Chapter 10

Liquids and Solids

1
 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"
3
 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

4
Intramolecular forces Intermolecular forces

5
 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
Intermolecular Model Basis of Attraction Energy Example
Force (kJ/mol)

10
 3 Types of van der Waals
Forces
• Dipole-Dipole forces

• London Dispersion forces

• Hydrogen bonding

11
12
 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.

13
 Dipole-Dipole Forces

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DIPOLE-DIPOLE FORCES

_
+
+

_ +

_ +

15
FYI

16
 DIPOLE-DIPOLE FORCES
• 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.

17
 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).

18
 Hydrogen Bonding

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20
 HYDROGEN BONDING
• These types of
bonds are super-
humanly strong.

• (unusually strong dipole

dipole 4X stronger that
diopole dipole)

21
HYDROGEN BONDING

22
 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. 23
 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).

24
25
 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

26
London Dispersion Forces

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 LONDON DISPERSION FORCES
• 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

29
 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.
30
 INSTANTANEOUS AND
INDUCED DIPOLES

Pg 454- 455 in text

31
 LONDON DISPERSION FORCES
(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.

32
 LONDON DISPERSION FORCES

LD forces are generally

the WEAKEST
intermolecular forces.

Molecules with more

electrons will
experience more LD
forces
33
CONCEPT CHECK!

Which molecule is capable of forming stronger

intermolecular forces?

N2 H2O

Explain.

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CONCEPT CHECK!

Draw two Lewis structures for the formula C2H6O

and compare the boiling points of the two
molecules.

H H H H
H C C O H H C O C H
H H H H

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CONCEPT CHECK!

Which gas would behave more ideally at the same

conditions of P and T?

CO or N2

Why?

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37
 n-pentane vs neopentane

• BP = 309.4 K BP = 282.7 K

• Same atomic masses different structure

38
 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.
39
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
C 2H 6 30 - 88.5
C 3H 8 44 - 42.1
C4H10 58 - 0.5
C5H12 72 36.1

40
 Which noble gas element has the
lowest boiling point?
He

Ne

Ar

Kr

Xe

41
 The chemical forces between HCl
is/are
• Dispersion All Molecules
Have

• Covalent bond
• Hydrogen bond
• Dipol-dipole Not symmetrical
Polar
• Two of the above

42
 Consider the following list of
compounds. How many of these
have hydrogen bonding as their
principle IMF
Hydrogen Bonding is
between H and highley EN
atoms such as N, O, F, and H
HCl
NH3
CH3OH
H2S
CH4
PH3
43
 Which of the following statements
are false or correct and why?
O2 is dipole dipole FALSE London
Dispersion
symmetric/nonpolar

HCl is hydrogen bonding FALSE Dipole dipole not

symmetric/polar

FALSE London
CO2 is dipole dipole Dispersion
symmetric/nonpolar

NH3 is hydrogen TRUE H + N,O, or F

44
 ION-DIPOLE FORCES
• Attraction between an ion and the partial
charge on the end of a polar molecule.

45
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

46
ION-DIPOLE FORCES AND THE
SOLUTION PROCESS

47
48
ANSWER

49
A.Identify the types of bonds in
1. Glucose
2. Cyclohexane

B.Glucose is soluble in water but

cyclohexane is not. Why?

50
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.

51
52
 Homework
• Pg 504-505

#’s : 35, 36, 37, 39 (you may need to read

10.1 for this part esp. LD portion)

53
CHARACTERISTICS OF LIQUIDS
• Surface tension

• Capillary action

• Viscosity

55
 COHESIVE FORCES
• Intermolecular forces that bind like
molecules to one another (e.g. hydrogen
bonding).

56
 Convex Meniscus Formed by
Nonpolar Liquid Mercury
• Which force dominates alongside the glass tube –
cohesive or adhesive forces?

cohesive forces

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

58
 Concave Meniscus Formed by
Polar Water
• Which force dominates alongside the glass tube –
cohesive or adhesive forces?

adhesive forces

Copyright © Cengage Learning. All rights reserved 59

 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.
60
 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. 61
 CAPILLARY ACTION
• Another way surface
tension manifests.

• The rise of liquids up

very narrow tubes.
This is limited by
adhesive and
cohesive forces.

62
 Formation of meniscus
• Water : adhesive
forces are greater
than cohesive forces

• Mercury: Cohesive
are greater than
adhesive forces.

63
 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

64
Larger molecules stronger IM
 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.
65
 Homework
• Pg 505
#’s 43-45 all

66
Four types of solids

67
 Ionic Solids
• 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

68
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 Molecular Solids

• Neutral molecules that form molecular

lattice structures
• Low MP and BP
• Non conductors in all states
70
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 Covalent network solids

Diamond

• 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
 Network Solids

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 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
76
 Malleability

Brittle

Nothing
flexible polymer

77
 Distillation
Takes advantage differences in IMF and thus vapor pressures and therefore
their BP .
Higher VP = Lower BP and weaker IMF

Condensation

Vaporization
AKA: steam

92
 Changes in state
• Liquid  Gas Vaporization Endothermic

• Gas  Liquid Condensation Exothermic

93
• Solid  Gas Sublimation Endothermic

• Gas  Solid Deposition Exothermic

94
• Solid  Liquid Melting Endothermic

• Liquid  Solid Freezing Exothermic

95
 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)

96
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
Methane Propane

CH4 C3H8

9.2 kJ/mol 18.1 kJ/mol

Larger molecule…greater IMF…greater Hvap 97

 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

98
• 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.
99
CONCEPT CHECK!

Which is larger for a given substance: ΔHvap or

ΔHfus?

Explain why.

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 Changes of State

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 Heating Curve
• A plot of the temperature versus time

102
 Heat of
Vaporization

Heat of Fusion

103
 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

104
 Water vapor
100°C Vapor
q4 150°C

Water 0°C
Ice water
0°C q3
q2

ICE -25 °C
q1

105
 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) 106
 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.
107
Which gas can be liquefied at 25ºC
Gas Critical Temp Critical
ºC Pressure
atm
Ammonia 132 112
Critical Temp above
Ethanol 25ºC
158 78

Argon -186 6 under

Critical Temp
25ºC
108
 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)

109
%
• At higher temperature more molecules
have enough energy
o • Higher vapor pressure.
f

T2
M
o
l
e
c
u
l Kinetic energy
e
• Liquids with high VP substance vapor
are volatile (alcohol pressure at
evaporates easily) 25oC
diethyl ether 0.7 atm
• Liquids that have C4H10O
strong IMF have low
Bromine 0.3 atm
vapor pressures.
Br2

• (take a lot of energy to

ethyl alcohol 0.08 atm
overcome IMF so it can
C2H5OH
evaporate) Water 0.03 atm
H 2O
111
 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

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 Dynamic 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
 VP example

In a closed container the number of partials changing

from liquid  vapor will eventually equal the number
changing from vapor  liquid.
117
 Boiling Point
The vapor pressure of the liquid = air pressure above the liquid

Note: The normal boiling point of water is 100oC. The term

normal refers to standard pressure or 1 atm, or also 101.3 118
kPa.
 Boiling Pts. of H2O at Various
Elevations
Altitude compared Boiling
to Sea Level Point
(m) (°C)

1609 98.3

177 100.3

119
 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)

120
 Boiling Point
↑ boiling pt by
↑ in IMF
Or
↓ VP

At high altitudes (low air pressure) water

boils at a lower temperature

121
 Normal Boiling Point
• Temperature at which something boils
when the vp =1 atm

• Note the lower the external pressure the

lower the boiling point.

122
 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.
123
 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

124
 Triple point = three phase are in equilibrium with
each other at the same time

1 atm

Boiling Point Melting Point

Critical point
125
 Critical Point: The temp beyond which
the ,molecules of a substance have to much kinetic
energy to stick together to form a liquid.
126
 Water
Not Water

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) 127
 Homework
• Pg 508
• #’s : 85, 87,89, 91,

128