Intermolecular Attractions & Properties of Liquids and Solids
Intermolecular Attractions & the Properties of Liquids & Solids CHAPTER 12
description
Transcript of Intermolecular Attractions & the Properties of Liquids & Solids CHAPTER 12
Intermolecular Attractions & the Properties of Liquids & Solids
CHAPTER 12 Chemistry: The Molecular Nature of Matter, 6th edition
By Jesperson, Brady, & Hyslop
2
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Understand, describe, and rank in order of strength the types of
intermolecular forces.
Difference between bonds and intermolecular forces
Changes of state: heat of vaporization, fusion, & sublimation
Clausius-Clapyron equation
Heating and cooling curves: ΔH, phase transition temperatures
Phase diagrams
Solids: Unit cell, stoichiometry, packing patterns, XRD, common
types and their properties
3
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Lecture Road Map:
① Properties of gas, liquids, solids
② Intermolecular forces
③ Changes of state
④ Dynamic Equilibrium
⑤ Structure & Characterization of a solid
4
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Properties of gases, liquids, &
solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E 5
Intermolecular ForcesImportant differences between gases, solids, and liquids:oGases
oExpand to fill their container oLiquids
oRetain volume, but not shapeoSolids
o Retain volume and shape
6
Intermolecular Forceso Physical state of molecule depends on
o Average kinetic energy of particlesoRecall KE Tave
o Intermolecular Forces oEnergy of Inter-particle attraction
o Physical properties of gases, liquids and solids determined by oHow tightly molecules are packed togethero Strength of attractions between
molecules
7
o Converting gas liquid or solidoMolecules must get closer together
oCool or compress
o Converting liquid or solid gasoRequires molecules to move farther apart
oHeat or reduce pressure
o As T decreases, kinetic energy of molecules decreaseso At certain T, molecules don’t have
enough energy to break away from one another’s attraction
Intermolecular Attractions
8
Inter vs. Intra-Molecular Forceso Intramolecular forces
o Covalent bonds within molecule o Strong o Hbond (HCl) = 431 kJ/mol
o Intermolecular forces o Attraction forces between moleculeso Weako Hvaporization (HCl) = 16 kJ/mol
Cl H Cl H
Covalent Bond (strong) Intermolecular attraction (weak)
9
Electronegativity Review
Electronegativity: Measure of attractive force that one atom in a covalent bond has for electrons of the bond
10
Bond Dipoleso Two atoms with different electronegativity
values share electrons unequallyo Electron density is uneven
oHigher charge concentration around more electronegative atom
o Bond dipoles o Indicated with delta (δ) notationo Indicates partial charge has arisen
H F
o 11
o Net Dipoleso Symmetrical molecules
o Even if they have polar bondso Are non-polar because bond dipoles cancel
o Asymmetrical molecules o Are polar because bond dipoles do not cancelo These molecules have permanent, net dipoles
o Molecular dipoles oCause molecules to interactoDecreased distance between molecules increases
amount of interaction
COVALENTBOND
IONICBOND
POLARCOVALENT
BOND
CHCl3
TiO2
F2
CaBr2
✔
✔
✔
✔
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
13
GroupProblem
Identify the overall dipole moment for CHCl3
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
14
GroupProblem
Identify the overall dipole moment for these molecules:
Solubility
15
LIKE DISSOLVES LIKEpolar molecules dissolve in polar solvents
nonpolar molecules dissolve in nonpolar solvents
Polar SolventsWater: H2OMethanol: CH3OHEthanol: CH3CH2OHAcetone: (CH3)2COAcetic Acid: CH3CO2HAmmonia: NH3
Acetonitrile: CH3CN
Nonpolar SolventsPentane: C5H12
Hexane: C6H14
Cyclohexane: C6H12
Benzene: C6H6
Toluene: CH3C6H5 Chloroform: CHCl3Diethylether: (CH3CH2)2O
Which molecule will dissolve in water?
16
Vitamin A
Vitamin B12
GroupProblem
17
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
IntermolecularForces
18
Intermolecular Forces The forces of attraction or repulsion between neighboring particles (atoms or molecules).
+ / - charges attract one another - / - or + / + forces repel each other
rrKE
19
Intermolecular Forces o When substance melts or boils
o Intermolecular forces are broken, not covalent bonds
o Responsible for non-ideal behavior of gaseso Responsible for existence of condensed
states of mattero Responsible for bulk properties of matter
o Boiling points and melting points reflect strength of intermolecular forces
20
Types of Intermolecular Forces
① London dispersion forces② Dipole-dipole forces③ Hydrogen bonds④ Ion-dipole forces
o Ion-induced dipole forces
21
London-Dispersion Forceso When atoms near one another,
their valence electrons interacto Repulsion causes electron clouds
in each to distort and polarizeo Instantaneous dipoles result from
this distortiono Effect enhanced with increased
volume of electron cloud sizeo Effect diminished by increased
distance between particles and compact arrangement of atoms
22
London Dispersion ForcesAffects ALL molecules, both polar & nonpolar
Boiling Point (BP) is an indication of relative intermolecular force strength.
Ease with which dipole moments can be induced and thus London Forces depend on
① Distance between particles② Polarizability of electron cloud③ Points of attraction
oNumber atomsoMolecular shape (compact or elongated)
23
Polarizability = Ease with which the electron cloud can be distorted
Larger molecules often more polarizableo Larger number of less tightly held
electrons o Magnitude of resulting partial
charge is largero Larger electron cloud
24
GroupProblem
Which is more polarizable?F2 or I2?
Table 12.1 Boiling Points of Halogens and Noble Gases
Larger molecules have stronger London forces and thus higher boiling points.
26
Number of Atoms in Molecule
o London dispersion forces increase with the number atoms in molecule because more points of attraction
Formula BP at 1 atm, C Formula BP at 1 atm, CCH4 –161.5 C5H12 36.1
C2H6 –88.6 C6H14 68.7C3H8 –42.1 : :C4H10 –0.5 C22H46 327
Hexane, C6H14
BP 68.7 °C27
Which of the following molecules will have the highest boiling point?
Propane, C3H8
BP –42.1 °C
GroupProblem
28
Molecular Shapeo Increased surface area available for contact =
increased points of contact = increase in London Dispersion forces.oMore compact molecules:
Less surface area to interact with other molecules
oLess compact molecules:More surface area to interact with other molecules
29
• Small area for interaction
• Larger area for interaction
More compact – lower BP Less compact – higher BP
Which of the following molecules experience the strongest Dispersion forces?
30
GroupProblem
31
Types of Intermolecular Forces
① London dispersion forces② Dipole-dipole forces③ Hydrogen bonds④ Ion-dipole forces
o Ion-induced dipole forces
32
Dipole-Dipole Attractions
o Occurs only between polar molecules
o Proportional to distance between molecules
o Polar molecules tend to align their partial charges: + / -
o As dipole moment increases, intermolecular force increases
+ +
+ +
+ +
33
Dipole-Dipole AttractionsTumbling molecules
oMixture of attractive and repulsive dipole-dipole forces
o Attractions (- -) are maintained longer than repulsions(- -)
oGet net attraction o ~1–4% of covalent bond
In the liquid state, which species has the strongest intermolecular forces, CH4, Cl2, O2 or HF?
HFThe polar molecule
34
GroupProblem
35
Types of Intermolecular Forces
① London dispersion forces② Dipole-dipole forces③ Hydrogen bonds④ Ion-dipole forces
o Ion-induced dipole forces
36
Hydrogen Bondso Very strong dipole-dipole attraction: ~10% of a covalent
bond
o Occurs between H and highly electronegative atom (O, N, or F): H—F, H—O, and H—N bonds very polaroElectrons are drawn away from H giving atoms high
partial chargesoH only has one electron, so +
H presents almost bare proton
o –X almost full –1 charge
oElement’s small size, means high charge density
37
Examples of Hydrogen BondingH O
H
H O
H
H O
H
H N
H
H
H F H O
H
H F H N
H
H
H N
H
H
H N
H
HH N
H
H
H O
H
38
Hydrogen Bonding in Water
Hydrogen Bonds are strong!o Responsible for the high boiling point of watero Responsible for expansion of water as it freezes o Hydrogen bonding (dotted lines) between
water molecules in ice form tetrahedral configuration
Hydrogen Bonding in Water
0.957 Å1.97 Å
List all intermolecular forces for CH3CH2OH.
Hydrogen-bonds, dipole-dipole attractions, London dispersion forces
40
GroupProblem
41
Types of Intermolecular Forces
① London dispersion forces② Dipole-dipole forces③ Hydrogen bonds④ Ion-dipole forces
o Ion-induced dipole forces
42
Ion-Dipole Attractionso Attractions between ion and charged end of
polar moleculeso Ions have full charges, increasing the attraction
(a) Negative ends of water dipoles surround cation (b) Positive ends of water dipoles surround anion
43
AlCl3·6H2O
o Positive charge of Al3+ ion attracts partial negative charges – on O of water molecules
o Ion-dipole attractions hold water molecules to metal ion in hydrateo Water molecules are found
at vertices of octahedron around aluminum ion
Attractions between ion and polar molecules
44
Ion-Induced Dipole Attractionso Attractions between ion and dipole it induces on
neighboring moleculesoDepends on
oIon charge and oPolarizability of its neighbor
o Attractions can be quite strong as ion charge is constant, unlike instantaneous dipoles of London-dispersion forces
45
GroupProblem
How many water molecules would be attracted to this molecule by Ion-Dipole interactions?
46
GroupProblem
List the intermolecular forces and rank in order of strength for the liquids of each molecule.
o Ion-DipoleoHydrogen BondingoDipole-DipoleoLondon Forces
• Larger, longer, and therefore heavier molecules often have stronger intermolecular forces
• Smaller, more compact, lighter molecules have generally weaker intermolecular forces
Weakest
Strongest
GroupProblem
Intermolecular Forces and Temperature
Decrease with increasing temperatureo Increasing kinetic energy overcomes attractive
forceso If allowed to expand, increasing temperature
increases distance between gas particles and decreases attractive forces
48
49
GroupProblem
GROUP PROBLEM SET 12.1
50
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
More properties of gases,
liquids, & solids
Compressibility Surface Tension
Diffusion
Retention of Volume & shape
Wetting Viscosity
MeltingPoint
BoilingPoint
51
Melting & Boiling PointOften can predict physical properties by comparing strengths of intermolecular attractions:
Boiling Point increases when intermolecular forces increase
Melting Point increases when intermolecular forces increase
52
Compressibility
Measure of the ability of a substance to be forced into smaller volume
oDetermined by strength of intermolecular forcesoGases highly compressible
oMolecules far apartoWeak intermolecular forces
o Solids and liquids nearly incompressibleoMolecules very close togetheroStronger intermolecular forces
53
Retention of volume and shape
o Solids retain both volume and shapeoStrongest intermolecular attractionsoMolecules closest
o Liquids retain volume, but not shapeoAttractions intermediate
oGases, expand to fill their containersoWeakest intermolecular attractionsoMolecules farthest apart
54
DiffusionIn Gases
o Molecules travel long distances between collisions
o Diffusion rapidIn Liquids
o Molecules closero Encounter more collisionso Takes a long time to move
from place to placeIn Solids
o Diffusion close to zero at room temperature
o Will increase at high temperature
55
Surface Tension
Inside body of liquido Intermolecular forces are
the same in all directionsMolecules at surface
o Potential energy increases when removing neighbors
o Molecules move together to reduce surface area and potential energy
sphere
Why does H2O bead up on a freshly waxed car instead of forming a layer?
56
Surface Tension
Liquids containing molecules with strong intermolecular forces have high surface tension
Allows us to fill glass above rimoGives surface rounded
appearanceoSurface resists expansion and
pushes back
o Surface tension increases as intermolecular forces increase
o Surface tension decreases as temperature increases
57
Wettingo Ability of liquid to spread
across surface to form thin film
o Greater similarity in attractive forces between liquid and surface, yields greater wetting effect
o Occurs only if intermolecular attractive force between surface and liquid about as strong as within liquid itself
58
Wetting: Surfactants (Detergents)o Detergents added to water to lower surface tension so water
can spread on greasy glasso Substances that have both polar and non-polar characteristicso Long chain hydrocarbons with polar tail
OS
O
O Na+
O
O
O Na+
o Nonpolar end dissolves in nonpolar greaseo Polar end dissolves in polar H2Oo Thus increasing solubility of grease in water
59
Viscosity o Resistance to flowo Measure of fluid’s
resistance to flow or changing form
o Decreases as Temp increases
o Not just a property of liquids: o Gas: respond to instantly
to form changing forceo Amorphous solids, like
glass
60
Viscosity
Acetone Polar molecule
oDipole-dipole ando London forces
Ethylene glycolPolar molecule
oHydrogen-bondingoDipole-dipole and o London forces
Which is more viscous?
GroupProblem
61
GroupProblem
For each pair given, which is has more viscosity?
CH3CH2CH2CH2OH, CH3CH2CH2CHO
C6H14, C12H26
NH3(l ), PH3(l )
62
GroupProblem
GROUP PROBLEM SET 12.2
63
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Changes of State
Heating/CoolingCurves ΔH
Phase Diagrams
64
Phase Changes = changes of physical state with temperature ( α to KE)
SOLID LIQUID GASfusion
freezing
evaporation
condensation
deposition
sublimation
endothermic
exothermic
System absorbs energy from surrounds in the form of heato Requires the addition of heat
System releases energy into surrounds in the form of heat or lighto Requires heat to be decreased
65
Phase Changes of Water
ICE WATER VAPORmelting
freezing
evaporation
forming dew
deposition
sublimation
66
Phase ChangesE
nerg
y of
Sys
tem
Gas
Solid
Liquid
Meltingor Fusion
Vaporization Condensation
Freezing
SublimationDeposition
Exothermic, releases heat Endothermic, absorbs heat
Heating CurveIf heat added at constant rate
Horizontal lineso Phase changeso Melting pointo Boiling point
Diagonal lines o Heating of solid,
liquid or gas
o68
Cooling CurveHeat removed at constant rate
Diagonal lines o Cooling of
solid, liquid or gas
Horizontal lineso Phase changeso Melting pointo Boiling point
Supercooling o Temperature of liquid dips below its freezing point
69
Boiling Point (bp)
Bp increases as strength of intermolecular forces increase
Normal Boiling Point • T at which vapor pressure of liquid = 1 atm
70
Rate of Evaporationo Depends on
o Temperatureo Surface areao Strength of
intermolecular attractions
o Molecules that escape from liquid have larger than minimum escape KE
o When they leave the average KE of remaining molecules is less and so T lower
71
Effect of Temperature on Evaporation Rate
For given liquid the rate of evaporation per unit surface area increases as T increases
Why?o At higher T, total fraction
of molecules with KE large enough to escape is larger
o Result: rate of evaporation is larger
72
Kinetic Energy Distributionin 2 different liquids
o Smaller intermolecular forces
o Lower KE required to escape liquid
o A evaporates faster
o Larger intermolecular forces
o Higher KE required to escape liquid
o B evaporates slower
A B
73
GroupProblem
What is an example of gas A and gas B?
74
Effects of Hydrogen Bonding
• Boiling points of hydrogen compounds of elements of Groups 4A, 5A, 6A, and 7A.
• Boiling points of molecules with hydrogen bonding are much higher than expected
SOLID LIQUID GASfusion
freezing
evaporation
condensation
deposition
sublimation
Hfus Hvap
HsubMolar heat of fusion (Hfus)Heat absorbed by one mole of solid when it melts to give liquid at constantT and P
Molar heat of vaporization (Hvap )Heat absorbed when one mole of liquid is changed to one mole of vapor at constant T and P
Molar heat of sublimation (Hsub )Heat absorbed by one mole of solid when it sublimes to give one mole of vapor at constant T and P
Energies of Phase Changes
76
Measuring Hvap
o Clausius-Clapeyron equation o Measure pressure at various temperatures, then
plot
o Two point form of Clausius-Clapeyron equationo Measure pressure at two temperatures and solve
equation
CTR
HP vap
1ln
1221 11ln
TTRH
PP vap
77
Vapor Pressure Diagram
RT = 25 C
o Variation of vapor pressure with T
o Ether o Volatile o High vapor pressure
near RTo Propylene glycol
o Non-volatileo Low vapor pressure
near RT
Temp (K) Vapor P 1/T lnP280 32.4 0.003571429 3.478158423300 92.5 0.003333333 4.527208645320 225 0.003125 5.416100402330 334 0.003030303 5.811140993340 483 0.002941176 6.180016654
CTR
HP vap
1ln
Slope = Hvap/R = -4288.1KHvap = 8.3145 Jmol/K x 4288.1K Hvap = 35.65 x 103 J/molHvap = 35.65 kJ/mol
79
The vapor pressure of diethyl ether is 401 mm Hg at 18 °C, and its molar heat of vaporization is 26 kJ/mol. Calculate its vapor pressure at 32 °C.
1221 11ln
TTRH
PP vap
6109.04928.021 e
PP
216109.0 PP
T1 = 273.15 + 18 = 291.15 KT2 = 273.15 + 32 = 305.15 K
Determine the enthalpy of vaporization, in kJ/mol, for benzene, using the following vapor pressure data.
T = 60.6 °C; P = 400 torrT = 80.1 °C; P = 760 torr
A. 32.2 kJ/molB. 14.0 kJ/molC. –32.4 kJ/molD. 0.32 kJ/molE. –14.0 kJ/mol
80
GroupProblem
81
GroupProblem
82
Phase Diagrams• Show the effects of both pressure and temperature
on phase changes • Boundaries between phases indicate equilibrium• Triple point:
– The temperature and pressure at which s, l, and g are all at equilibrium
• Critical point: – The temperature and pressure at which a gas can no
longer be condensed– TC
= temperature at critical point– PC = pressure at critical point
Phase Diagram
X axis – temperatureY axis – pressure
o As P increases(T constant), solid most likely more compact
o As T increases (P constant), gas most likely higher energy
o Each point = T and Po B = o E =o F =
E
0.01 °C, 4.58 torr100 °C, 760 torr–10 °C, 2.15 torr
F
84
Phase Diagram of Water
AB = vapor pressure curve for iceBD = vapor pressure curve for liquid waterBC = melting point lineB = triple point: T and P where all three phases are in equilibriumD = critical point
T and P above which liquid does not exist
85
Phase Diagram – CO2
o Now line between solid and liquid slants to right
o More typicalo Where is triple
point?o Where is critical
point?
86
Supercritical Fluid
o Substance with temperature above its critical temperature (TC) and density near its liquid density
o Have unique properties that make them excellent solvents
o Values of TC tend to increase with increased intermolecular attractions between particles
87
At 89 °C and 760 mmHg, what physical state is present?
A.SolidB.LiquidC.GasD.Supercritical fluidE.Not enough information
is given
GroupProblem
88
GroupProblem
GROUP PROBLEM SET 12.3
89
The Before & After of Phase Changes
SOLID LIQUID GASfusion
freezing
evaporation
condensation
deposition
sublimation
endothermic
exothermic
System absorbs energy from surrounds in the form of heato Requires the addition of heat
System releases energy into surrounds in the form of heat or lighto Requires heat to be decreased
90
The molar heat of a phase change (H) describes the heat needed for a phase change to go to completion.
The specific heat of a phase change (q) describes the heat needed for an amount of a substance to completely undergo a phase change.
q = n x H
91
Enthalpy Of Phase ChangesEndothermic Phase Changes
1. Must add heat2. Energy entering system (+)
Sublimation: Hsub > 0Vaporization: Hvap > 0Melting or Fusion: Hfus > 0
Exothermic Phase Changes3. Must give off heat4. Energy leaving system (–)
Deposition: H < 0 = –Hsub Condensation: H < 0 = –Hvap Freezing: H < 0 = –Hfus
92
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Dynamic Equilibria
93
Equilibria Exist During a Phase Change
• Fraction of molecules in condensed state is higher when intermolecular attractions are higher
• Intermolecular attractions must be overcome to separate the particles, while separated particles are simultaneously attracted to one another
condensedphase
separatedphase
94
Le Chatelier’s Principle
o Equilibria are often disturbed or upseto When dynamic equilibrium of system is
upset by a disturbanceoSystem responds in direction that tends to
counteract disturbance and, if possible, restore equilibrium
o Position of equilibrium oUsed to refer to relative amounts of substance on
each side of double (equilibrium) arrows
95
Liquid Vapor Equilibrium Liquid + Heat Vapor
• Increasing T – Increases amount of vapor – Decreases amount of liquid
• Equilibrium has shifted – Shifted to the right– More vapor is produced at expense of liquid
• Temperature-pressure relationships can be represented using a phase diagram
96
Equilibrium & Phase Diagrams
T1 = 78°CP1 = 330 atm
To increaseT2 = 100°CThe system must respond by increasing P2 = 760 to restore equilibrium:o T is highero Volume of liquid
is lower o P of vapor higher
Le Chatelier’s Principle
Liquid + Heat Vapor
Initial V1 T1 P1
ChangeVolume lost
in evaporation
Increase Temperature
Pressure increases
Final V2 T2 P2
98
Evaporation Rate
99
Before System Reaches Equilibrium
o Liquid is placed in empty, closed, containero Begins to evaporate
o Once in gas phaseoMolecules can condense
by o Striking surface of liquid
and giving up some kinetic energy
100
System At Equilibrium
o Rate of evaporation = rate of condensation
o Occurs in closed systems where molecules cannot escape
101
Enthalpy Of Phase ChangesEndothermic:
Liquid+ heat of vaporization ↔ GasLiquid + Hvap ↔ Gas
Solid + heat of fusion ↔ LiquidSolid + Hfus ↔ Liquid
Solid + heat of sublimation ↔ GasSolid + Hsub ↔ Liquid
Exothermic:Liquid ↔ Gas - Hvap Solid ↔ Liquid - Hfus Solid ↔ Liquid - Hsub
102
liquid + heat of vaporization ↔ gasEquilibrium Vapor Pressure
o Pressure of gas when liquid or solid is at equilibrium with its gas phase
o Usually referred to as simply vapor pressureo Increasing temperature increases vapor pressure
because vaporization is endothermic
103
Vapor Pressure Diagram
RT = 25 C
• Variation of vapor pressure with T
• Ether – Volatile – High vapor pressure
near RT• Propylene glycol
– Non-volatile– Low vapor pressure
near RT
104
Effect of Volume on Vapor Pressure
Initial(equilibrium
exists)
Volume of Container
Volume of liquid P1
ChangeVolume
manually increased
Rate condensation
decreases
Pressure decreases
System changes to establish new
equilibrium
Volumeof container
greater
Volume of liquid
decreases
P2
(P2 = P1)
105
Similar Equilibria Reached in Melting
Melting Point (mp)o Solid begins to change
into liquid as heat added
Dynamic equilibria exists between solid and liquid states
oMelting (red arrows) and freezing (black arrows) occur at same rate
o As long as no heat added or removed from equilibrium mixture
106
Equilibria Reached in Sublimation
At equilibrium molecules sublime from solid at same rate as molecules condense from vapor
107
Do Solids Have Vapor Pressures?
o At given temperature some solid particles have enough KE to escape into vapor phase
o When vapor particles collide with surface they can be captured
o Yes equilibrium vapor pressure of solid exists
108
CHAPTER 12 Intermolecular Attractions & the Properties of Liquids & Solids
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E
Solid Structures
109
Types of Solids• Crystalline Solids
– Solids with highly regular arrangements of components
• Amorphous Solids– Solids with considerable disorder in their
structures
110
Crystalline Solids
• Unit Cell– Smallest
segment that repeats regularly
– Smallest repeating unit of lattice
– Two-dimensional unit cells
111
Crystal Structures Have Regular Patterns
• Lattice– Many repeats of unit cell – Regular, highly
symmetrical system– Three (3) dimensional
system of points designating positions of components• Atoms• Ions• Molecules
112
Three Types Of 3-D Unit Cells • Simple cubic
– Has one host atom at each corner– Edge length a = 2r – Where r is radius of atom or ion
• Body-centered cubic (BCC)– Has one atom at each corner and one in
center– Edge length
• Face-centered cubic (FCC)– Has one atom centered in each face, and one
at each corner– Edge length
Most efficient arrangement of spheres in two dimensions
Each sphere has 6 nearest neighbors Second layer with atoms in holes on the first
layer 113
Close Packing of Spheres1st layer 2nd layer
114
Two Ways to Put on Third Layer
1. Directly above spheres in first layer
2. Above holes in first layer
Remaining holes not covered by second layer
Cubic lattice: 3-dimensional arrays
115
3-D Simple Cubic Lattice
Portion of lattice—open view
Unit Cell
Space filling model
Other Cubic Lattices
116
Face Centered Cubic
Body Centered Cubic
117
Ionic Solids Lattices of alternating charges• Want cations next to anions
– Maximizes electrostatic attractive forces– Minimizes electrostatic repulsions
• Based on one of three basic lattices:– Simple cubic– Face centered cubic– Body centered cubic
Common Ionic SolidsRock salt or NaCl– Face centered cubic lattice of Cl– ions (green)– Na+ ions (blue) in all octahedral holes
118
119
Other Common Ionic Solids
Cesium Chloride, CsCl
Zinc Sulfide, ZnS
Calcium Fluoride, CaF2
120
Spaces In Ionic Solids Are Filled With Counter Ions
• In NaCl– Cl– ions form face-
centered cubic unit cell
– Smaller Na+ ions fill spaces between Cl–ions
• Count atoms in unit cell – Have 6 of each or
1:1 Na+:Cl– ratio
121
Counting Atoms per Unit Cell• Four types of sites in unit cell
– Central or body position – atom is completely contained in one unit cell
– Face site – atom on face shared by two unit cells– Edge site – atom on edge shared by four unit cells– Corner site – atom on corner shared by eight unit cells
Site Counts as Shared by X unit cellsBody 1 1Face 1/2 2Edge 1/4 4
Corner 1/8 8
122
Example: NaCl
Site # of Na+ # of Cl–
Body 1 0Face 0Edge 0
Corner 0Total 4 4
36 21 312 41
18 81
FaceEdge Corner
Center
123
1:1CsCl
Determine the number of each type of ion in the unit cell.
4:4ZnS
4:8CaF2
124
Some Factors Affecting Crystalline Structure
• Size of atoms or ions involved• Stoichiometry of salt• Materials involved
– Some substances do not form crystalline solids
125
Amorphous Solids (Glass)• Have little order, thus referred to as “super cooled
liquids”• Edges are not clean, but ragged due to the lack of order
126
X-Ray Crystallography
• X rays are passed through crystalline solid
• Some x rays are absorbed, most re-emitted in all directions
• Some emissions by atoms are in phase, others out of phase
• Emission is recorded on film
127
X-ray Diffraction
Experimental Setup Diffraction Pattern
128
Interpreting Diffraction Data• As x rays hit atoms
in lattice they are deflected
• Angles of deflections related to lattice spacing
• So we can estimate atomic and ionic radii from distance data
129
Interpreting Diffraction DataBragg Equation• nλ=2d sinθ
– n = integer (1, 2, …)– = wavelength of
X rays– d = interplane spacing
in crystal– = angle of incidence
and angle of reflectance of X rays to various crystal planes
130
Example: Diffraction DataThe diffraction pattern of copper metal was measured with X-ray radiation of wavelength of 131.5 pm. The first order (n = 1) Bragg diffraction peak was found at an angle θ of 50.5°. Calculate the spacing between the diffracting planes in the copper metal.
1(131.5 pm) = 2 × d × sin(50.5)
n = 2d sin
d = 283 pm
131
Example: Using Diffraction DataX-ray diffraction measurements reveal that copper crystallizes with a face-centered cubic lattice in which the unit cell length is 362 pm. What is the radius of a copper atom expressed in picometers?
This is basically a geometry problem.
132
Ex. Using Diffraction Data (cont.)
diagonal = 4 rCu = 512 pm
rCu = 128 pm
Pythagorean theorem: a2 + b2 = c2 Where a = b = 362 pm sides and c = diagonal
2a2 = c2 and aac 22 2
133
Ionic Crystals (e.g. NaCl, NaNO3)
• Have cations and anions at lattice sites• Are relatively hard• Have high melting points• Are brittle• Have strong attractive forces between ions • Do not conduct electricity in their solid
states• Conduct electricity well when molten
Potassium chloride crystallizes with the rock salt structure. When bathed in X rays, the layers of atoms corresponding to the surfaces of the unit cell produce a diffracted beam of X rays (λ=154 pm) at an angle of 6.97°. From this, calculate the density of potassium chloride in g/cm3.
134
GroupProblem
135
Covalent Crystals• Lattice positions occupied by atoms that are
covalently bonded to other atoms at neighboring lattice sites
• Also called network solids – Interlocking network of covalent bonds extending
all directions• Covalent crystals tend to
– Be very hard – Have very high melting points – Have strong attractions between covalently
bonded atoms
136
Ex. Covalent (Network) Solid • Diamond (all C)
– Shown
• SiO2 silicon oxide– Alternating Si and O– Basis of glass and quartz
• Silicon carbide (SiC)
137
Metallic Crystals• Simplest models
– Lattice positions of metallic crystal occupied by positive ions
– Cations surrounded by “cloud” of electrons
• Formed by valence electrons• Extends throughout entire solid
138
Metallic Crystals• Conduct heat and electricity
– By their movement, electrons transmit kinetic energy rapidly through solid
• Have the luster characteristically associated with metals– When light shines on metal– Loosely held electrons vibrate easily – Re-emit light with essentially same frequency
and intensity
139
GroupProblem
GROUP PROBLEM SET 12.3