Chapter 11
Liquids and Solids
Kinetic-Molecular Description
Kineticof Liquids and Solids

Solids and liquids are condensed states.
◦ The atoms, ions, or molecules in solids and liquids
are much closer to one another than in gases.
◦ Solids and liquids are highly incompressible.

Liquids and gases are fluids.

The intermolecular (forces between
molecules) electrical attractions in liquids
and solids are strong
◦ They easily flow.
2
Kinetic-Molecular Description of
KineticLiquids and Solids
If we compare the strengths of interactions
among particles and the degree of ordering of
particles, we see that
Gases<< Liquids < Solids
 Miscible liquids are soluble in each other.


Examples of miscible liquids:
 Alcohol dissolves in water
 Oil dissolves in motor gasoline
3
1
Intermolecular Attractions
 Ion
– Ion
 Ion - Dipole
p
p
 Dipole
– Dipole
◦ Hydrogen Bonding
 Dispersion
Forces
Ion – Ion Intermolecular Attractions

Applies to Ionic solids like NaCl or CaBr2

Ions are arranged in arrays so that they are very
close together
Because of the small distances between ions of
opposite charge (full positive and full negative
charge), the energies of attraction are very high
Most ionic bonding in solids is strong and as a
result have very high melting points
Attraction of 2+ for a 2- is stronger than the
attraction of 1+ for 1-
◦ Complete transfer of electrons



Dipole – Dipole Intermolecular
Attractions



A dipole is a separation of charge between two
covalently bonded atoms
Dipole-dipole interactions occur between
polar covalent molecules because of the
attraction of the ppartiallyy ppositivelyy charge
g
atoms of one molecule for the partially
negatively charged atoms of a second molecule
Energy of dipole-dipole interactions is about 4
kJ per mole of bonds (fairly weak)
2
Dipole – Dipole Intermolecular
Attractions

Examples of polar molecules include CH3Cl,
H2S and (CH3)2C=O
Ion-Dipole Intermolecular
IonAttractions



Forces between a charged ion and a polar
molecule (a molecule with a dipole)
+2 ions have a stronger effect than +1 ions
Positive ions are usually smaller than
negative
ti ions,
i
and
d th
thus hhave a hi
higher
h charge
h
density
◦ Positive ions interact more strongly with dipoles
than negative ions

Hydration is of salts in water is an example
of a ion-dipole interaction
Ion--Dipole Forces
Ion
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
9
3
Hydrogen Bonding




Hydrogen Bonding is a special case of very strong
dipole-dipole interaction
Hydrogen Bonding occurs among polar covalent
(partially ionic) molecules containing H bonded
to one of the following elements: F,
F O,
ON
F, O, N are small and highly electronegative
The energy of the hydrogen bond is 15 – 20
kJ/mole of bonds (4 to 5 times larger than normal
dipole-dipole interactions)
Hydrogen Bonding
H 2O
CH3OH
NH3
Figure 13-4
General Chemistry
Whitten, Davis,
Peck
6th Ed.
The Liquid State
Compound MW
BP (°C)
HF
20
19.5
HCl
37
-85.0
HBr
81
-67.0
HI
128
-34
What kind of
intermolecular
forces are at
work here?
4
The Liquid State
Hydrogen Halides
Boiling Point (C)
40
20
HF
0
-20 0
50
100
150
-40
-60
HCl
HBr
HI
-80
-100
Molar Mass
So, what is going on with HF?
The Liquid State
Compound
MW
BP (°C)
H2O
18
100
H2S
34
-61
H2Se
81
-42
H2Te
130
-2
What kind of
intermolecular
forces are at
work here?
The Liquid State
Boiling Pointt (C)
VIA Hydrides
120
100
80
60
40
20
0
-20 0
-40
-60
-80
50
100
150
Molar Mass
5
Dispersion Forces
Intermolecular Attractions

Dispersion forces result from the
attraction of the positively charged
nucleus for the electron cloud of an atom
in a near byy molecule
◦ Also called London Forces

Can be viewed at temporary, induced
dipoles (some times called induced dipole
interactions)
Dispersion Forces
Intermolecular Attractions
Dispersion forces are present in all
molecules
 Weaker than dipole-dipole forces for
most small molecules
 Dispersion forces increase in strength for
larger molecules

◦ The size of the electron cloud increases and is
easier to polarize
◦ Forces can become more significant than
dipole-dipole forces
Summary of Intermolecular Forces
Interacting Species
Intermolecular Force Type
Ion + Ion
Ion-Ion
Ion + Polar Molecule
Ion-Dipole
Ion + Nonpolar Molecule
Ion-Dispersion (Ion-Induced
Dipole)
Polar Molecule + Polar Molecule
Dipole- Dipole
Polar molecule + Nonpolar
molecule
Dipole- Dispersion (Induced
Dipole Interaction)
Nonpolar Molecule + Nonpolar
Molecule
Dispersion- Dispersion (Induced
Dipole-Induced Dipole)
6
Summary

All molecules experience dispersion
forces
◦ For a system that only contains nonpolar
molecules, this is the only forces experienced
Polar
dipoleP l molecules
l l experience
i
di
l
dispersion forces
 Polar molecules also experience dipoledipole forces

◦ Polar molecules containing O-H, N-H, F-H
experience hydrogen bonding
Relative Strength of Intermolecular
Attractions
Strongest
Ion – Ion >>
Hydrogen
y g Bondingg >>
Dipole – Dipole >>
Dispersion Forces*
Weakest
*For small molecules and atoms
Liquid State

Intermolecular forces of attraction are great enough
that disordered clustering occurs
◦ The stronger the intermolecular forces in a liquid,
 The lower the vapor pressure
 The higher the boiling point, surface tension, & viscosity




Particles
P
i l are close
l
together
h
Difficult to compress
Particles are able to slide past one another so that
liquids, so that the material flows, and assumes the
shape of its container
Liquids slowly diffuse into other liquids in which they
are miscible (mixable)
7
Surface Tension
Molecules at the surface
are only affected by
intermolecular forces
from the interior
molecules pulling the
surface layer toward the
center
 Most stable shape is
where surface area is
minimized – droplets or
spheres

Viscosity
Viscosity is the resistance to flow
Units of viscosity centipoise (cP) at T°C
 Higher intermolecular forces of attraction,
the higher the viscosity
 Hydrogen bonding increases viscosity
 The larger the molecule, the larger the
viscosity
 Viscosity decreases with increasing
temperature


Capillary Action
A capillary is very thin tube
When a glass capillary is inserted in water,
the liquid creeps up the sides of the tube to
the point where the adhesive forces balance
the weight of the liquid
 A similar capillary action forces water to
feed from the roots of a plant to the top of
the plant (along with osmotic pressure)


8
Capillary Action
 The
forces of attraction between a
liquid & another surface are called
adhesive forces
◦ Water adheres to gglass – the ppositive
hydrogen atoms are attracted to the negative
oxygen atoms in glass
 Forces
of attraction between like
molecules are called cohesive forces
Capillary Action

If adhesive forces are stronger than the
cohesive forces, the liquid is pulled up the
capillary
◦ Liquid level rises until adhesive force balances
the weight of the liquid
◦ Water in a tube forms a meniscus (concave
shape) due to this effect

If the cohesive forces are stronger than
the adhesive forces, the liquid is pulled
down the capillary
◦ Mercury is an example – convex meniscus
Capillary Action
Adhesion >> Cohesion
Cohesion >> Adhesion
Water
Mercury
9
Water

Many properties are due to the high
degree of hydrogen bonding that occurs
in water
◦
◦
◦
◦
y g bonds per
p molecule
2 hydrogen
Solid is less dense of than liquid
Relatively high specific heat
Dissolves polar molecules and many ionic
solids
Types of Solids

Molecular Substances

Network Covalent Substances
◦ Held together by dispersion or dipole forces
◦ Held together by a network of covalent
bonds

Ionic Substances
◦ Held together by electrostatic attraction
between ions

Metals
◦ Held together by a “sea” of electrons around
positive ions
Ionic Solids
 Held
together by electrostatic
attraction between the ions
◦ High melting points
◦ Soluble in ppolar solvents
◦ Solids are poor electrical conductors
 Examples
◦ Sodium Chloride
◦ Cesium Chloride
10
Network Covalent Solids

Held together by covalent bonds (shared
electrons)
◦ High melting and boiling points – strong
covalent bonds
gy
◦ Insoluble in most solvents – too much energy
required to break covalent bonds upon
dissolution
◦ Poor electrical conductors

Examples
◦ Diamond
◦ Graphite
Metals

Held together by a “sea” of electrons around
positive ions
◦
◦
◦
◦

Electrically and thermally conductive
Ductile (hammerable) and malleable (pulled into wire)
L
Lustery
Insoluble in most solvent
Examples
◦ Gold
◦ Silver
◦ Platinum
Molecular Solids

Held together by weak dispersion or
dipole forces
◦ Low Melting and boiling points
◦ Non-conductor
◦ Typically
T
ll soluble
l bl in non-polar
l solvents
l

Examples
◦ Carbon Dioxide – “Dry Ice” (CO2)
◦ Iodine (I2)
11
Evaporation

Process where higher energy molecules on the
surface of a liquid escape into the gas phase
T 2 > T1
• Only higher
energy
molecules
l
l
have sufficient
energy to
evaporate
• The liquid
becomes
cooler
Condensation
 Liquefaction
of the vapor phase
 Opposite of evaporation
 Exothermic (gives off energy)
Dynamic Equilibrium
In a closed system composed of a liquid
and its vapor, a dynamic equilibrium will
eventually be achieved
 The rate of evaporation will equal the rate
of condensation
Liquid
Vapor
 No change the amount of liquid or vapor
occurs
 Rates are NOT zero, just equal

12
Vapor Pressure
The partial pressure of vapor molecules above
the surface of a liquid at equilibrium at a given
temperature
 Vapor pressure always increases with increasing
temperature
 Liquids which vaporize easily are called volatile –
they have high vapor pressures
 The weaker the forces of attraction, the
higher the vapor pressure (the lower the
boiling point)

Clausis--Clapeyron Equation
Clausis

Vapor Pressure is dependent upon Temperature (K)
As Temperature increases Vapor Pressure
increases
 H vap
 P 
ln  2  
R
 P1 
 1
1 



T2 
 T1
• R is the ideal gas constant = 8.314 J/mole-K
• Hvap is the heat of vaporization
• Compound specific (J/mol), but constant for
each compound
Boiling Point


Boiling Point is defined as the temperature where
the vapor pressure of liquid is equal to the external
pressure (when taken at 760 torr or 1 atm it is
called
ll d the
h normall boiling
b ili point)
i )
Boiling points are directly related to the strength of
the intermolecular forces involved
◦ The higher the boiling point the higher the strength of
the intermolecular forces
13
Effect of Pressure on BP
Location
Atm Pressure BP of H2O
Galveston
760 mm Hg
100°C
Denver
630 mm Hg
95°C
Mt Everest
280 mm Hg
70°C
Heating Curve
Heat Calculation
 Calculate
the amount of heat required
to convert 100 g of water at 10.0°C to
steam at 105.0 °C.
◦ Specific heat of liquid water = 4.18 J/ (g -°C)
◦ Specific heat of steam = 2.03 J/ (g -°C)
◦ Specific heat of vaporization of water = 2260 J/g
14
Heat Calculation
100 g
x 4.18 J/(g-°C)
100 g
x 2260 J/g
100 g
X 2.03 J/(g-°C)
x (100 – 10) °C
=
37620 J
= 226,000
226 000 J
X (105 –100) °C
=
Total
1015 J
= 264,635 J
= 265 kJ
Phase Diagrams
Critical
Point
Critical
Pressure
Melting Curve
Supercritical
Fluid
Pressure
Liquid
Solid
Vapor Pressure
Curve
Triple Pt
Vapor
Sublimation
Curve
Triple Point - gas,
liquid, solid co-exist
at equilibrium
Critical Temp – T
above which a gas
can NOT be liquefied
Critical Pressure – P
required to liquefy a
gas at its Critical
Temp
Critical Point – The
point defined by C.T.
and C.P.
Temperature
Critical
Temp
Definitions
Critical
Point
Critical
Pressure
Melting Curve
Liquid
Pressure
Evaporation
/Condensation
 Sublimation
/Deposition
 Fusion /Freezing

Solid
Supercritical
Fluid
Vapor Pressure
Curve
Triple Pt
Vapor
Sublimation
Curve
Temperature
Critical
Temp
15
You Can’t Skate On Dry Ice
Water
Melting Curve
Pressure
Liquid
Supercritical
Fluid
Vapor Pressure
Curve
Critical
Pressure
Melting Curve
Liquid
Pressure
Critical
Point
Critical
Pressure
Solid
Carbon Dioxide
Solid
Critical
Point
Supercritical
Fluid
Vapor Pressure
Curve
Triple Pt
Triple Pt
Vapor
Vapor
Sublimation
Curve
Sublimation
Curve
Temperature
Critical
Temp
Temperature
Critical
Temp
16