Transcript Chapter 9: Intermolecular Attractions and the Properties

Lecture Notes
Chem 150 - K. Marr
Chapter 12
Intermolecular Attractions &
the Properties of Liquids & Solids
Silberberg 3 ed
Intermolecular Forces: Liquids, Solids, and Phase Changes
12.1 An Overview of Physical States and Phase Changes
12.2 Quantitative Aspects of Phase Changes
12.3 Types of Intermolecular Forces
12.4 Properties of the Liquid State
12.5 The Uniqueness of Water
12.6 The Solid State: Structure, Properties, and Bonding
12.7 Advanced Materials
Table 12.1
A Macroscopic Comparison of Gases, Liquids, and Solids
State
Shape and Volume
Compressibility Ability to Flow
Gas
Conforms to shape and
volume of container
high
high
Conforms to shape of
container; volume limited by
surface
very low
moderate
Maintains its own shape and
volume
almost
none
almost
none
Liquid
Solid
Chapter 12: Intermolecular Attractions &
the Properties of Liquids & Solids
1.
Why do the gas laws work with almost any gas?
2. Gases are alike
»
Mainly empty space  Weak intermolecular attractions
Solid: Maintains its
own shape and
volume
Liquid: Conforms to
shape of container;
volume limited by
surface
Gas: Conforms
to shape and
volume of
container
Why aren’t there liquid laws and solid laws?
1.
Little empty space between molecules:
Particles close together in S’s and L’s
2.
Stronger and quite varied intermolecular
attractions in S’s & L’s than in gases...Why?
» Attractions decrease as distance between molecules
increase:
I.M.F. a 1/ d2
3.
Attractions dependent on chemical composition
– Polar vs Nonpolar molecules
e.g. Water vs Carbon Dioxide (sublimes. @ -58.5 oC)
Intermolecular Attractions:
Bonds between Molecules
1.
2.
3.
Much weaker than Chemical Bonding
within molecules
Chemical Bonds (ionic and covalent)
determine chemical properties
Intermolecular bonds determine physical
properties
e.g. density, mp, bp, solubility, vapor pressure, etc.
Kinds of Intermolecular Attractions
1.
2.
3.
4.
5.
Dipole-Dipole Attractions
Hydrogen Bonds (H-FON Bonds)
London Forces (dispersion forces)
Ion-Dipole (e.g. spheres of hydration) Chapter 13
Induced Dipole forces Chapter 13
a. Ion induced
b. Dipole induced
How do IMF’s affect Heats of
Vaporization and Fusion?
Relative magnitudes of
forces in Molecular compounds
Covalent bonds >>>>> Hydrogen bonding >>
Dipole-dipole interactions >>>>> London forces
What kind of IMF??
Dipole-Dipole Attractions
1.
Dipoles are polar molecules
– Molecules w/ polar bonds and asymmetric distribution
of charge
– What determines if a bond is nonpolar, polar or ionic?
– What determines if a molecule with polar bonds is
polar or nonpolar
2.
3.
Much weaker than covalent bonds
Important in maintaining the shape of many
biological molecules: e.g. Proteins
Hydrogen Bonds (H-FON Bonds)
1.
Special kind of dipole-dipole interaction
»
2.
3.
4.
Found in HF and molecules containing O-H or N-H
bonds
5x’s the strength of a typical dipole-dipole bond
~ 5% the strength of a covalent bond
Important in biological molecules
e.g. DNA, proteins
Myoglobin
Hemoglobin
Hydrogen bonding is responsible
for the expansion of water when it freezes.
Exercise: Which of the following molecules
display hydrogen bonding?
1.
2.
3.
4.
Methane, CH4
methyl ether, CH3OCH3
Hydrogen peroxide, H2O2
methyl alcohol, CH3OH
London Forces:
Attractions between temporary dipoles
Electrostatic
Attraction
Random movement of electrons
may cause temporary charge imbalances
London Force
London
(dispersion)
forces
between
nonpolar
molecules
Why are London
forces the greatest in
large molecules?
London Forces (dispersion forces)
exist in all molecules
1.
2.
Ave. strength <<< Dipole-Dipole
interactions
Result from temporary charge imbalances
a. Due to the random movement of electrons
b. Nucleus of one atom attracts electrons from a
neighboring atom.
c. At the same time, the electrons in one particle
repel the electrons in the neighbor and create a
short lived charge imbalance.
Relationship between atomic size and the
strength of London forces
1.
2.
Greatest in large atoms
» Electron clouds more easily distorted
» Halogens and Noble Gases BP increase w/
molar mass
Ion Induced Dipoles
» dipoles can be induced by ions
» attractions exist between ions and dipoles
London Forces
Effect of molecular surface area
Cyclopentane, BP = 49.3 oC
IMF’s in
Nonpolar Organic Molecules
1.
2.
3.
What kind of attractive forces are present?
What role do molecular size and surface area
play?
Linear molecules have more surface area than
if they are folded into a sphere.
• Linear molecules have higher melting and boiling
points because of the increased attractions.
Predicting the
Relative Boiling Points of Substances
1.
2.
The substance with the strongest
intermolecular attractions will have the
higher BP. Why?
More energy needed to separate molecules 
higher boiling temperature
e.g. Halogens and noble gases
Cooling Curve: H2O (g)  H2O (s)
What is happening in to KE and PE at each part of the curve?
DHvap = + 40.7 kJ/mol
DHfus = + 6.01 kJ/mol
Molar Heat of Vaporization, DHvap
Heat absorbed when one mole liquid is changed to
one mole of vapor at constant T and P
» Depends on strength of IMF’s
» Endothermic (results in PE elevation)
» For water: DHvap = + 40.7 kJ/mol @ 100oC
Molar Heat of fusion, DHvap
Heat absorbed when 1 mole solid is changed to 1
mole of liquid at constant T and P.
» Depends on strength of IMF’s
» Endothermic (results in PE elevation)
» For water: DHfus = + 6.01 kJ/mol
Quantitative Aspects of Phase Changes
Within a phase
A change in heat is associated with a change in average KE
and, therefore, a change in temperature.
q = (mass)(Specific Heat)(Dt)
During a phase change
A change in heat occurs at constant temperature, which is associated
with a change in PE, as the average distance between molecules
changes--Bond IMF formation is exothermic, IMF breaking is
endothermic
q = (moles of substance)(enthalpy of phase change)
Calculation of the Heat of Fusion of Ice
Use the data below to calculate the heat of fusion of water.
1.
–
–
–
–
2.
A piece of ice at zero Celsius melts in 100.0 g water until the
water’s temperature also becomes zero.
o
Initial water temp. = 44.0 C
Mass of ice that melted = 56.0 g.
o
Specific heat of water, CH2O = 4.184 J/g C
Calculate the % error and explain the source of the error.
DHfus H2O = + 6.01 kJ/mol
Application Questions
1.
2.
o
The molar heat of vaporization of water at 25 C is
43.99 kJ/mol. How many kilojoules of heat would be
o
required to vaporize 125 mL (125 g) of water at 25 C?
» Answer: 305kJ
How much heat would be needed to convert 125 mL
o
water at 25.oC to steam at 100.0 C? The heat of
vaporization of water at 100.0 oC is 40.657 kJ/mol and
o
the specific heat of water is 4.184 J/g C.
» Answer: 321 kJ
General Properties of Liquids and Solids
1.
2.
Macroscopic properties depend on
Microscopic properties
Microscopic properties of L’s and S’s
» Molecules tightly packed
» Strong intermolecular attractions
Macroscopic properties of S’s and L’s
1.
2.
Compressibility
» Little to none. Why?
Diffusion (ability to flow)
» Slow in liquids
(T6)
– Like moving in a crowded room
» Nonexistent in Solids
– Particles not free to move
Macroproperties Liquids
Why are raindrops spherical?
»
Increases stability by maximizing the number of
IMF’s and decreasing surface tension
»
Surface Tension =
• E needed to increase the surface area of a liquid by a
given amount (J/m2)
• Depends on nature of intermolecular forces
The Molecular Basis of Surface Tension
Surface molecules experience fewer intermolecular forces than interior molecules
Liquids minimize surface area by forming spherical surfaces  Lowers PE, thus
increases stability ( e.g. Raindrops, overfilled glass)
• Surface molecules are at higher P.E. than interior molecules
Recall: Bond formation results in P.E. lowering
Macroproperties Liquids
1.
2.
Wetting of a Surface by a Liquid
» Spreading of a Liquid across a surface
» Caused by attraction of liq. molecules to surface
molecules
Why are Liquids with a Low Surface Tension Good
Wetters?........
– Low Surface Tension means weak IM Forces
3.
Which wets solids surfaces better, hydrocarbons
(gasoline/oil ) or water?.......
Evaporation and Sublimation
Where molecules leave surface and enter vapor space
around them
» Evaporation: L  Vapor
» Sublimation: S  Vapor
Evaporation:
1.
2.
3.
4.
L Vapor
(T9)
Factors that affect the rate of evaporation
» Surface area, Temp., Strength of IMF’s
Why does evaporation occur at temp.’s below
the BP?
Why does an increase in temp. increase the
rate of evaporation?
Why does sweating cool you?
» Why does a fan cool you?
Effect of Temperature on the
Distribution of Molecular Speeds in a Liquid
Evaporation: L Vapor :
Application Questions (T10)
1.
2.
If the liquids are water and ethanol, which one is
liquid A? Liquid B?
Why does the evaporation of ethanol from the skin
feel cooler than the evap. of water?
» Which removes more K.E.?
3.
4.
Which liquid will evaporate more quickly, acetone or
ethanol?
Which is more difficult to clean up, a marine spill of
low or high MW crude oil?
Application Questions (T12)
1.
2.
How can snow or ice cubes in the
freezer disappear w/o melting?
Why do moth balls (naphthalene) need
to be replaced periodically?
Changes of State &
Dynamic Equilibrium
1.
Change of State (T11)
» Change from one physical state to another
S
L
G or S
G
– Occur under conditions of Dynamic Equilibriium
– Two opposing events occurring @ equal rates
2.
Application Questions
» Why does fanning cause you to feel cooler?
» Why does sweating in the tropics cool you less than
sweating in the desert?
Vapor Pressure of Solids & Liquids
1.
Vapor Pressure
Pressure exerted by a vapor in a closed flask in
equilibrium w/ its liquid (T11)
2.
Measurement of vapor pressure (T 13a)
Rateevap >
Ratecond
Rateevap =
Ratecond
Liquid – Gas Equilibrium
•Only Temperature
and Intermolecular
Forces affect Vapor
Pressure
•The % of molecules
with sufficient K.E.to
escape the liquid
surface increases
with Temperature
Factors affecting Vapor Pressure
1.
2.
Strength IMF’s (T13b)
Temperature
» At higher temps a higher % of the molecules have
suffiecient K.E. to escape the liquid surface
3.
Factors not Affecting VP
» Surface Area
– Increases the rate of evaporation and condensation
equally
» Amount of Liquid
– Evaporation occurs at surface
VP in Solids
1.
Due to vibrations of surface molecules 
Sublimation (T12)
Examples
• Snow and ice cubes
• Dry ice (solid carbon dioxide)
Boiling Points of Liquids
1.
Boiling Point
• Temp. at which VP of a liquid equals atmospheric
pressure (T14)
• Depends on
o
o
2.
3.
Strength of IMF’s
Atmospheric pressure
Why do liquids Boil?
Why do bubbles form on sides first?
Application Questions
1.
2.
3.
Explain why water can exist at temps above
its normal BP in a car’s radiator.
Explain why a pressure cooker can cook a
beef stew faster than in a normal pot.
Explain why water, hydrogen fluoride and
ammonia have much higher BP’s than one
might predict by size alone (T15)
Clausius-Clapeyron Equation
1.
Relates the Vapor Pressure of a liquid to the heat of
vaporation, -DHvap
Ln P = -DHvap/RT + constant
or
Ln (P1/P2) = -DHvap /R(1/T2 - 1/T1)
Note: P = Vap Pressure; R = 8.314 J/mol K
2.
Our next lab experiment will involve the use of the
Clausius-Clapeyron Equation to calculate the
DHvap for methanol and ethanol
Clausius-Clapeyron Eqn:
ln P = -DHvap/RT + c
P = Vapor Pressure
= 8.314 J/mol K
R
Dynamic Equilibrium and
Le Chatelier’s Principle
Le Chatelier’s Principle: When a stress is applied
to a system at equilibrium, the equilibrium will
shift to relieve that stress and, if possible, restore
equilibrium.
» Position of Equilibrium: Refers to relative
amount of reactants and products
Application Questions.................
1.
Le Chatelier’s Principle : Application Questions
1.
2.
3.
Use L.C.P. to predict how an increase in temperature
will affect the vapor pressure of a solid.
Use L.C.P. to predict how a decrease in temperature
will affect the vapor pressure of a liquid.
Use L.C.P. to predict how an increase in atmospheric
pressure will affect the vapor pressure of a liquid.
Phase Diagrams
Objective:
Interpret phase diagrams and show how a
phase diagram can be used to represent the
thermodynamic relationship between the
three states of matter for a particular
substance.
Phase Diagrams (T16)
1.
2.
3.
4.
Lines represent equilibrium between phases
Triple point
» T and P at which all three phases present @ equilibrium
Critical Temperature
» Temp at which liquid phase can not be distinguished
from its vapor
Supercritical Fluid
» Fluid at a Temp > Critical Temp.
– e.g. Supercritical carbon dioxide
Used to decaffeinate coffee and tea
Application Questions
1.
2.
3.
What phase will occur if water at –20. C and 2.15 torr
is heated to 50. oC under constant pressure?
What phase will water be in if it is at a pressure of 330
o
torr and a temperature of 50 C?
Use L.C.P. to explain why the MP of ice decreases as
pressure increases.
o
Application Questions (T17)
1.
2.
At what temperature does Dry ice sublime?
What effect does an increase in pressure have on
the melting point of carbon dioxide. Use L.C.P. to
explain why.
Crystalline Solids
Objective:
•
•
Describe how atoms, molecules, or ions are arranged
in crystalline solids
Unit Cells to Know
– Simple cubic (T18)
– Face-centered cubic (T19, 20, 21)
– Body-centered cubic (T21)
X-Ray Diffraction
Objective:
Describe the use of X-Ray diffraction to
determine the structure of crystals (T23a)
Physical Properties and Crystal Types
Objective: Relate the properties of solids to crystal type
» Understand Table 12.5, (T24a)
Crystal Types
• Ionic
• Molecular
• covalent (Network)
• Metallic (T24c)
Application Questions
1.
2.
Boron nitride, which has the empirical
formula BN,
o
melts under pressure at 3000 C and is as hard as
diamond. What is the probable crystal type of BN?
Crystals of elemental sulfur are easily crushed and
o
melt at 113 C to give a clear yellow liquid that does
not conduct electricity. What is the probable crystal
type for solid sulfur?
Physical Properties of Graphite vs. Diamond
Property
Graphite
Diamond
2.27
Very soft
3.51
Very hard
Shiny black
Colorless/transparent
Electrical
Conductivity
High
None
DHcomb (kJ/mol)
-393.5
-395.4
Density (g/mL)
Hardness
Color
Which Structure is Diamond? Graphite?
Uses of Unit Cells
•
Unit cells may be used to determine the
a. Empirical formula of an ionic compound
E.g. Problems 12.90 and 12.91, page 480 Silberberg 3rd ed.
b. Molar Mass of a substance
E.g. Problem 12.93, page 480 Silberberg 3rd ed.
c. Density of a substance
(divide the mass of the atoms per unit cell by the volume
of the unit cell)