intern_forc_biceps2.ppt

1
Chapter 10
Liquids and Solids

3
Van Der Waals Forces
• These are
intermolecular
forces of attraction
between neutral
molecules.
• The Nobel Prize in
Physics 1910
(Johannes van der
Waals)
• "for his work on the
equation of state for
gases and liquids"

4
intER vs. intRA molecular forces
• Intramolecular forces are the forces
within a molecule or ionic compound
Example: Individual therapy
NaCl Ionic bond between atom of Na and atom of Cl
• Intermolecular forces are the forces
between molecules or ions and molecules
Example: couples therapy
Solid liquid gas

5
Intramolecular forces Intermolecular forces

Strength
Intramolecular bonds > intermolecular forces
Intramolecular bonds are stronger because it
would take a lot more energy to overcome
covalent bonds and break apart the molecule
than to overcome intermolecular forces in
between the atoms (to make it become a liquid or
gas).
6

Phase Changes
• When a substance changes from solid to
liquid to gas, the molecules remain intact.
• The changes in state are due to changes in
the forces among molecules rather than in
those within the molecules.
7

10
Intermolecular
Force
Model Basis of Attraction Energy
(kJ/mol)
Example

11
3 Types of van der Waals
Forces
• Dipole-Dipole forces
• London Dispersion forces
• Hydrogen bonding

13
DIPOLE-DIPOLE FORCES
• These are forces of attraction that occur between
polar molecules. (big difference in electron
negativity)
• These forces are effective only when polar
molecules are very close. As distance increase
strength of bond decreases.
• For molecules of approximately equal mass and size,
the strength of force of attraction increases as the
polarity increases.
• Radius have an effect on strength of dipole.

Dipole-Dipole Forces
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• Molecules with larger
dipole moments have
higher melting and boiling
points (hard to break)
than those with small
dipole moments.
• Dipole attractions are
relatively weak and tend
to be liquids or gas at
room temperature.

18
HYDROGEN BONDING
• A special type of dipole-dipole
interaction between the hydrogen atom
in a polar bond and an unshared
electron pair of an element that is very
electronegative usually a F, O, or N
atom on another molecule
• (note that all of these have very high
EN’s and small atomic radii).

Hydrogen Bonding

21
HYDROGEN BONDING
• These types of
bonds are super-
humanly strong.
• (unusually strong dipole
dipole 4X stronger that
diopole dipole)

23
WHY HYDROGEN BONDING IS
EFFECTIVE
• F, O, & N are extremely small and very
electronegative atoms.
• Hydrogen atoms are very small and have no inner
core of electrons, therefore, the positive side of the
bond dipole has the concentrated charge of the
partially exposed, nearly bare proton of the
nucleus.
• …in other words, the atoms have a large difference
in electronegativity and their nuclei can get really
close.

24
IMPORTANCE OF HYDROGEN
BONDING
• Are important biologically, in stabilizing
proteins and keeping DNA together.
• Also explains why ice is less dense than
water (see text).

26
LONDON DISPERSION FORCES
• Fritz London
• These are forces that
arise as a result of
temporary dipoles
induced in the atoms or
molecules.( it’s a
temporary accident!)
• All molecules have
some degree of LD
forces

London Dispersion Forces

29
• LD forces occur between neutral non-polar molecules.
(nobles gases and nonpolar compounds)
• Occurs in all molecules, including nonpolar ones.
• LD forces are weak
• The greater the number of electrons the greater the LD
force. (ie the greater the melting and boiling pt.).
Say This: The larger the electron cloud the more polarizable an the
greater the strength of the interaction
• LD force molecules have Low melting and boiling pts

30
See Graphic on next slide
• The motion of electrons in an atom or
molecule can create an instantaneous
dipole moment.
• EX: in a collection of He (g) the average
distribution of electrons about a nucleus is
spherical, the molecules are non-polar and
there is no attraction.

31
INSTANTANEOUS AND
INDUCED DIPOLES
Pg 454- 455 in text

32
(CONT)
• These forces tend to increase in strength with an
increase in molecular weight (The size of the
molecule generally increases with mass and the
electrons are less tightly held…allows the
electron cloud to be more easily distorted.
• These forces are stronger in linear molecules
than comparable “bunched up” molecules.

33
LD forces are generally
the WEAKEST
intermolecular forces.
Molecules with more
electrons will
experience more LD
forces

Which molecule is capable of forming stronger
intermolecular forces?
N2 H2O
Explain.
CONCEPT CHECK!

Draw two Lewis structures for the formula C2H6O
and compare the boiling points of the two
molecules.
C
H
H C
H
H
H
O H C
H
H C
H
H
H
O H
CONCEPT CHECK!

Which gas would behave more ideally at the same
conditions of P and T?
CO or N2
Why?
CONCEPT CHECK!

38
n-pentane vs neopentane
• BP = 309.4 K BP = 282.7 K
• Same atomic masses different structure

39
Generalizations Regarding
Relative Strengths of IM Forces
• If molecules have comparable molecular
weights and shapes, dispersion forces are
approximately equal. Any difference in
attractive forces is due to dipole-dipole
attractions.
• If molecules differ widely in molecular
weight, dispersion forces are the decisive
factor. The most massive molecule has
the strongest attractions.

40
Because melting points (MPs) and boiling points (BPs) of
covalent molecules increase with the strengths of the forces
holding them together, it is common to use MPs and BPs as a
way to compare the strengths of intermolecular forces.
This is shown below, with the molecular formulas, molar masses
and normal BPs of the first five straight-chain hydrocarbons.
Molecular Formula Molar Mass Normal BP (C)
CH4 16 - 161.5
C2H6 30 - 88.5
C3H8 44 - 42.1
C4H10 58 - 0.5
C5H12 72 36.1

41
Which noble gas element has the
lowest boiling point?
He
Ne
Ar
Kr
Xe

42
The chemical forces between HCl
is/are
• Dispersion
• Covalent bond
• Hydrogen bond
• Dipol-dipole
• Two of the above
All Molecules
Have
Not symmetrical
Polar

43
Consider the following list of
compounds. How many of these
have hydrogen bonding as their
principle IMF
HCl
NH3
CH3OH
H2S
CH4
PH3
Hydrogen Bonding is
between H and highley EN
atoms such as N, O, F, and H

44
Which of the following statements
are false or correct and why?
O2 is dipole dipole
HCl is hydrogen bonding
CO2 is dipole dipole
NH3 is hydrogen
FALSE Dipole dipole not
symmetric/polar
FALSE London
Dispersion
symmetric/nonpolar
TRUE H + N,O, or F
FALSE London
Dispersion
symmetric/nonpolar

45
ION-DIPOLE FORCES
• Attraction between an ion and the partial
charge on the end of a polar molecule.

46
ION-DIPOLE FORCES (CONT)
• The magnitude of attraction increases as
either the charge of the ion increases or
magnitude of the dipole moment
increases.
• Ion-dipole forces are important in solutions
of ionic substances in polar liquids (e.g.
water)
• Stronger than Hydrogen bonding

47
ION-DIPOLE FORCES AND THE
SOLUTION PROCESS

50
A.Identify the types of bonds in
1. Glucose
2. Cyclohexane
B.Glucose is soluble in water but
cyclohexane is not. Why?

51
A.Identify the types of bonds in
1. Glucose
H, LD, VanderWal, Dip-dip
2. Cyclohexane
LD only
B.Glucose is soluble in water but
cyclohexane is not. Why?
Glucose is polar and cyclohexane is
nonpolar. Polar compounds are soluble
in polar solvent and visversa.

53
Homework
• Pg 504-505
#’s : 35, 36, 37, 39 (you may need to read
10.1 for this part esp. LD portion)

55
CHARACTERISTICS OF LIQUIDS
• Surface tension
• Capillary action
• Viscosity

56
COHESIVE FORCES
• Intermolecular forces that bind like
molecules to one another (e.g. hydrogen
bonding).

• Which force dominates alongside the glass tube –
cohesive or adhesive forces?
cohesive forces
Convex Meniscus Formed by
Nonpolar Liquid Mercury

58
ADHESIVE FORCES
• Intermolecular forces that bind a
substance to a surface.

Concave Meniscus Formed by
Polar Water
• Which force dominates alongside the glass tube –
cohesive or adhesive forces?
adhesive forces

60
SURFACE TENSION
• A measure of the
inward forces that
must be overcome in
order to expand the
surface area of a
liquid and resist and
external force.
• The greater the forces
of attraction between
molecules (IMF) of
the liquid, the greater
the surface tension.

61
Surface Tension Cont.
• Surface tension of a
liquid decreases with
increasing
temperature.
• The stronger the
intermolecular forces
the stronger the
surface tension.
Water has a high surface
tension do to hydrogen
bonding.

62
CAPILLARY ACTION
• Another way surface
tension manifests.
• The rise of liquids up
very narrow tubes.
This is limited by
adhesive and
cohesive forces.

63
Formation of meniscus
• Water : adhesive
forces are greater
than cohesive forces
• Mercury: Cohesive
are greater than
adhesive forces.

64
VISCOSITY
• The resistance of a liquid to flow.
• The less “tangled” a molecule is expected to be, the less viscous it
is.
Water = less Viscosity
syrup = high Viscosity
Larger molecules stronger IM

65
Viscosity Cont.
• Viscosity decreases with increasing
temperature (molecules gain kinetic
energy and can more easily overcome
forces of attraction).
• Viscosity Increases as pressure increases.
• Liquids with strong IMF have a higher
viscosity.

66
Homework
• Pg 505
#’s 43-45 all

Ionic Solids
68
• Formed by cation and anion. Typicaly metals and non
metals.
• Crystalline solids
• High melting points and Boiling point due to strong
attractions
• Poor conductors in solid form good conductors in
solution.
• Brittle

Molecular Solids
• Neutral molecules that form molecular
lattice structures
• Low MP and BP
• Non conductors in all states
70

Covalent network solids
• Distinct atoms all bound covalently.
• High MP and BP
• Made of carbons and Si Ge and B
• Poor conductors except for graphite sp2 hybridization
and delocalized electrons
72
Diamond

Network Solids

Metallic Solids
• Metallic bonding
• Great conductors or heat and electricity
• Ductile and malleable
• MP and BP vary
74

ID The type of Solid based on
the formula
75

77
Malleability
Brittle
Nothing
flexible polymer

92
Distillation
Vaporization
AKA: steam
Condensation
Takes advantage differences in IMF and thus vapor pressures and therefore
their BP .
Higher VP = Lower BP and weaker IMF

93
Changes in state
• Liquid  Gas Vaporization Endothermic
• Gas  Liquid Condensation Exothermic

94
• Solid  Gas Sublimation Endothermic
• Gas  Solid Deposition Exothermic

95
• Solid  Liquid Melting Endothermic
• Liquid  Solid Freezing Exothermic

96
Changes of state
• The energy involved it phase changes is
calculated using
– Heat of fusion (solid  liquid or liquid solid)
– Heat of vaporization (liquid gas or gas liquid)

97
Energy Changes and Phase
Changes
Heat of Vaporization: Vaporization is an
endothermic process ( it requires heat). Energy is
required to overcome intermolecular forces to turn
liq to gas. (AKA evaporate)
Hvap is an Indicator of strength of IMF
Larger molecule…greater IMF…greater Hvap
Methane Propane
CH4 C3H8
9.2 kJ/mol 18.1 kJ/mol

98
Question
How much energy does it take to
vaporizer 111 g of water?
Given: Hvap water= 40.67 kJ/mol
111 g H2O 1 mol x 40.6kJ = 250kJ
18g 1mol

99
• Heat of Fusion: the enthalpy change
associated with melting. (Solid to liquid.)
• Hfusion water= 6.01 kJ/mol
NOTE: heat of fusion is always
smaller than heat of
vaporization. This makes sense
think about the level of “order”
in the molecules in these phases.

Which is larger for a given substance: ΔHvap or
ΔHfus?
Explain why.
CONCEPT CHECK!

Changes of State

102
Heating Curve
• A plot of the temperature versus time

103
Heat of
Vaporization
Heat of Fusion

104
Example
Calculate the enthalpy change associated
with converting 1.00 mole of ice -25ºC to
water 150ºC at 1 atm. Specific heat of ice,
water, and steam are 2.09 J/g ºC and
4.184 J/g ºC, 1.84 J/g ºC . The heat of
fusion of ice is 6.01 kJ/mol and heat of
vaporization of water is 40.67 kJ/mol

105
ICE -25 °C
q1
Ice water
0°C
q2
Water 0°C
q3
Water vapor
100°C
q4
Vapor
150°C

106
1 mol ice  1 mol ice 1 mol water  1 mol water  1 mol steam
T= -25ºC 0ºC 0ºC 100ºC 100ºC -> 150
qtotal = q1 + q2 + q3 + q4 + q5
1.)q = 2.09(18g)(-25-0)
2.) q = 6.02 KJ/mol (convert heat of fusion)
3.) q = 4.184(18g) (100-0)
4.) q = 40.7 KJ/mol (convert heat of
vaporization)
5. ) q = 1.84(18 g) (150-100)

107
Critical Stuff
• Critical Temperature: The temperature above
which it is impossible to liquefy the gas under
study no matter how high the applied pressure.
• Critical Pressure: The pressure required to
liquefy a gas as at its critical temperature
NOTE: the critical temp of a gas gives an indication
of the strength of the IMF of that gas. A
substance with weak attractive forces would have
a low critical temp.

108
Which gas can be liquefied at 25ºC
Gas Critical Temp
ºC
Critical
Pressure
atm
Ammonia 132 112
Ethanol 158 78
Argon -186 6
Critical Temp above
25ºC
Critical Temp under
25ºC

109
Vapor Pressure (vp)
Vapor Pressure: Pressure
exerted by molecules that have
enough energy to escape the
surface.
As T ↑ VP ↑evaporation ↑
Liquids with high VP are volatile
(alcohol evaporates easily)
Liquids that have strong IMF have
low vapor pressures.
(take a lot of energy to overcome
IMF so it can evaporate)

Kinetic energy
%
o
f
M
o
l
e
c
u
l
e
s
T2
• At higher temperature more
molecules have enough energy
• Higher vapor pressure.

111
• Liquids with high VP
are volatile (alcohol
evaporates easily)
• Liquids that have
strong IMF have low
vapor pressures.
• (take a lot of energy to
overcome IMF so it can
evaporate)
substance vapor
pressure at
25oC
diethyl ether
C4H10O
0.7 atm
Bromine
Br2
0.3 atm
ethyl alcohol
C2H5OH
0.08 atm
Water
H2O
0.03 atm

Evaporation
• Molecules at the surface break away
and become gas.
• Only those with enough KE
escape
• Evaporation is a cooling
process.
• It requires heat.
• Endothermic.

Condensation
Change from gas to liquid
Achieves a dynamic equilibrium with
vaporization in a closed system.
What the heck is a
“dynamic equilibrium?”

Liquid/Vapor Equilibrium

When first sealed the
molecules gradually escape
the surface of the liquid
As the molecules build up
above the liquid some
condense back to a liquid.
Dynamic equilibrium

As time goes by the rate of
vaporization remains constant
but the rate of condensation
increases because there are
more molecules to condense.
Equilibrium is reached when
Rate of Vaporization = Rate of Condensation
Dynamic equilibrium

117
VP example
In a closed container the number of partials changing
from liquid  vapor will eventually equal the number
changing from vapor  liquid.

118
Boiling Point
Note: The normal boiling point of water is 100oC. The term
normal refers to standard pressure or 1 atm, or also 101.3
kPa.
The vapor pressure of the liquid = air pressure above the liquid

119
Boiling Pts. of H2O at Various
Elevations
Altitude compared
to Sea Level
(m)
Boiling
Point
(°C)
1609 98.3
177 100.3

120
How to make something boil
1. Increase the VP of the liquid (heat it) so
that the VP of the liquid is > that of the
atmosphere.
2. Lower the atmospheric pressure
(pressure above the liquid) (put a lid on
it)

121
Boiling Point
↑ boiling pt by
↑ in IMF
Or
↓ VP
At high altitudes (low air pressure) water
boils at a lower temperature

122
Normal Boiling Point
• Temperature at which something boils
when the vp =1 atm
• Note the lower the external pressure the
lower the boiling point.

123
Freezing point/melting point
• They are the same but in opposite directions.
• When heated the particles vibrate more rapidly
until they shake themselves free of each other.
• Ionic solids have strong intermolecular forces so
a high mp.
• Covalent/molecular solids have weak
intermolecular forces so a low mp.

124
Phase Diagram
• A graphical way to summarize the
conditions under which equilibrium exists
between different states of matter.
• Allows you to predict the phase of a
substance that is stable at a given
temperature and pressure

125
Triple point = three phase are in equilibrium with
each other at the same time
Boiling Point Melting Point
1 atm
Critical point

126
Critical Point: The temp beyond which the ,molecules
of a substance have to much kinetic energy to stick
together to form a liquid.

127
Phase diagrams of substances other than water the slope of
the solid liquid line slopes forward. (positive)
In water the slope of the solid-liquid lines slopes downward.
(negative)
Water
Not Water

128
Homework
• Pg 508
• #’s : 85, 87,89, 91,

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