Get documents about "Liquids and Solids
Shared by: ert554898
-
Stats
- views:
- 2
- posted:
- 3/19/2012
- language:
- pages:
- 76
Document Sample
Liquids and Solids
States of Matter
• The fundamental difference between states of matter is
the distance between particles.
2
States of Matter
• Because in the solid and liquid states particles are closer
together, we refer to them as condensed phases.
3
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.
4
Kinetic-Molecular Description of Liquids and
Solids
• Miscible liquids are soluble in each
other.
• Examples of miscible liquids:
• Water dissolves in alcohol.
• Gasoline dissolves in motor
oil.
• Immiscible liquids are insoluble in QuickTime™ and a
Cinepak decompressor
each other. are neede d to see this picture.
• Two examples of immiscible
liquids:
• Water does not dissolve in
oil.
• Water does not dissolve in
cyclohexane.
5
Intermolecular Forces
• Have studied intramolecular
forces—the forces holding
atoms together to form
molecules.
• Now turn to forces between
molecules — intermolecular
forces.
• Forces between molecules,
between ions, or between
molecules and ions.
6
Intermolecular Forces
• The attractions between molecules are not nearly as strong as the
intramolecular attractions that hold compounds together.
• They are, however, strong enough to control physical properties such
as boiling and melting points, vapor pressures, and viscosities.
• These intermolecular forces as a group are referred to as van der
Waals forces.
7
Summary of Intermolecular Forces
Energy of
Interaction Factors
interactions
Ion-ion Ion charge; size 400-4000 kJ/mol
Ion charge; dipole
Ion-dipole 40-600 kJ/mol
moment
Hydrogen bond Hydrogen bond 25-200 kJ/mol
Dipole-dipole Dipole moment 5-25 kJ/mol
Dipole-induced Dipole moment;
2-10 kJ/mol
dipole polarizability
Induces dipole polarizability 0.05-40 kJ/mol
8
Intermolecular Forces: Ion-ion Interactions
• Na+ — Cl- in salt.
• These are the strongest
forces.
• Lead to solids with high
melting temperatures.
• NaCl, mp = 800°C
• MgO, mp = 2800°C
9
Ion–ion Interactions
• Ion–ion interactions are those electrostatic attractions that
form between and ion and a oppositely charged ion of an
ionic molecule.
• Force of attraction between two oppositely charged ions is
determined by coulomb’s law.
10
Intermolecular Attractions and Phase Changes
• Coulomb’s law & the attraction energy determine:
• melting & boiling points of ionic compounds
• the solubility of ionic compounds
11
Attraction Between Ions and Permanent Dipoles
12
Attraction Between Ions and Permanent Dipoles
• Water is highly polar and can
interact with positive ions to
give hydrated ions in water.
13
Dipole-Dipole Interactions
• 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.
• Dipole–dipole forces exist
between neutral molecules.
• For molecules of approximately
equal mass and size, the
intermolecular attractions
increase with increasing polarity.
14
Dipole–dipole Forces
• Influence of dipole-dipole forces is seen in the boiling
points of simple molecules.
• Compd mol. Wt. Boil point
• N2 28 -196°C
• CO 28 -192°C
• Br2 160 59°C
• ICl 162 97°C
15
Hydrogen Bonding
• The dipole-dipole interactions
experienced when H is bonded to N,
O, or F are unusually strong.
• We call these interactions hydrogen
bonds.
16
Hydrogen Bonding
17
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.
18
London Dispersion Forces
• At that instant, then, the helium atom is polar, with an
excess of electrons on the left side and a shortage on the
right side.
19
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.
20
London Dispersion Forces
• London dispersion forces, or dispersion forces, are
attractions between an instantaneous dipole and an
induced dipole.
21
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.
22
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.
23
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.
24
Forces Involving Induced Dipoles
• The induced forces between I2 molecules are very weak,
so solid I2 sublimes (goes from a solid to gaseous
molecules).
25
Summary of Intermolecular Forces
Energy of
Interaction Factors
interactions
Ion-ion Ion charge; size 400-4000 kJ/mol
Ion charge; dipole
Ion-dipole 40-600 kJ/mol
moment
Hydrogen bond Hydrogen bond 25-200 kJ/mol
Dipole-dipole Dipole moment 5-25 kJ/mol
Dipole-induced Dipole moment;
2-10 kJ/mol
dipole polarizability
Induces dipole polarizability 0.05-40 kJ/mol
26
Test Your Knowledge
• Predict the order of increasing boiling points for: H2S;
H2O; CH4; H2; KBr.
Hint: What intermolecular forces exist?
H2 < CH4 < H2S < H2O < KBr
27
Liquids
• Molecules are in constant
motion
• There are appreciable
intermolecular forces
• Molecules are close together
• Liquids are almost
incompressible
• Liquids do not completely fill
the container
28
Liquids
• Evaporation is the process in which molecules escape
from the surface of a liquid and become a gas.
• The molecules must have sufficient energy to break the
intermolecular forces.
• Evaporation is temperature dependent.
29
The Liquid State
• Only a small fraction of the molecules in a liquid have
enough energy to escape.
• However, as the temperature increases, the fraction of the
molecules with “escape energy” increases.
• The higher the temperature, the faster the rate of
evaporation.
30
Dynamic Equilibrium
• In a closed container, once the rates of vaporization and
condensation are equal, the total amount of vapor and
liquid will not change.
• Evaporation and condensation are still occurring, but
because they are opposite processes, there is no net gain
or loss or either vapor or liquid.
• When two opposite processes reach the same rate so that
there is no gain or loss of material, we call it a dynamic
equilibrium.
• This does not mean that there are equal amounts of
vapor and liquid—it means that they are changing by
equal amounts.
31
The Liquid State
• Vapor Pressure
• Vapor pressure is the pressure exerted by a liquid’s vapor on
its surface at equilibrium.
32
The Liquid State
• Boiling Points
• The boiling point is the
temperature at which the
liquid’s vapor pressure is
equal to the applied pressure.
• The normal boiling point is the
boiling point when the
pressure is exactly 1 atm.
33
• Liquid boils when its vapor pressure equals atmospheric pressure.
Liquid boils when its vapor pressure
equals atmospheric pressure.
34
Boiling Point at Lower Pressure
• When pressure is lowered, the vapor pressure can equal
the external pressure at a lower temperature.
QuickTime™ and a
Cinepak decompressor
are neede d to see this picture.
35
Consequences of Vapor Pressure Changes
• When can cools, the vapor pressure of water drops. The
pressure in the can is less than that of atmosphere, so the
can is crushed.
QuickTime™ and a
Cinepak decompressor
are neede d to see this picture.
36
The Liquid State
• Viscosity
• Viscosity is the resistance to
flow.
• For example, compare how
water pours out of a glass
compared to molasses, QuickTime™ and a
syrup or honey. Cinepak decompressor
•
are neede d to see this picture.
Oil for your car is bought
based on this property.
• 10W30 or 5W30 describes
the viscosity of the oil at
high and low temperatures.
37
The Liquid State
• Surface tension is a measure
of the unequal attractions that
occur at the surface of a
liquid.
• The molecules at the surface
are attracted unevenly.
38
The Liquid State
• Floating paper clip demonstration of surface tension.
QuickTime™ and a
Cinepak decompressor
are neede d to see this picture.
39
The Liquid State
• Capillary action is the ability of a liquid to rise (or fall) in a
glass tube or other container
QuickTime™ and a
Graphics decompres sor
are needed to see this picture.
40
The Liquid State
• Cohesive forces are the forces that hold liquids together.
• Adhesive forces are the forces between a liquid and
another surface.
• Capillary rise implies that the:
•Adhesive forces > cohesive forces
• Capillary fall implies that the:
•Cohesive forces > adhesive forces
41
The Liquid State
• The curving of the liquid surface in a thin tube is
due to the competition between adhesive and
cohesive forces.
• The meniscus of water is concave in a glass tube
because its adhesion to the glass is stronger
than its cohesion for itself.
• The meniscus of mercury is convex in a glass
tube because its cohesion for itself is stronger
than its adhesion for the glass.
• Metallic bonds are stronger than intermolecular
attractions.
42
Amorphous Solids and Crystalline Solids
• Amorphous solids do not have a well ordered molecular
structure.
• Examples of amorphous solids include waxes, glasses,
asphalt.
• Crystalline solids have well defined structures that consist of
extended array of repeating units called unit cells.
• Crystalline solids display X-ray diffraction patterns which
reflect the molecular structure.
• The Bragg equation, detailed in the textbook, describes
how an X-ray diffraction pattern can be used to determine
the interatomic distances in crystals.
43
Structure of Crystals
• Unit cells are the smallest repeating unit of a crystal.
• As an analogy, bricks are repeating units for buildings.
• There are seven basic crystal systems.
• We shall look at the three variations of the cubic crystal
system.
44
Structure of Crystals
• In a simple cubic unit cell
each atom, ion, or molecule
at a corner is shared by 8 unit
cells
• Thus 1 unit cell contains
8(1/8) = 1 atom, ion, or
molecule.
45
Structure of Crystals
• Body centered cubic (bcc)
has an additional atom, ion,
or molecule in the center of
the unit cell.
• On a body centered cubic unit
cell there are 8 corners + 1
particle in center of cell.
• 1 bcc unit cell
• contains 8(1/8) + 1 = 2
particles.
46
Structure of Crystals
• A face centered cubic (fcc)
unit cell has a cubic unit cell
structure with an extra atom,
ion, or molecule in each face.
• A face centered cubic unit cell
has 8 corners and 6 faces.
• 1 fcc unit cell contains
• 8(1/8) + 6(1/2) = 4
particles.
47
Unit Cell Summary
48
Bonding in Solids
• Molecular Solids have molecules in each of the positions
of the unit cell.
• Molecular solids have low melting points, are volatile,
and are electrical insulators.
• Examples of molecular solids include:
• water, sugar, carbon dioxide, benzene
49
Bonding in Solids
• Covalent Solids have atoms that are covalently bonded to
one another
• Some examples of covalent solids are:
• Diamond, graphite, SiO2 (sand), SiC
50
Bonding in Solids
• Ionic Solids have ions that occupy the positions in the unit
cell.
• Examples of ionic solids include:
• CsCl, NaCl, ZnS
51
Bonding in Solids
• Metallic Solids may be thought of as positively charged
nuclei surrounded by a sea of electrons.
• The positive ions occupy the crystal lattice positions.
• Examples of metallic solids include:
• Na, Li, Au, Ag, …
52
Pythagoras
a = edge
d = face diagonal =
D (d2 = a2 + a2 = 2a2)
a d
a D = body diagonal =
a (D2 = d2 + a2 = 2a2 + a2 = 3a2)
55
Face Centered Cubic
1/8 atom
x 8 corners = 1atom
corner
4 atoms/unit cell
1/2 atom
x 6 facess = 3atom
face
atoms in contact along face diagonal (d)
∴ d = 4r =
a= or r=
V = a3 =
56
Examples
• The unit cell of Ag is FCC with a side of 4.086Å.
• Calculate the volume of a Ag unit cell.
57
Examples
• Calculate the mass of a Ag unit cell.
• Calculate the density of Ag.
7.165x1022 g
d 23 3
10.50g / cm3
6.822x10 cm
58
Phase Diagrams (P versus T)
• Phase diagrams are a convenient way to display all of the
different phase transitions of a substance.
• This is the phase diagram for water.
60
Phase Diagrams (P versus T)
• Compare water’s phase diagram to carbon dioxide’s
phase diagram.
61
Transitions Between Phases
• As P and T increase, you finally reach the CRITICAL T
and P
Above critical T no
liquid exists no
matter how high the
pressure.
62
Heat Transfer
• Specific heat, J/g°c and molar heat capacity, J/mol°c is
the amount of heat that must be added to the stated mass
of liquid to raise its temperature by 1°C.
63
Heat Transfer
• The molar heat of vaporization, ΔHvap, is the amount of
heat that must be added to one mole of the liquid at its
boiling point to convert it to vapor with no temperature
change.
64
Heat Transfer
• The heat of condensation is the amount of heat that must
be removed from one mole of the gas at its boiling point to
convert it to liquid with no temperature change.
• Vapor → liquid + heat
65
Example
• Calculate the amount of heat, joules, required to convert
180g of water at 10.0°C to steam at 105°C.
q = smΔT = 180g (4.18J/g-C) (100.0- 10.0)C =
6.77 x 104 J
q = nΔHvap = 180g(2.26 x 103 J/g)
q = 4.07 x 105 J
q = smΔT = 180g(2.03 J/g-C)(105.0 - 100.0)C q = 1.8 x
103 J
q = 4.76 x 105 J
66
The Clausius–Clapeyron Equation
• The Clausius–Clapeyron equation gives the relationship
between temperature and vapor pressure.
67
The Liquid State
• In Denver the normal atmospheric pressure is 630 torr. At
what temperature does water boil in Denver?
P1 H 1 1
ln T T
P2 R 1 2
3 J
630torr 40.7x10 mol 1 1
ln
760torr J 372K T
8.314 2
mol K
1 1
ln 0.829 4895K
373K T2
4895
0.188 13.1
T2
4895
T2 368K or 95C
13.3
68
Changes of State
• Energy changes accompany all changes of state.
• When a change produces a less–ordered state, energy
must be supplied to overcome the intermolecular
forces.
• Thus, melting and vaporization are endothermic
processes.
69
Changes of State
• The heat of fusion is the
enthalpy (heat) change
associated with melting a
substance and is usually
expressed in kJ/mol.
70
Cooling Curves-melting Point
• Liquid → solid
• The of heat required to melt one gram of a solid at its
melting point is the heat of fusion and the molar heat of
fusion is the amount of heat required to melt one mole of a
solid at its melting point.
• ΔHsolidification = -ΔHfusion
71
Example
• Calculate the amount of heat that must be absorbed by
50.0g of ice at -12.0°C to convert it to water at 20.0°C.
q = smΔT = 50.0g(2.09 J/g-C)(0.0 - (-12.0)) = 1.25 x 103 J
q = nΔHfus = 50.0g(334 J/g) = 1.67 x 104 J
q = smΔT = 50.0g(4.18 J/g-C)(20.0 -0.0)C = 4.18 x 103 J
q = 2.21 x 104 J or 22.1 kJ
72
Phases
74
Heat and Changes of State
• What quantity of heat is required to melt 500. g of ice and heat
the water to steam at 100oC?
• To melt ice
• q = (500. g)(333 J/g) = 1.67 x 105 J
• To raise water from 0oC to 100oC
• q = (500. g)(4.2 J/g•C)(100 - 0)C = 2.1 x 105 J
• To evaporate water at 100oC
• q = (500. g)(2260 J/g) = 1.13 x 106 J
• Total heat energy = 1.51 x 106 J = 1510 kJ
75
Phases
76
