Transcript Intermolecular Forces (a) Particles in solid (b) Particles in liquid (c) Particles in gas.

Intermolecular Forces
(a) Particles in solid
(b) Particles in liquid
(c) Particles in gas
Properties of Solids, Liquids, and Gases
Property
Solid
Liquid
Gas
Shape
Has definite shape
Takes the shape of
the container
Volume
Has a definite volume
High densities
Has a definite volume Fills the volume of
High densities
the container
Low densities
Bonding
Ionic, Metallic, Covalent
Covalent
Arrangement of
Particles
Fixed, very close
Random, close
Crystalline or amorphous
Interactions Between
Particles
Very strong forces:
(i.e. Melting point,
malleability, ductility,
conductivity…)
Strong forces:
(i.e. Boiling point,
Surface Tension,
Viscosity, Vapor
pressure…)
Takes the shape
of its container
Covalent
Random, far apart,
Collisions
Essentially none
Surface Tension
• molecules minimize
their surface area (“skin”)
Liquid
• molecules at surface
interact only with molecules
in the interior of liquid
• surface molecules
subjected to inward force,
so surface is under tension
• surface tension increases
with increasing intermolecular
forces
H2O(l) Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31
Viscosity
• resistance of a liquid to flow
Liquid
• greatest in substances with
strong intermolecular forces,
which hinder flow
• longer molecules higher
viscosity than shorter ones
H2O(l) Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31
Acetone
Gasoline (Hexane)
Corn Syrup
Motor Oil
Molasses
Intermolecular Forces
Dispersion (London Force)
• (a) Interaction of any two
atoms or molecules.
Electrons unevenly
distributed. Creates
instantaneous (temporary
dipole). Polarization
increases with size.
-
+
-
-
-
-
+
-
+
-
+
Attraction
Repulsion
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 442
+
+
-
• (b) interaction of many
dipoles. WEAKEST forces!
+
-
+
-
+
+
Intermolecular Forces
Dipole-Dipole
- +
• (a) Interaction of two
polar molecules. Polar
molecules have permanent
dipoles from electronegativity
difference. Higher melting and
boiling points due to
+
stronger IM forces.
• (b) interaction of many
dipoles in a liquid.
-
-
+
+
-
+
-
-
+
Attraction
Repulsion
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 442
+
-
+
-
+
+
Intermolecular Forces
Hydrogen Bonding
• Strong intermolecular
forces of attraction
between molecules
containing fluorine,
oxygen, or nitrogen
bonded to hydrogen
• Results from large
electronegativity
difference and small
atomic size of hydrogen
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 442
Hydrogen
bonds
H
Chemical
Bonds
O
H
O
H
H
Chemical
Bonds
Solubility
Dissolving of NaCl in Water
Na+
ions
Water molecules
Clions
NaCl(s) + H2O  Na+(aq) + Cl-(aq)
Ethanol is Polar
dH
d+
O
H
Polar bond
H
C
H
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 469
C
H
H
Ethanol and Water are Soluble
H
H
H
H
C
O
O
H
C H
‘Like dissolves like’
H
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 469
H
“Oil and Water Don’t Mix”
• Oil is nonpolar
• Water is polar
“Like dissolves like”,
nonpolar dissolves
nonpolar, nonpolar
does not dissolve
polar
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 470
Interstitial Spaces
Oil
Oil
Oil
Oil
Oil
Oil
Oil
Non-polar
"immiscible"
Layer
dissolved
solid
Water
Water
Water
Water
Water
Water
Water
Water
Polar
red food
coloring
Cleaning Action
of Soap
Micelle
Timberlake, Chemistry 7th Edition, page 573
Evaporation (Vaporization)
• Molecules must have sufficient energy to break
IM bonds.
• Molecules at the surface break away and
become gas (“volatility”).
• Only molecules with enough KE escape.
• Breaking IM bonds absorbs energy.
Evaporation is endothermic.
• Rate of evaporation increases with
increasing surface area, increasing
temperature, and weaker IM forces
Condensation
• Forming IM bonds from gas to liquid
• Condensation is exothermic because energy
is released.
• Dynamic equilibrium: rate of vaporization
equals rate of condensation (gas molecules
above liquid becomes constant).
• Vapor pressure: partial pressure
of gas in dynamic equilibrium with liquid
• Vapor pressure increases
with increasing temperature
and weaker IM forces
Boiling point: temperature
at which the vapor pressure
of a liquid is equal to the
pressure above it
Microscopic view of a liquid near its
surface
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 446
Formation of a bubble is opposed by
the pressure of the atmosphere
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 452
Effect of Pressure on Boiling Point
Boiling Point of Water at Various Locations
Feet
Boiling
Location
above Patm (kPa)
Point
sea level
(C)
Top of Mt. Everest, Tibet
29,028
32.0
70
Top of Mt. Denali, Alaska
20,320
45.3
79
Top of Mt. Whitney, California
14,494
57.3
85
Leadville, Colorado
10,150
68.0
89
Top of Mt. Washington, N.H.
6,293
78.6
93
Boulder, Colorado
5,430
81.3
94
Madison, Wisconsin
900
97.3
99
New York City, New York
10
101.3
100
-282
102.6
100.3
Death Valley, California
Boiling Points of Covalent Hydrides
H2O
100
Boiling point (oC)
0
H2Te
H2Se
SnH4
H2S
-100
GeH4
SiH4
-200
CH4
50
Molecular mass
100
150
Heating Curves
• Temperature Change
– change in KE (molecular motion)
– depends on heat capacity
• Phase Change
– change in PE (molecular arrangement)
– temperature remains constant
• Heat Capacity
– energy required to raise the temp of 1 gram of a
substance by 1°C (q = mC∆T)
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Heating Curve for Water
Temperature (oC)
vaporization
D
100
condensation
melting
C
liquid
B
0
A
solid
freezing
Heat added
LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487
E
gas
Heating Curves
140
120
Gas - KE 
Temperature (oC)
100
Boiling - PE 
80
60
40
20
0
-20
Liquid - KE 
Melting - PE 
-40
-60
-80
Solid - KE 
-100
Energy
Heating Curves
• Heat of Fusion (Hfus)
– energy required to melt 1 gram of a substance at its
melting point. Breaking intermolecular forces in the solid.
Water = 6.02 kJ/mol (@ M.P.)
• Heat of Vaporization (Hvap)
– energy required to vaporize 1 gram of a substance at its
boiling point.
– usually larger than Hfus (requires complete separation of
molecules)
– higher temperatures = lower (Hvap)
Water = 40.7 kJ/mol (@ B.P.)
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Calculating Energy Changes Heating Curve for Water
H = mol x Hvap
140
120
H = mol x Hfus
Temperature (oC)
100
80
∆H = mCgasT
60
40
20
0
∆H = mCliquidT
-20
-40
-60
-80
∆H = mCsolidT
-100
Energy
Specific Heat Capacities
Substance
Specific Heat
J/g°C
Aluminum
0.903
Carbon (dia)
0.508
Carbon (gra)
0.708
Copper
0.385
Gold
0.128
Iron
0.449
Lead
0.128
Silver
0.235
Ethanol
2.42
Water (l)
4.184
Water (s)
2.03
Water (g)
2.02
Tro's "Introductory Chemistry",
Chapter 3
States of Matter
Energy Changes Accompanying Phase Changes
Gas
Energy of system
Vaporization
Condensation
Sublimation
Liquid
Melting
Freezing
Solid
Brown, LeMay, Bursten, Chemistry 2000, page 405
Deposition
Phase Diagrams
• Show the phases of a substance at different
temps and pressures.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Intermolecular Forces
• Determine the kinds of IM forces that are
present in the following: Kr, N2, CO, HF, NCl3,
NH3, SiH4, CCl4
• Which of the following has the highest boiling
point. Why? CH4, CH3CH3, CH3CH2CH3,
CH3CH2CH2CH3
• A flask containing a mixture of NH3(g) and
CH4(g) is cooled. At -33.30C a liquid begins to
form in the flask. What is the liquid?
Melting Points
1
H
Mg
-259.2
2
3
4
Li
Be
180.5
1283
650
K
Ca Sc
Rb Sr
38.8
6
> 3000
98
850
770
710
B
oC
2000 - 3000
oC
Al
660
Ti
V
C
N
O
F
Y
1500 1852 2487 2610 2127 2427 1966 1550
920
Ta
P
1423 44.2
420 29.78 960
Zr Nb Mo Tc Ru Rh Pd Ag Cd
Hf
Si
S
119
Ne
W
Re Os
Ir
961
In
Ar
-101 -189.6
Kr
817 217.4 -7.2 -157.2
Sn Sb Te
I
Xe
321 156.2 231.9 630.5 450 113.6 -111.9
Pt Au Hg
Tl
Pb Bi
Po At Rn
2222 2997 3380 3180 2727 2454 1769 1063 -38.9 303.6 327.4 271.3 254
Ralph A. Burns, Fundamentals of Chemistry , 1999, page 1999
Cl
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
1423 1677 1917 1900 1244 1539 1495 1455 1083
Cs Ba La
28.6
-269.7
2027 4100 -210.1 -218.8 -219.6 -248.6
Na Mg
63.2
5
650
He
Symbol
Melting point oC
-71
Densities of
Elements
1
2
3
4
H
He
0.071
0.126
Li
Be
B
C
N
O
0.53
1.8
2.5
2.26
0.81
1.14
Na Mg
Al
Si
P
S
0.97
2.70
2.4 1.82w 2.07 1.557 1.402
K
0.86
5
Ca Sc
Ti
V
1.55
4.5
5.96
Rb Sr
(2.5)
2.6
I
Xe
4.93
3.06
7.86
8.9
8.90
8.92
7.14
5.91
5.36
5,7
4.7
6.4
8.4
10.2
8.6
7.3
7.3
Cs Ba La
Hf
Ta
W
Pt Au Hg
Tl
Pb Bi
1.90
13.1
16.6
19.3
8.0 – 11.9 g/cm3
Mg
1.74
W
Ar
3.119
7.4
5.51
6.7
Cl
7.1
Sn Sb Te
3.5
1.11 1.204
Kr
In
2.6
Y
Ne
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
Zr Nb Mo Tc Ru Rh Pd Ag Cd
1.53
6
1.74
F
11.5
12.5
Re Os
12.5
Ir
12.0
10.5
21.4 22.48 22.4 21.45 19.3 13.55 11.85 11.34
12.0 – 17.9 g/cm3
6.7
9.8
6.1
Po At Rn
9.4
> 18.0 g/cm3
Symbol
Density in g/cm3C, for gases, in g/L
---
4.4
Molecular Structure of Ice
Hydrogen
bonding
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 455
Water boiling on a stove
Hot vs. Cold Tea
Many molecules have an
intermediate kinetic energy
Low temperature
(iced tea)
High temperature
(hot tea)
Percent of molecules
Few molecules have a
very high kinetic energy
Kinetic energy