66
Figure 11.31
• Length of sides a, b, and c as well as angles a, b, g vary to
give most of the unit cells. Return to unit cells
8–67
Chapter 11-67
Unit Cells in Crystalline Solids
• Metal crystals made up of atoms in regular arrays – the smallest
of repeating array of atoms is called the unit cell.
• There are 14 different unit cells that are observed which vary in
terms of the angles between atoms some are 90°, but others are
not. Go to Figure 11.31
8–68
Chapter 11-68
Packing of Spheres and the Structures of Metals
•
•
•
Arrays of atoms act as if they are spheres. Two or more layers produce
3-D structure.
Angles between groups of atoms can be 90° or can be in a more
compact arrangement such as the hexagonal closest pack (see below)
where the spheres form hexagons.
Two cubic arrays one directly on top of the other produces simple cubic
(primitive) structure.
– Each atom has 6 nearest neighbors (coordination number of 6); nearest
neighbor is where an atom touches another atom.
– 54% of the space in a cube is used.
•
Offset layers produces a-b-a-b arrangement since it takes two layers to
define arrangement of atoms.
– BCC structure an example.
– Coordination # is 8.
8–69
Chapter 11-69
Packing of Spheres and the Structures of Metals
• FCC structure has a-b-c-a-b-c
stacking. It takes three layers to
establish the repeating pattern
and has 4 atoms per unit cell and
the coordination number is 12.
8–70
Chapter 11-70
Metallic Crystals
 can
be viewed as metals
atoms (spheres) packed
together in the closest
arrangement possible
 closest packing- when each
sphere has 12 neighbors



6 on the same plane
3 in the plane above
3 in the plane below
71
Bonding of Metals
 the
highest energy level for most
metal atoms does not contain many
electrons
 these vacant overlapping orbitals
allow outer electrons to roam freely
around the entire metal
72
Bonding of Metals
 these
roaming electrons
form a sea of electrons
around the metal atoms
 malleability and ductility


bonding is the same in every direction
one layer of atoms can slide past another
without friction
 conductivity

from the freedom of electrons to move
around the atoms
73
Metal Alloys
substance that is a mixture of elements and
has metallic properties
 substitutional alloy




host metal atoms are replaced by other metal
atoms
happens when they have similar sizes
interstitial alloy


metal atoms occupy spaces created between host
metal atoms
happens when metal atoms have large difference
in size
74
Examples
 Brass


substitutional
1/3 of Cu atoms
replaced by Zn
 Steel



interstitial
Fe with C atoms in
between
makes harder and
less malleable
75
Chapter 11 Overview
• Changes of State
– Phase transitions
– Phase Diagrams
• Liquid State
Exam on Friday
We will begin Chp 14
Thursday
– Properties of Liquids; Surface tension and viscosity
– Intermolecular forces; explaining liquid properties
• Solid State
–
–
–
–
–
Classification of Solids by Type of Attraction between Units
Crystalline solids; crystal lattices and unit cells
Structures of some crystalline solids
Calculations Involving Unit-Cell Dimensions
Determining the Crystal Structure by X-ray Diffraction
8–76
Chapter 11-76
Comparison of Gases, Liquids and Solids
– Gases are compressible fluids. Their molecules are widely
separated.
– Liquids are relatively incompressible fluids. Their molecules are
more tightly packed.
– Solids are nearly incompressible and rigid. Their molecules or ions
are in close contact and do not move.
8–77
Figure 11.2 States of Matter
Chapter 11-77
Phase Transitions
• Melting: change of a solid to a liquid. H2O(s)  H2O(l)
• Freezing: change a liquid to a solid. H2O(l)  H2O(s)
• Vaporization: change of a solid or
liquid to a gas. Change of solid to
vapor often called sublimation.
H2O(l)  H2O(g)
or
H2O(s)  H2O(g)
• Condensation: change of a gas to a H2O(g)  H2O(l)
liquid or solid. Change of a gas to a or
H2O(g)  H2O(s)
solid often called deposition.
8–78
Chapter 11-78
Vapor Pressure
• In a sealed container,
some of a liquid
evaporates to establish a
pressure in the vapor
phase.
• Vapor pressure: partial
pressure of the vapor over
the liquid measured at
equilibrium and at some
temperature.
• Dynamic equilibrium
8–79
Chapter 11-79
Temperature Dependence of Vapor Pressures
• The vapor pressure above
the liquid varies
exponentially with changes
in the temperature.
• The Clausius-Clapeyron
equation shows how the
vapor pressure and
temperature are related. It
can be written as:
ln P  
Hvap
RT
1
 C
T
8–80
Chapter 11-80
Clausius – Clapeyron Equation
•
•
•
A straight line plot results when ln P
vs. 1/T is plotted and has a slope of
Hvap/R.
Clausius – Clapeyron equation is
true for any two pairs of points.
Hvap 1
ln P  
 C
RT
T
Write the equation for each and
combine to get:
Hvap  1
P
1
ln 2 
   
P1
RT
 T1 T2 
8–81
Chapter 11-81
Using the Clausius – Clapeyron Equation
• Boiling point the temperature at which the vapor
pressure of a liquid is equal to the pressure of the
external atmosphere.
• Normal boiling point the temperature at which the
vapor pressure of a liquid is equal to atmospheric
pressure (1 atm).
E.g. Determine normal boiling point of chloroform if its heat of
vaporization is 31.4 kJ/mol and it has a vapor pressure of 190.0
mmHg at 25.0°C.
E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it
has a vapor pressure of 100.0 mmHg. What is the heat of
8–82
vaporization?
Chapter 11-82
Energy of Heat and Phase Change
• Heat of vaporization: heat
needed for the vaporization of a
liquid.
H2O(l) H2O(g) H = 40.7 kJ
• Heat of fusion: heat needed
for the melting of a solid.
H2O(s) H2O(l) H = 6.01 kJ
• Temperature does not change
during the change from one
phase to another.
E.g. Start with a solution consisting of 50.0 g of H2O(s) and 50.0 g of
H2O(l) at 0°C. Determine the heat required to heat this mixture to
8–83
100.0°C and evaporate half of the water.
Chapter 11-83
Phase Diagrams
• Graph of pressure-temperature
relationship; describes when
1,2,3 or more phases are present
and/or in equilibrium with each
other.
• Lines indicate equilibrium state
two phases.
• Triple point- Temp. and press.
where all three phases co-exist in
equilibrium.
• Critical temp.- Temp. where
substance must always be gas,
no matter what pressure.
• Critical pressure- vapor pressure at critical temp.
• Critical point- point where system is at its critical pressure and
temp.
8–84
Chapter 11-84
Properties of Liquids
• Surface tension: the energy
required to increase the surface area
of a liquid by a unit amount.
• Viscosity: a measure of a liquid’s
resistance to flow.
• Surface tension: The net pull toward
the interior of the liquid makes the
surface tend to as small a surface
area as possible and a substance
does not penetrate it easily.
• Viscosity: Related to mobility of a
molecule (proportional to the size
and types of interactions in the
liquid).
– Viscosity decreases as the temperature increases since increased
temperatures tend to cause increased mobility of the molecule.
8–85
Chapter 11-85
Intermolecular Forces
• Intermolecular forces: attractions and repulsions between
molecules that hold them together.
• Intermolecular forces (van der Waals forces) hold molecules
together in liquid and solid phases.
– Ion-dipole force: interaction between an ion and partial charges in
a polar molecule.
– Dipole-dipole force: attractive force between polar molecules with
positive end of one molecule is aligned with negative side of other.
– London dispersion Forces: interactions between instantaneously
formed electric dipoles on neighboring polar or nonpolar molecules.
– Polarizability: ease with which electron cloud of some substance
can be distorted by presence of some electric field (such as another
dipolar substance). Related to size of atom or molecule. Small
atoms and molecules less easily polarized.
8–86
Chapter 11-86
Boiling Points vs. Molecular Weight
• Hydrogen bonds: the
interaction between hydrogen
bound to an electronegative
element (N, O, or F) and an
electron pair from another
electronegative element.
Hydrogen bonding is the
dominate force holding the
two DNA molecules together
to form the double helix
configuration of DNA.
8–87
Chapter 11-87
Comparisonof Energies for Intermolecular Forces
Interaction Forces
:Approximate Energy
Intermolecular
London
1 – 10 kJ
Dipole-dipole
3 – 4 kJ
Ion-dipole
5 – 50 kJ
Hydrogen bonding
10– 40 kJ
Chemical bonding
Ionic
100 – 1000 kJ
Covalent
100 – 1000 kJ
8–88
Chapter 11-88
Structure of Solids
• Types of solids:
– Crystalline – a well defined arrangement of atoms; this
arrangement is often seen on a macroscopic level.
• Ionic solids – ionic bonds hold the solids in a regular
three dimensional arrangement.
• Molecular solid – solids like ice that are held together by
intermolecular forces.
• Covalent network – a solid consists of atoms held
together in large networks or chains by covalent
networks.
• Metallic – similar to covalent network except with metals.
Provides high conductivity.
– Amorphous – atoms are randomly arranged. No order exists8–89
in the solid. Example: glass
Chapter 11-89