You Ever Ask???
• Why is water usually a liquid and
not a gas?
• Why does liquid water boil at
such a high temperature for
such a small molecule?
• Why does ice float on water?
• Why do snowflakes have 6
sides?
• Why is I2 a solid whereas Cl2 is
a gas?
• Why are NaCl crystals little
cubes?
1
Chapter 11: Liquids, Solids and
Intermolecular Forces
2
Chapter
Topics
Chapter
Objectives
• Kinetic-Molecular Description of Liquids & Solids
• Intermolecular (IMF) Attractions & Phase
Changes
• Properties of Liquids
– Viscosity, Surface Tension, Capillary Action
– Evaporation, Vapor Pressure,
– Boiling Points & Distillation, Heat transfer
• Phase Changes
The following will NOT be covered in class
Types of crystals
3
States of Matter
• List all the differences between
– Solids
– Liquids
– Gases
• Kinetic Energy?
4
States of Matter (KE of Matter)
The fundamental difference between states of
matter is the distance between particles.
Intermolecular attractions in liquids & solids are strong:
KE of molecules << IMF
5
States of Matter
Because in the solid and liquid states
particles are closer together, we refer to them
as condensed phases (depends on T and P).
6
The States of Matter
• The state a substance is in
at a particular temperature
and pressure depends on
two antagonistic entities:
– The kinetic energy of
the particles
– The strength of the
attractions between the
particles
7
Examples
1. What is the difference between
intermolecular forces and intramolecular
forces?
2. List all the intermolecular forces you are
familiar with.
3. List all the intramolecular forces you are
familiar with.
4. Which forces are stronger?
8
Intermolecular Forces
The attractions between molecules are not nearly as strong
as the intramolecular attractions that hold compounds
together.
Intermolecular forces are weaker than intramolecular forces ( to
break 2 O-H in bonds in water: 930 kJ/mol; to vaporize water:
43 kJ/mol)
Responsible for the existence of condensed states (liquids, solids)
9
Intermolecular Forces
They are, however, strong enough to control physical
properties such as boiling and melting points, vapor
pressures, surface tension, and viscosities (reflect the
strength of the bond).
INTRAmolecular forces—the forces that holds atoms
together o form molecules
INTERmolecular forces: the forces between molecules, ions
and molecules-ions.
10
Intermolecular Forces
These intermolecular forces as a group are
referred to as van der Waals forces (electrostatic
forces).
11
IMF Problem
For each of the molecules below,
Determine the geometry of the molecule
Determine the polarity of the molecule
List the types of intermolecular force which act
between pairs of these molecules.
(a) CH4
(b) PF3
(c) CO2
(d) HCN,
(e) HCOOH (methanoic acid)
12
Types of IMF
• Ion – Ion
• Van der Waals Forces
Dipole – dipole (for molecules with dipole moments)
Dipole – induced dipole
Dispersion forces (London)
Hydrogen Bond (special case of dipole-dipole (IMF)
London Dispersion Forces (induced dipole-dipole)
• Ion – induced dipole
• Ion – dipole
• Total attraction between molecules may depend on
more than one type of intermolecular force.
13
Ion – Ion Forces
08M16VD1
14
Ion-Ion Forces
• Ion-ion forces: electrostatic forces of attraction
between __________________ of ionic
compounds. Generally very strong → 250 kJ.
(not a true intermolecular force)
• Ionic compounds: metal and nonmetal or
polyatomic anions (NH4+)
• Coulomb’s law & the attraction energy
determine:
– melting & boiling points of ionic compounds
– the solubility of ionic compounds
15
IMF: Ionic Solids
• Ion-ion interactions
– force of attraction between two oppositely charged
ions is determined by Coulomb’s law

F
q+
 
q-
d2
• Energy of attraction between two ions is given by:
q q   d
Fd =
+
E =
q q 
=
+
-
d2
-
d
D:\Media\Movies\08M17AN1.MOV
17
Ion-Ion Forces
for comparison of magnitude
Na+—Cl- in salt
These are the strongest
forces.
Lead to solids with high
melting temperatures.
NaCl (lattice energy = 788
kJ/mol), mp = 800 oC
MgO (lattice energy = 3890
kJ/mol) , mp = 2800 oC
18
Covalent Bonding Forces
for comparison of magnitude
C=C, 610 kJ/mol
C–C, 346 kJ/mol
C–H, 413 kJ/mol
CN, 887 kJ/mol
19
Ion – Dipole Force
20
Ion-Dipole Interactions
• Ion-dipole interactions are an important force in
solutions of ions.
• The strength of these forces are what make it
possible for ionic substances to dissolve in polar
solvents.
Na+(g) + 6H2O(l)→ [Na(H2O)6]+(aq)
Hydrated Ions?
ΔHrxn = -405 kJ
Coordination Number?
21
Attraction Between Ions and Permanent
Dipoles
••
••
water
-
dipole
O
H
H +
22
Attraction Between Ions and
Permanent Dipoles
Attraction between ions and dipole depends on ion
charge and ion-dipole distance.
Measured by ∆H for Mn+ + H2O → [M(H2O)x]n+
- H
O
H
+
•••
Mg2+
-1922 kJ/mol
- H
O
H
+
- H
O
H
+
•••
Na +
-405 kJ/mol
•••
Cs+
-263 kJ/mol
23
Dipole – Dipole Force
D:\Media\Movies\13M04AN2.MOV
24
Dipole-Dipole Forces
Molecules that have permanent dipoles are attracted to
each other.
The positive end of one is attracted to the negative end
of the other and vice-versa.
These forces are only important when the molecules
are close to each other.
Note the difference between solid and liquid.
Liquid
liquid
Solid
25
Effect of Dipole Moment on BP
Substance
MM
BP, K
44
Dipole
moment
0.1
C3H8
46
1.3
248
44
2.7
294
41
3.9
355
231
propane
CH3OCH3
dimethyl ether
CH3CHO
Acetyl aldehyde
CH3CN
acetonitrile
26
Dipole-Dipole Interactions
The more polar the molecule (higher μ), the
higher is its boiling point.
• Basic attraction : electrostatic, Coulomb’s Law
Examples: HCl, CO, SO2, NF3, etc
27
IMF: Dipole-Dipole
• Dipole-dipole are of the order of 5 to 20
kJ/mol. (KE due to temp at 25oC about 4
kJ/mol). Cmpds that have these forces
(dipole-dipole) are frequently solids and
liquids at room temp.
• The stronger the forces, the ______ the
melting and boiling points of the
compounds.
28
Hydrogen Bond: type of dipole-dipole
force
13M07AN2
29
Boiling Points of Simple Hydrogen-Containing
Compounds
The nonpolar
series (SnH4 to
CH4) follow the
expected trend.
The polar series
follows the trend
from H2Te
through H2S, but
water is quite an
anomaly.
EXPLAIN!
30
Intermolecular Forces: H-bond
Which of these are capable of forming
hydrogen bonds among themselves?
a)
b)
c)
d)
e)
f)
g)
CH3OH
C2H4
CH3NH2
HCN
NH4+
KF
CH3COOH
31
Hydrogen Bonding
• The dipole-dipole
interactions experienced
when H is bonded to N, O,
or F (HIGH
ELECTRONEGATIVITY)
are unusually strong.
• Hydrogen nucleus is
exposed.
• We call these interactions
hydrogen bonds.
32
Hydrogen Bonding
33
H-Bonding Between Methanol and Water
-
H-bond
+
-
34
Hydrogen Bonding in H2O
H-bonding is especially
strong (40 kJ/mol) in water
because
• the O—H bond is very
polar
• there are 2 lone pairs on
the O atom
Accounts for many of water’s
(and other molecules such
as DNA, proteins) unique
properties such as
anomalous high BP and
high viscosity.
35
Hydrogen Bonding in H2O
Ice has open
lattice-like
structure.
Ice density is
< liquid.
And so solid
floats on
water.
Snow flake:
www.snowcrystals.com
36
Hydrogen Bonding in H2O
Ice has open lattice-like structure.
Ice density is < liquid and so solid floats on water.
One of the VERY few
substances where
solid is LESS DENSE
than the liquid.
37
Hydrogen Bonding
H bonds leads to
abnormally high
boiling point of water.
D:\Media\Movies\13M07AN
1.MOV
See Screen 13.7
38
Boiling Point of Hydrides in ºC
Group VIA
Group VIIA Group IVA
Group VA
H2O
100
HF
20
NH3
-33
CH4
-161
H2S
-65
HCl
-85
PH3
-87
SiH4
-112
H2Se
-45
HBr
-69
AsH3
-60
GeH4 -90
H2Te
-15
HI
-35
SbH3
-25
39
Hydrogen Bonding in Biology
H-bonding is especially strong in biological
systems — such as proteins and DNA.
D:\Media\Movies\09S03AN1.MOV
DNA — helical chains of phosphate groups and
sugar molecules. Chains are helical because
of tetrahedral geometry of P, C, and O.
Chains bind to one another by specific
hydrogen bonding between pairs of
Lewis bases.
—adenine with thymine
—guanine with cytosine
40
Double helix
of DNA
Portion of a
DNA chain
41
Base-Pairing through H-Bonds
42
Induced Dipole –Induced Dipole
(London Dispersion Forces)
43
London Dispersion Forces
While the electrons in the 1s orbital of helium
would repel each other (and, therefore, tend
to stay far away from each other), it does
happen that they occasionally wind up on the
same side of the atom.
44
London Dispersion Forces
Instantaneous dipole
The helium atom becomes polar, with an excess
of electrons on the left side and a shortage on the
right side. Instantenous dipole forms (for an
instant)
45
London Dispersion Forces
Another helium nearby, then, would have a
dipole induced in it, as the electrons on the
left side of helium atom 2 repel the electrons
in the cloud on helium atom 1.
46
London Dispersion Forces
London dispersion forces, or dispersion
forces, are attractions between an
instantaneous dipole and an induced dipole.
47
London Dispersion Forces
• These forces are present in all molecules,
whether they are polar or nonpolar.
• The tendency of an electron cloud to distort in
this way is called POLARIZABILITY.
48
Forces Involving Dipole -Induced Dipole
• Process of inducing a
dipole is polarization
• Degree to which
electron cloud of an
atom or molecule can
be distorted in its
polarizability.
49
IMF: London Dispersion Forces
• Induced Dipoles: the temporary separation of positive
and negative charges in a neutral particle due to the
proximity of an ion, dipole, or another induced dipole.
On average μ = 0.
• London Dispersion Forces: attractive forces
(electrostatic in origin) that arise as a result of temporary
dipoles induced in atoms or molecules (instantaneous
dipoles). Weak: 0.1- 5 kJ/mol
• Dispersion forces allow non-polar molecules to
condense.
• Exist in all molecules!!!!! Importance depends on
the type of intermolecular forces.
50
IMF: London Dispersion Forces
• London Forces
very weak
only attractive force in nonpolar molecules
Ar atom
Cluster of Ar atoms
51
Factors Affecting London Forces
• The shape of the molecule
affects the strength of dispersion
forces: long, skinny molecules
(like n-pentane tend to have
stronger dispersion forces than
short, fat ones (like neopentane).
• This is due to the increased
surface area in n-pentane.
52
Effect of Geometry (shape) on BP of Molecules
Compound
N-butane
Isobutane
MM
58
58
N-pentane
72
2-methyl butane
72
2,2 methyl propane 72
BP, ºC
-0.45
-12.0
36.1
27.8
9.5
•Given the same molecular mass, GEOMETRY is important
53
Factors Affecting London Forces
• The strength of dispersion forces tends to
increase with increased molecular weight.
• Larger atoms have larger electron clouds, which
are easier to polarize (larger polarizability.
54
Boiling Points of Hydrocarbons
Molecule
CH4 (methane)
C2H6 (ethane)
C3H8 (propane)
C4H10 (butane)
MM
16
38
44
58
BP (oC)
- 161.5
- 88.6
- 42.1
- 0.5
C4H10
Note: linear
relationship
between BP and
MM (Polarizability
increases).
C3H8
C2H6
CH4
55
Which Have a Greater Effect:
Dipole-Dipole Interactions or Dispersion Forces?
• If two molecules are of comparable size
and shape, dipole-dipole interactions will
likely be the dominating force.
• If one molecule is much larger than
another, dispersion forces will likely
determine its physical properties.
56
Dipole – Induced Dipole
57
FORCES INVOLVING INDUCED DIPOLES
How can non-polar molecules such as O2 and I2
dissolve in water?
The water dipole INDUCES a dipole
in the O2 electric cloud.
Dipole-induced
dipole
58
Forces Involving Dipole -Induced DIPOLE
Solubility increases with mass of the gas
59
Forces Involving Dipole -induced Dipole
Consider I2
dissolving
in ethanol,
CH3CH2OH.
-
I-I
- O
R
H
+
I-I
The alcohol
temporarily
creates or
INDUCES a
dipole in I2.
+
- O
R
H
+
60
Summary of Intermolecular Forces
• Ion-dipole forces (very strong; solubility of ions in water)
• Dipole-dipole forces (larger dipole moments)
– Special dipole-dipole force: hydrogen bonds
• Induced dipoles (occur in all substances; important for
nonpolar molecules); increase with molar mass (glues);
depend on geometry. For large molecules may exceed
the force of dipole-dipole force (polymers, glues)
In general: ionic forces the strongest; then
hydrogen bonding; dipole-dipole; and lastly
dispersion for species of similar molar mass.
London Dispersion Forces exist in all molecules
and ions.
61
Summarizing Intermolecular Forces
62
Summary of dipole forces
63
Intermolecular Forces Summary
64
Intermolecular Forces
Figure 13.13
65
Example
What type of intermolecular forces exist between
the following pairs?
1. HBr and H2S
2. Cl2 and CBr4
3. I2 and NO34. NH3 and C6H6
1. Dipole-dipole; dispersion;
2. dispersion;
3. ion-induced dipole; dispersion
4. dipole-induced dipole; dispersion
66
Intermolecular Forces: Examples
5. Order the following compounds in order of increasing MP and BP:
a) N2, O2, H2
b) Cl2, F2, I2, Br2
c) SiH4, GeH4, SnH4, CH4
Order of MP and BP:
(a) H2< N2<O2
(b) F2<Cl2<Br2<I2
(c) CH4<SiH4<GeH4<SnH4
•
Always compare like species.
6.
7.
8.
But what about HF, HCl, HBr, HI
and H2O, H2S, H2Se, H2Te
and NH3, PH3, AsH3 and SbH3.
67
Examples
1. List the IMF and arrange the substances
BaCl2, H2, CO, and Ne in order of
increasing boiling points.
2. In which of the following substances is
hydrogen bonding possible?
a.
b.
c.
d.
e.
f.
Methane
Methyl alcohol
Hydrazine (H2NNH2)
Methyl fluoride
Hydrogen sulfide
Carboxylic acid
68
Intermolecular Forces: Examples
4. Arrange in order of increasing BP:
CO2 CH3OH
CH3Br , RbF
5. Which one in each pair has the higher BP?
a)
b)
c)
d)
e)
CH4 and C2H6
H2S and H2Te
NH3 and PH3
HCl and HF
I2 and ICl
69
Example (London Forces)
6. Arrange the following molecules in order of
increasing strength of intermolecular forces.
F2
Br2
Cl2
I2
7. Explain the trend in the normal boiling points of
these liquids in terms of intermolecular forces.
CH4 normal boiling point: -161.5°C
CF4 normal boiling point: -28°C
CCl4 normal boiling point: +77°C
CBr4 normal boiling point: +190°C
70
Intermolecular Forces: Examples
8. (11.100) Which of the following
substances has the highest
polarizability?
CH4, H2, CCl4, SF6, H2S
71
Identify “intermolecular” forces
9. H2O
10. CH2Cl2
11. KBr
12. F- + H2O
13. I2
14. CH3OH
15. PCl3
16. C6H6
17. Fe
18. CS2
19. BCl3
20. Na+ + NH3
21. Dimethyl ether (CH3OCH3) and ethanol (C2H5OH)
have the same formula (C2H6O) but the BP of the ether is
-25ºC and of the ethanol 78ºC. Explain.
72