Liquids%20and%20Solids

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13

• Liquids and Solids

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Chapter Goals

• Kinetic-Molecular Description of Liquids and
  Solids
• Intermolecular Attractions and Phase Changes
  ??????????
• The Liquid State
• Viscosity ??
• Surface Tension ????
• Capillary Action ????
• Evaporation ??
• Vapor Pressure ???
• Boiling Points and Distillation ?????
• Heat Transfer Involving Liquids ??????

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Chapter Goals

• The Solid State
• Melting Point ??
• Heat Transfer Involving Solids ??????
• Sublimation and the Vapor Pressure of Solids
  ?????????
• Phase Diagrams (P versus T) ??
• Amorphous Solids and Crystalline Solids?????????
• Structures of Crystals ?????
• Bonding in Solids ?????
• Band Theory of Metals??????

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Kinetic-Molecular Description of 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 ???.
• They easily flow.
• The intermolecular attractions in liquids and
  solids are strong. ??????????????

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Kinetic-Molecular Description of Liquids and
Solids

• Schematic representation of the three common
  states of matter.

• The process in which a liquid changes to a solid
• Solidification
• Crystallization
• A more specific term
• The formation of a very ordered solid material

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Kinetic-Molecular Description of Liquids and
Solids

• If we compare the strengths of interactions among
  particles and the degree of ordering of
  particles, we see that
• Gaseslt Liquids lt Solids
• Miscible????liquids are soluble in each other.
• Examples of miscible liquids
• Water dissolves in alcohol.
• A drop of red ink in the water
• Gasoline dissolves in motor oil.

Diffusion
The miscibility of two liquids refers to their
ability to mix and produced a homogeneous
solution ????
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• Diffusion in solids very slowly

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Kinetic-Molecular Description of Liquids and
Solids

• Immiscible liquids are insoluble in each other.
• Two examples of immiscible liquids
• Water does not dissolve in oil.
• Water does not dissolve in cyclohexane ???

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Intermolecular Attractions and Phase Changes

• Intermolecular forces ??????
• The forces between individual particles (atoms,
  molecules, ions) of a substance
• Intramolecular forces ??????
• Covalent and ionic bonds within compounds

Covalent bonds ????
Intramolecular forces
Intermolecular forces
hydrogen bonding ??
12

• ?????????????? (??????????) ?????,??????????????,?
  ????????????????????

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Intermolecular Attractions and Phase Changes

• There are four important intermolecular
  attractions.
• This list is from strongest attraction to the
  weakest attraction.
• Ion-ion interactions ??- ?????
• The force of attraction between two oppositely
  charged ions is governed by Coulombs law ????.

Energy has the units of forces x distance, F x
d The energy of attraction, E
F ?
q and q- are the ion charges. d is the distance
between the e ions F is the force of attraction
E ?
? ????????,????????????????????????????????,????
?????
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Intermolecular Attractions and Phase Changes

• Most ionic bonding is strong
• ?have relatively high melting points
• Ionic substances containing multiply charged
  ions, such as Al3, Mg2, O2-, and S2-, usually
  have higher melting and boiling points than ionic
  compounds containing only singly charged ions,
  such as Na, K, F-, Cl-.
• For a series of ions of similar charges, the
  closer approached ions of smaller ions result in
  stronger interionic attractive forces and higher
  melting points (NaF, NaCl, NaBr)

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Intermolecular Attractions and Phase Changes

• Coulombs law determines
• The melting and boiling points of ionic
  compounds.
• The solubility of ionic compounds.
• Example 13-1 Arrange the following ionic
  compounds in the expected order of increasing
  melting and boiling points.
• NaF, CaO, CaF2
• NaF- lt Ca2F2-ltCa2O2-

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Intermolecular Attractions and Phase Changes

• Dipole-dipole interactions ??- ????
• ??????,??????,????????????????????,???
• ?????????????????
• ??-???????????????????????,???????????????????
• BrF

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Intermolecular Attractions and Phase Changes

• Hydrogen bonding ??
• Very strong dipole-dipole interaction
• Strong hydrogen bonding occurs among polar
  covalent molecules containing H (hydrogen bond
  donor) and one of the three small, highly
  electronegative elements F, O or N (hydrogen
  bond acceptor)
• Consider H2O a very polar molecule.

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Intermolecular Attractions and Phase Changes
Hydrogen bond
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Intermolecular Attractions and Phase Changes

• London Forces (Dispersion forces ???)
• They are the weakest of the intermolecular
  forces.
• This is the only attractive force in nonpolar
  molecules such as O2, N2 and monatomic species
  such as the noble gases.
• Without dispersion forces, such substance could
  not condense to form liquids or solidify to form
  solids

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Intermolecular Attractions and Phase Changes

• In a group of Ar atoms the temporary dipole in
  one atom induces other atomic dipoles.
• The positively charged nucleus
• The electron clouds of an atom in nearby
  molecules
• Dispersion forces are generally stronger for
  molecules that are larger
• They exist in all substances

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????? (London dispersion forces) ?????????????????
???,???????????????,???????????? (instantaneous
dipole moments),???????????????????????,????????,?
?????????-?????????
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• ?????????
• ?????????????,?????????????,?????????????????,????
  ??????????????-??????
• ????????????????????????????????,?????????????????
  ??
• ??????????????,????????????????

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?????
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??????
?????? ???????? ??
??-????? ?????(???????????????????) H2O?HCl
?? X?H????Y ?????X?H????????Y??????????,????????????-?????(?X ? F?N?O) H2O???? H2O
??-??????? ?????????????????? H2O????I2
????? ?????? I2????I2
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?????????
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Intermolecular Attractions and Phase Changes

• Example 13-1 Intermolecular Forces
• Identify the type of intermolecular forces that
  are present in a condensed phase sample of each
  of the following. For each, make a sketch,
  including a few molecules, that represents the
  major type of force. (a) water, H2O (b) iodine,
  I2 (c) nitrogen dioxide, NO2.
• water, H2O
• polar
• H, O ?hydrogen bonds
• The London forces
• (b) iodine, I2
• Nonpolar ?the London forces
• (c) nitrogen dioxide, NO2
• polar
• Dipole-dipole interactions
• The London forces

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The Liquid State

• Viscosity ??
• Viscosity is the resistance to flow.
• For example molasses??, syrup??or honey.
• Oil for your car is bought based on this
  property.
• 10W30 or 5W30 describes the viscosity of the oil
  at high and low temperatures.
• Measure by Viscometer.
• The stronger of the intermolecular forces of
  attraction? the more viscosity
• Increasing the size and surface area of molecule
  genernally results in increased viscosity
• Temperature increase
• ?viscosity decrease

Pentane C5H12??
????????????????,????????,????????????????????????
?????,????????,?????
Dodecane C12H26???
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The Liquid State

• Surface Tension ????
• 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.

??????????????????!?????????????????????????,?????
??????????,??????????????????????,??????
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The Liquid State

• Capillary Action ????
• Capillary action is the ability of a liquid to
  rise (or fall) in a glass tube or other container
  ????????????????????,?????????,???????????????????
  
• 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 gt cohesive forces
• Capillary fall implies that the
• Cohesive forces gt adhesive forces
• The smaller the bore, the higher the liquid
  climbs

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The Liquid State

• Capillary action also affects the meniscus of
  liquids.????????????????

concave
convex
Adhesive forces gt cohesive forces
Adhesive forces lt cohesive forces
???????????????????,?????????
???????????????????,?????????
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The Liquid State

• Evaporation ??
• Evaporation is the process in which molecules
  escape from the surface of a liquid and become a
  gas.????????????????

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The Liquid State

• Evaporation is temperature dependent.
• The rate of evaporation increases as temperature
  increases

Only the higher-energy molecules can escape from
the liquid phase
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The Liquid State

• Condensation ?? ??
• In a closed container
• ?Dynamic equilibrium

evaporation
vapor
liquid
condensation

• LeChateliers Principle ??????
• If the vessel were left open to the air, the
  equilibrium could not be reached.
• A liquid can eventually evaporated entirely

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The Liquid State

• Vapor Pressure ???
• Vapor pressure is the pressure exerted by a
  liquids vapor on its surface at equilibrium.
• Vapor Pressure (torr) and boiling point for three
  liquids at different temperatures.
• 0oC 20oC 30oC normal boiling
  point
• diethyl ether?? 185 442 647 36oC
• Ethanol ?? 12 44 74 78oC
• water ? 5 18 32 100oC
• Vapor pressures of liquids always increase as
  temperature increase

????????????????????????
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The Liquid State

• Vapor Pressure ???
• Easily vaporized liquids are called volatile
  liquids?????
• Stronger cohesive forces tend to hold molecules
  in the liquid state
• Methanol molecules are strongly linked by
  hydrogen bond ?lower vapor pressure
• Dispersion forces increase with increasing
  molecular size? larger molecules have lower vapor
  pressure

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The Liquid State

• Vapor Pressure ???
• Vapor pressure can be measured with manometers???

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The Liquid State

• Vapor Pressure as a function of temperature

????????????????????????
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The Liquid State

• Boiling Points ?? and Distillation??
• The boiling point is the temperature at which the
  liquids vapor pressure is equal to the applied
  pressure (usually atmospheric???).
• The normal boiling point???? is the boiling point
  when the pressure is exactly 1 atm (760 torr).
  ????????????,?????????
• If the applied pressure is lower than 1 atom
• ?water boil below 100oC.
• ??????????????????????????????,?????
• Distillation ??is a method we use to separate
  mixtures of liquids based on their differences in
  boiling points.

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The Liquid State

• Distillation??
• Different liquids have different vapor pressure
  and boil at different temperature.
• Distillation is a process in which a mixture or
  solution is separated into its components on the
  basis of the differences in boiling points of the
  components.
• Distillation is another vapor pressure
  phenomenon.

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The Liquid State

• Distillation??

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The Liquid State

• Heat Transfer Involving Liquids
• Heat must added to a liquid to raise its
  temperature
• The amount of heat that must be added to the
  stated mass of liquid to raise its temperature by
  one degree
• ?Specific heat (J/goC) ??
• molar heat capacity (J/moloC)?????
  (?1mol????1oC?????
• Molar heat (enthalpy?) of vaporization?????
• ???1 mol ??????????????
• ??100oC ???????40.7 kJ/mol

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The Liquid State
x
40.7KJ/mol 40.7x
2.26x103 J/g
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The Liquid State

• Heat Transfer Involving Liquids
• Condensation
• Heat of condensation ???
• liquid heat vapor
  

evaporation
condensation
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The Liquid State

• Example 13-2
• How much heat is released by 2.00 x 102 g of H2O
  as it cools from 85.0oC to 40.0oC? The specific
  heat of water is 4.184 J/goC.
• qmC?T
• ?J 2.00 x 102 g x (4.184J/goC) x(
  85.0-40oC)
• ?J 3.76x104 J 37.6kJ

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The Liquid State

• Example 13-3
• The molar heat capacity of ethyl alcohol, C2H5OH,
  is 113 J/moloC. How much heat is required to
  raise the T of 125 g of ethyl alcohol from 20.0oC
  to 30.0oC?
• 1 mol C2H5OH 46.0 g

?mol C2H5OH 125g x
2.72 mol C2H5OH
?J 2.72mol x
x (30.0-20.0oC)
3.07 KJ
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The Liquid State

• The calculations we have done up to now tell us
  the energy changes as long as the substance
  remains in a single phase.
• Next, we must address the energy associated with
  phase changes.
• For example, solid to liquid or liquid to gas and
  the reverse.
• Heat of Vaporization is the amount of heat
  required to change 1.00 g of a liquid substance
  to a gas at constant temperature.
• Heat of vaporization has units of J/g.
• Heat of Condensation is the reverse of heat of
  vaporization, phase change from gas to liquid.

1.0g H2O(l) at 100oC
1.0g H2O(g) at 100oC
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The Liquid State

• Molar heat of vaporization or DHvap
• The DHvap is the amount of heat required to
  change 1.00 mole of a liquid to a gas at constant
  temperature.
• DHvap has units of J/mol.
• Molar heat of condensation
• The reverse of molar heat of vaporization is the
  heat of condensation.

1.0mol H2O(l) at 100oC
1.0mol H2O(g) at 100oC
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The Liquid State
??/???? (heating-cooling curve)
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The Liquid State

• Example 13-4
• How many joules of energy must be absorbed by
  5.00 x 102 g of H2O at 50.0oC to convert it to
  steam at 120oC? The molar heat of vaporization
  of water is 40.7 kJ/mol and the molar heat
  capacities of liquid water and steam are 75.3
  J/mol oC and 36.4 J/mol oC, respectively.

50oC H2O(l)
100oC H2O(l)
100oC H2O(g)
120oC H2O(g)
?mol H2O 500g x
27.8 mol H2O
?J 27.8mol x
x (100.0-50.0oC)
1.05x105 J
?J 27.8mol x
11.31x105 J
?J 27.8mol x
X (120.0-100.0oC)
0.20x105 J
Total J 1.05x105 11.31x105 0.2x105
12.56 x105 J or 1.26x103 kJ
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The Liquid State

• Example 13-5
• If 45.0 g of steam at 140oC is slowly bubbled
  into 450 g of water at 50.0oC in an insulated
  container, can all the steam be condensed?

?mol steam 45.0g x
2.5 mol steam
?mol H2O 450g x
25.0 mol H2O
Calculate the amount of heat required to condense
the steam
2.5mol x(36.4J/moloC)x(140-100oC)2.5 mol x
(40.7kJ/mol) 105.4kJ
Calculate the amount of heat available in the
liquid water
25.0mol x (75.3J/moloC)x(100.0-50oC) 94.1kJ
Amount of heat to condense all of the steam is
105kJ Amount of heat that the liquid water can
absorb is 94.1kJ Thus all of the steam cannot be
condensed
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The Liquid State

• Example 13-6
• Arrange the following substances in order of
  increasing boiling points.
• C2H6, NH3, Ar, NaCl, AsH3
• Ar lt C2H6 lt AsH3 lt NH3 lt NaCl
• nonpolar nonpolar polar very polar
  ionic
• London London dipole-dipole H-bonding
  ion-ion

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Example 13-2 Heat of Vaporization Calculate the
amount of heat, in joules, required to convert
180.0grams of water at 10.0oC to steam at 105.0oC.
10oC H2O(l)
100oC H2O(l)
100oC H2O(g)
105oC H2O(g)
?J 180.0g x
x (100.0-10.0oC)
0.67x105 J
?J 180.0g x
4.07x105 J
?J 180.0g x
x (105.0-100.0oC)
0.01820x105 J
Total J 0.67x105 4.07x105 0.0182x105
4.76 x105 J
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Example 13-3 Heat of Vaporization Compare the
amount of cooling experienced by an individual
who drinks 400ml of ice water (0.0oC) with the
amount of cooling experienced by an individual
who sweats out 400ml of water. Assume that the
sweat is essentially pure water and that all of
it evaporates. The density of water is very
nearly 1.00g/ml at both 0.0oC and 37oC, average
body temperature. The heat of vaporization of
water is 2.41kJ/g at 37oC. The amount of heat
lost by perspiration the amount of heat
required to vaporize 400g of water at
37oC Raising the temperature of 400.0g of water
from 0oC to 37oC
?J 400.0g x
x (37.0-0.0oC)
6.19x104 J or 61.9kJ
Evaporating 400.o ml of water at 37oC requires
?J 400.0ml x
x (2.41x103 J/g)
9.64x105 J or 964kJ
Sweating removes more heat than drinking ice water
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• Example 13-5 Boiling Points Versus
  Intermolecular Forces
• Predict the order of increasing boiling points
  for the following H2S H2O CH4 H2 KBr.
• KBr is ionic, it boils at the highest
  temperature
• Hydrogen bond H2O the next highest temperature
• polar covalent substance H2S
• CH4 H2 are nonpolar .
• CH4 is larger than H2
• so the dispersion forces are
  stronger in CH4
• ? CH4 boils at a higher
  temperature than H2
• H2 lt CH4 lt H2S lt H2O lt KBr

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The Solid State

• Normal Melting Point
• The normal melting point is the temperature at
  which the solid melts (liquid and solid in
  equilibrium) at exactly 1.00 atm of pressure.
• Freezing point ? melting point
• The melting point increases as the strength of
  the intermolecular attractions increase.

melting
liquid
solid
freezing
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The Solid State

• Which requires more energy?

Ion-ion interaction
or
Hydrogen bond
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Heat Transfer Involving Solids

• Heat of Fusion
• Heat of fusion ???is the amount of heat required
  to melt one gram of a solid at its melting point
  at constant temperature.

Heat of fusion
334J
1.00g H2O(l) at 0oC
1.00g H2O(s) at 0oC
-334J
Heat of solidification

• Heat of solidification??? is the reverse of
  the heat of fusion.

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Specific heat
Heat of vaporization
Specific heat ??
Heat of fusion ???
Specific heat ??
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Heat Transfer Involving Solids

• Molar heat of fusion ????? or ?Hfusion
• The molar heat of fusion is the amount of heat
  required to melt a mole of a substance at its
  melting point.
• The molar heat of solidification is the reverse
  of molar heat of fusion

Molar heat of fusion
6012J
1.00mole H2O(l) at 0oC
1.00mole H2O(s) at 0oC
-6012J
Molar heat of solidification
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Heat Transfer Involving Solids

• The heat of fusion depends on the intermlacular
  forces of attraction in the solid state
• Heats of fusion are usually higher for substances
  with higher melting points.

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Heat Transfer Involving Solids

• Example 13-7
• Calculate the amount of heat required to convert
  150.0 g of ice at -10.0oC to water at 40.0oC.
• specific heat of ice is 2.09 J/goC

-10oC H2O(s)
0oC H2O(s)
0oC H2O(l)
40oC H2O(l)
?J 150.0g x
x (10.0oC)
3.14x103 J
?J 150.0g x
5.01x104 J
?J 150.0g x
x (40.0-0.0oC)
2.51x104 J
Total J 3.14x103 5.01x104 7.83x104
7.83x104J
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Example 13-6 Heat of Fusion The molar heat of
fusion, ?Hfus, of Na is 2.6kJ/mol at its melting
point, 97.5oC, How much neat must be absorbed by
5.0g of sodium Na at 97.5oC to melt it?
?J 5.0g x
x
0.57 kJ
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Example 13-2 Heat of Fusion Calculate the amount
of heat that must be absorbed by 50.0g of ice at
-12.0oC to convert to water at 20.0oC.
-10oC H2O(s)
0oC H2O(s)
0oC H2O(l)
20oC H2O(l)
?J 50.0g x
x (0.0-(-12).0oC)
1.25x103 J
?J 50.0g x
16.7x103 J
?J 50.0g x
x (20.0-0.0oC)
4.18x103 J
Total J 1.25x103 16.7x103 4.18x103
22.1 kJ
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Sublimation and the Vapor Pressure of Solids

• Sublimation??
• In the sublimation process the solid transforms
  directly to the vapor phase without passing
  through the liquid phase.
• Solid CO2 or dry ice does this well.
• iodine

sublimation
gas
solid
deposition
Sublimation can be used to purify volatile solids
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Exothermic ??
??
??
??
Endothermic ??
Transitions among the three states of matter
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Phase Diagrams (P versus T)??

• Phase diagrams are a convenient way to display
  all of the different phase transitions of a
  substance.
• The equilibrium pressure-temperature (??-??)
  relationships among the different phases of a
  given pure substance in a closed system (????).

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Phase Diagrams (P versus T)??

• This is the phase diagram for water.

• Negative slope of line AB
• Ice is less dense than liquid
• The network of hydrogen bonding in ice is more
  extensive than that in a liquid water

AB Melting curve AD Sublimation curve
???
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Phase Diagrams (P versus T)

• Compare waters phase diagram to carbon dioxides
  phase diagram.

Critical point
???
Liquid CO2 cannot exist at atmospheric pressure
For CO2 critical point is at 31oC and 73 atm For
H2O critical point is at 374oC and 218 atm
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To illustrate the use of a phase diagram in
determining the physical state or states of a
system under different sets of pressures and
temperatures
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Amorphous????Solids andCrystalline????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.

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All crystal contain regularly repeating
arrangements of atoms, molecules or ions

• Unit cell ????
• The Lengths of its edges (a,b,c)
• The angles between the edges
• (a,b,g)
• Three-dimensional arrangement
• Lattice point ??

???
??
???
???
???
???
???
??
????
????
??
???
???
?????
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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.

??
???
???
???
???
???
??
?????
???
???
????
????
???
???
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? 14.14
???????????????????????,????????????????????(a)???
??????????????????(b)????????????,????????Na
?Cl- ???(c)????????,???????????H2O
?????????????????
P.425
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? 14.3
?????????
P.426
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Structure of Crystals

• We shall look at the three variations of the
  cubic crystal system.
• Simple cubic unit cells???
• The balls represent the positions of atoms, ions,
  or molecules in a simple cubic unit cell.
• 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.

77
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.

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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.

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• Isomorphous
• refers to crystals having the same atomic
  arrangement
• Polymorphous
• Refers to the substances that crystallize in more
  than one crystalline arrangement

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Bonding in Solids

• Four categories
• Metallic solids???? - Ionic solids????
• Molecular solids???? - Covalent solids

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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

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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???

??
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Bonding in Solids

• Ionic Solids have ions that occupy the positions
  in the unit cell.
• Examples of ionic solids include
• CsCl, NaCl, ZnS

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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, ..

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• ???????????????
• ????????????????
• ????? (electron sea model)??????????????,???????
  ??????????????????
• ?????????,???????????,??????????????????

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• ?? (alloy)?????????????????????
• ?????????
• ????? (substitutional alloy)?????????????????,??
  ??,?????1/3 ???????????,
• ????? (interstitial alloy)????????????????? (??)
  ????????????????????????????????

P. 428