Gaya Antar Molekul

Chapter 12:
Intermolecular Attractions
and the Properties of Liquids
and Solids
Chemistry: The Molecular
Nature of Matter, 6E
Jespersen/Brady/Hyslop

Jespersen/Brady/Hyslop Chemistry: The Molecular Nature of Matter, 6E 2
Chapter 12 Intermolecular Forces
 Important differences between gases,
solids, and liquids:
 Gases
 Expand to fill their container
 Liquids
 Retain volume, but not shape
 Solids
 Retain volume and shape

Chapter 12 Intermolecular Forces
 Physical State of molecule depends
on
 Average kinetic energy of particles
 Recall KE ∝ Tave
 Intermolecular Forces
 Energy of Inter-particle attraction
 Physical Properties of Gases, Liquids
and Solids determined by
 How tightly molecules are packed together
 Strength of attractions between
molecules

 Converting gas → liquid or solid
 Molecules must get closer together
 Cool or compress
 Converting liquid or solid → gas
 Requires molecules to move farther apart
 Heat or reduce pressure
 As T↓, Kinetic Energy of molecules ↓
 At certain T, molecules don’t have
enough energy to break away from one
another’s attraction
Intermolecular Attractions

Inter vs. Intra-Molecular Forces
 Intramolecular forces
 Covalent bonds within molecule
 Strong
 ∆Hbond (HCl) = 431 kJ/mol
 Intermolecular forces
 Attraction forces between molecules
 Weak
 ∆Hvaporization (HCl) = 16 kJ/mol
Cl H Cl H
Covalent Bond (strong) Intermolecular attraction (weak)

Electronegativity Review
Electronegativity: Measure of attractive
force that one atom in a covalent bond has for
electrons of the bond

Bond Dipoles
 Two atoms with different electronegativity
values share electrons unequally
 Electron density is uneven
 Higher charge concentration around more
electronegative atom
 Bond dipoles
 Indicated with delta (δ) notation
 Indicates partial charge has arisen
H F
δ+
δ−

Net Dipoles
 Symmetrical molecules
 Even if they have polar bonds
 Are non-polar because bond dipoles cancel
 Asymmetrical molecules
 Are polar because bond dipoles do not cancel
 These molecules have permanent, net dipoles
 Molecular dipoles
 Cause molecules to interact
 Decreased distance between molecules increases
amount of interaction

Intermolecular Forces
 When substance melts or boils
 Intermolecular forces are broken
 Not covalent bonds
 Responsible for non-ideal behavior of
gases
 Responsible for existence of condensed
states of matter
 Responsible for bulk properties of matter
 Boiling Points and Melting Points
 Reflect strength of intermolecular forces

Three Important Types of
Intermolecular Forces
1. Dipole-dipole forces
 Hydrogen bonds
2. London dispersion forces
3. Ion-dipole forces
 Ion-induced dipole forces

I. Dipole-dipole Attractions
 Occur only between polar
molecules
 Possess dipole moments
 Molecules need to be
close together
 Polar molecules tend to
align their partial charges
 + to –
 As dipole moment ↑,
intermolecular force ↑
+ − + −
− + − +
+ − + −

I. Dipole-dipole Attractions
 Tumbling molecules
 Mixture of attractive and
repulsive dipole-dipole
forces
 Attractions (- -) greater
than repulsions(- -)
 Get net attraction
 ~ 1% of covalent bond

Dipole - Dipole Attractions
 Interactions between net dipoles in polar
molecules
 About 1% as strong as a covalent bond
 Decrease as molecular distance increases
 Drops off as 1/d3
(d = distance between
dipoles)
 Dipole-dipole forces ↑ with ↑ polarity

Hydrogen Bonds
 Special type of Dipole-Dipole Interaction
 Very strong dipole-dipole attraction
 ~40 kJ/mol
 Occurs between H and highly electronegative atom
(O, N, or F)
 H—F, H—O, and H—N bonds very polar
 e−
s lie closer to X than to H, so high partial charges
 H only has 1 e−
, so δ+
H presents almost bare proton
 δ−
X almostfull −1 charge
 Element’s small size, means high charge density
 Positive end of one can get very close to negative end of
another

Examples of Hydrogen Bonding
H 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
H
H N
H
H
H O
H

Effects of Hydrogen Bonding
 Boiling points of H
compounds of
elements of Groups
IVA, VA, VIA, and
VIIA.
 Boiling points of
molecules with H
bonding are higher
than expected.
 Don’t follow rule that
BP ↑ as MM ↑
(London forces ↑)
BoilingPoint(°C)

Hydrogen Bonding in Water
 Responsible for expansion of water as it freezes
 Hydrogen bonding produces strong attractions in
liquid
 Hydrogen bonding (dotted lines) between
water molecules in ice form tetrahedral configuration

II. London Dispersion Forces
 Intermolecular forces between
nonpolar molecules
 Two neutral molecules (atoms)
can affect each other
 Nucleus of 1 molecule (atom)
attracts e−
’s of adjacent molecule
(atom)
 Electron cloud distorts
 Temporary or instantaneous
dipole forms
 One instantaneous dipole can
induce another in adjacent molecule
(atom)
 Results in net attractive force
e−
e−
2+
e−
e−
Electrostatic
attraction
He atom 1 He atom 2
2+
δ−
δ−
δ+
δ+

London Forces
 When atoms near one another,
their valence electrons interact
 Repulsion causes electron
clouds in each to distort and
polarize
 Instantaneous, induced
dipoles result from this
distortion
 Effect enhanced with increased
particle mass
 Effect diminished by increased
distance between particles

London Dispersion Forces
 Instantaneous dipole-induced dipole
attractions
 London Dispersion Forces
 London forces
 Dispersion forces
 Decrease as 1/d6
(d = distance between
molecules)
 Operate between all molecules
 Neutral or net charged
 Nonpolar or polar

London Dispersion Forces
 Ease with which dipole moments can be
induced and thus London Dispersion
Forces depends on
1. Polarizability
2. Molecular size
 Number atoms
 Molecular mass
1. Molecular Shape

1. Polarizability
 Ease with which electron
distribution in neutral atom
(or molecule) can be distorted
 Larger molecules = more polarizable
 Larger number of e−
s ∝ greater ease of
distorting electron cloud
 Magnitude of resulting partial charge is larger
 London Forces ↑ as MM ↑
 More e−
, less tightly held
 London Forces ↑ as electron cloud volume ↑
 Depends on size of atoms

Jespersen/Brady/Hyslop Chemistry: The Molecular Nature of Matter, 6E
Your Turn!
List all intermolecular forces for CH3CH2OH.
A. H-bonding
B. H-bonding, Dipole-Dipole, London
C. Dipole-Dipole
D. London
E. London, H-bonding
23

Table 12.1 Boiling Points of
Halogens and Noble Gases
Larger molecules have stronger London forces
and thus higher boiling points.

2. Number of Atoms in Molecule
 London forces depend on number atoms in
molecule
 Boiling point of hydrocarbons demonstrates this
trend
Formula BP at 1 atm, °C Formula BP at 1 atm, °C
CH4 −161.5 C5H12 36.1
C2H6 −88.6 C6H14 68.7
C3H8 −42.1 : :
C4H10 −0.5 C22H46 327

How Intermolecular Forces
Determine Physical Properties
Propane, C3H8
BP –42.1o
C
Hexane, C6H14
 BP 68.7o
C
 More sites (marked with *) along its chain where
attraction to other molecules can occur

3. Molecular Shape
 Increased surface area available for contact
= Increased London Forces
 London dispersion forces between
spherical molecules are lower than
chain-like molecules
 More compact molecules
 H’s not as free to interact with H’s on other
molecules
 Less compact molecules
 H’s have more chance to interact with H’s
on other molecules

Physical Origin of Shape Effect
 Small area for
interaction
 Larger area for
interaction
More compact – lower BP Less compact – higher BP

Your Turn!
Which species has a higher boiling point, Cl2 or
HCl; F2 or HF ?
A. HCl; F2
B. Cl2; F2
C. HCl; HF
D. Cl2 ; HF
29

III. Ion-dipole Attractions
 Attractions between ion and charged end of
polar molecules
 Attractions can be quite strong as ions have full
charges
(a) Negative ends of water dipoles surround cation
(b) Positive ends of water dipoles surround anion

Ex. Ion-dipole Attractions
AlCl3·6H2O
 Positive charge of Al3+
ion
attracts partial negative
charges δ–
on O of water
molecules
 Ion-dipole attractions hold
water molecules to metal
ion in hydrate
 Water molecules are found
at vertices of octahedron
around aluminum ion
 Attractions between ion and polar molecules

Ion-induced Dipole Attractions
 Attractions between ion and dipole it induces
on neighboring molecules
 Depends on
 Ion charge and
 Polarizability of its neighbor
 Attractions can be quite strong as ion charge does
NOT flicker on and off like instantaneous dipoles of
ordinary London forces
 Ex. I–
and Benzene

Summary of Intermolecular
Attractions
Dipole-dipole
 occur between neutral molecules with permanent
dipoles;
 about 1% - 5% of covalent bond
 Mid range in terms of intermolecular forces
Hydrogen bonding
 Special type of dipole-dipole interaction
 Occur when molecules contain N—H,
H—F and O—H bonds
 About 5% to 10% of a covalent bond

Summary of Intermolecular
Attractions
London dispersion
 Present in all substances
 Weakest intermolecular forces
 Weak, but can lead to large net attractions
Ion-dipole
 Occur when ions interact with polar molecules
 Strongest intermolecular attraction
Ion-induced dipole
 Occur when ion induces dipole on neighboring particle
 Depend on ion charge and polarizability of its neighbor

Using Intermolecular Forces
 Often can predict physical properties (like BP
and MP) by comparing strengths of
intermolecular attractions
 Ion-Dipole
 Hydrogen Bonding
 Dipole-Dipole
 London Dispersion Forces
 Larger, longer, heavier molecules have stronger
IMFs
 Smaller, more compact, lighter molecules have
weaker IMFs
Weakest
Strongest

Learning Check
 Identify the kinds of intermolecular forces present in
the following compounds
 Rank them in order of increasing boiling point: H2S,
CH3OH, CBr4, and Ne
H
S
H
H
C
O H
H
H
Br
C
Br
Br
Br
Ne
dipole-
dipole
Hydrogen
bonding
London
forces
London
forces
MM=331.6
MM=20.2
CH3OH >H2S > CBr4 > Ne

Physical Properties that Depend on
How Tightly Molecules Pack
 Compressibility
 Measure of ability of substance to be forced into
smaller volume
 Determined by strength of intermolecular forces
 Gases highly compressible
 Molecules far apart
 Weak intermolecular forces
 Solids and liquids nearly incompressible
 Molecules very close together
 Stronger intermolecular forces

Intermolecular Forces Determine
Strength of Many Physical
Properties Retention of Volume and Shape
 Solids keep both volume and shape
 Strongest intermolecular attractions
 Molecules closest
 Gases, keep nothing
 Weakest intermolecular attractions
 Molecules farthest apart
 Liquids keep volume, but not shape
 Attractions intermediate

Diffusion
 Movement that
spreads one gas
though another gas to
occupy space
uniformly
 Spontaneous
intermingling of
molecules of one gas
with molecules of
another gas
Occurs more rapidly in gases than in liquids
Hardly at all in solids

Diffusion
 In Gases
 Molecules travel long
distances between
collisions
 Diffusion rapid
 In Liquids
 Molecules closer
 Encounter more
collisions
 Takes a long time to
move from place to
place
 In Solids
 Diffusion nonexistent at
room temp
 Will ↑ at high Temp

Surface Tension
 Why does H2O bead up
on a freshly waxed car
instead of forming a
layer?
 Inside body of liquid
 Intermolecular forces in
all directions
 Molecules at surface
 IMF only pull down and
to side
 Fewer attractions, so
free to expand in
direction with no forces

Surface Tension
 Tendency of liquid
to take shape that
minimizes surface
area
 Molecules at surface
have higher potential
energy than those in
bulk of liquid
 Energy required to
expand or increase
surface by unit area
 Wax = nonpolar
 H2O = polar
 Water beads as
wants to maximize
attractions

Surface Tension
 Liquids containing
molecules with strong
intermolecular forces have
high surface tension
 Allows us to fill glass above
rim
 Gives surface rounded
appearance
 Surface acts as “skin” that lets
water pile up
 Surface resists expansion and
pushes back
 Surface Tension ↑
as IMF ↑
 Surface Tension ↓
as IMF ↓

Wetting
 Ability of liquid to spread
across surface to form
thin film
 Greater similarity in
attractive forces
between liquid and
surface, yields greater
wetting effect
 Occurs only if
intermolecular attractive
force between surface
and liquid about as
strong as within liquid
itself

Wetting
Ex. H2O wets clean glass surface as it forms H
bonds to SiO2 surface
 Does not wet greasy glass, as grease interacts
weakly with water
 Only London dispersion forces
 Forms bead instead
Surfactants
 Added to detergents to lower surface tension of H2O
 Now water can get better access to surface to be
cleaned

Surfactants
 Substances that have polar and non-polar
characteristics
 Long chain hydrocarbons with polar tail
O
S
O
O−
Na+
O
O
O−
Na+
 Nonpolar end interacts with grease
 Polar end interacts with H2O
 Thus increasing solubility of grease in water

Viscosity
 Resistance to flow
 Measure of fluid’s
resistance to flow or
changing form
 Larger molecules
collide and interact
more often, impeding
their flow
 Also called internal friction
 Depends on intermolecular attractions
 Molecular shape
www.chemistryexplained.com

Viscosity
 Viscosity ↓ when Temperature ↑
 Most people associate liquids with viscosity
 Molasses more viscous than water
 Gases have viscosity
 Respond almost instantly to form-changing forces
 Solids, such as rocks
 Normally respond very slowly to forces acting to
change their shape
 For same size molecules, viscosity
increases as strength of Intermolecular
Forces increases

Effect of Intermolecular Forces on
Viscosity
Acetone
 Polar molecule
 Dipole-dipole and
 London forces
Ethylene glycol
 Polar molecule
 Hydrogen-bonding
 Dipole-dipole and
 London forces
Which is more viscous??

Your Turn!
For each pair given, which is more viscose ?
 CH3CH2CH2CH2OH, CH3CH2CH2CHO;
 C6H14, C12H26; NH3(l ), PH3(l )
A. CH3CH2CH2CH2OH; C6H14; NH3(l )
B. CH3CH2CH2CH2OH; C12H26; NH3(l )
C. CH3CH2CH2CHO; C6H14; PH3(l )
D. CH3CH2CH2CHO; C12H26; NH3(l )
E. CH3CH2CH2CH2OH; C12H26; PH3(l )
50

Solubility
 “Like dissolves like”
 To dissolve polar substance, use polar solvent
 To dissolve nonpolar substance, use nonpolar
solvent
 Compare relative polarity of two substances
 Similar polarity means greater ability to interact
with each other
 Differing polarity means that they don’t interact;
move past each other
 Surfactants
 Both polar and non-polar characteristics
 Used to increase solubility

Your Turn!
Which of the following are not expected
to be soluble in water?
A. HF
B. CH4
C. CH3OH
D. All are soluble

Phase Changes
 Changes of physical state
 Deal with motion of molecules
 As temperature changes
 Matter will undergo phase changes
 Liquid → Gas
 Evaporation
 As heat H2O, forms steam or water vapor
 Requires energy or source of heat to occur

Phase Changes
 Solid → Gas
 Sublimation
 Ice cubes in freezer, leave in long enough disappear
 Endothermic
 Gas → Liquid
 Cooling or Condensation
 Dew is H2O vapor condensing onto cooler ground
 Exothermic

Rate of Evaporation
 Depends on
 Temperature
 Surface area
 Strength of
intermolecular
attractions
 Molecules that escape
from liquid have larger
than average KE’s
 When they leave
 Average KE of
remaining
molecules is less
 T lower

Effect of Temperature on
Evaporation Rate
 For given liquid
 Rate of evaporation per
unit surface area ↑
as T ↑
 Why?
 At higher T, total
fraction of molecules
with KE large enough to
escape is larger
 Result: rate of
evaporation is larger

Kinetic Energy Distribution in Two
Different Liquids
 Smaller IMF’s
 Lower KE required to
escape liquid
 A evaporates faster
 Larger IMF’s
 Higher KE required
to escape liquid
 B evaporates
slower
A B

Changes Of State Involve Equilibria
 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
condensed
phase
separated
phase

Before System Reaches
Equilibrium
 Liquid is placed in
empty container
 Begins to evaporate
 Once in gas phase
 Molecules can
condense by
 Striking surface of liquid
and giving up some
kinetic energy

System At Equilibrium
 Rate of evaporation =
rate of condensation
 Can occur in system
where molecules are
constrained to remain
close to liquid surface

Similar Equilibria Reached in
Melting
Melting Point (mp)
 Solid begins to change
into liquid as heat added
 Dynamic Equilibria
exists between solid and
liquid states
 Melting (red arrows) and
freezing (black arrows)
occur at same rate
 As long as no heat added or
removed from equilibrium
mixture

Equilibria Reached in
Sublimation
At equilibrium
 Molecules evaporate
from solid at same
rate as molecules
condense from vapor
 Molecules sublime
and condense on
crystal at same rate

Phase ChangesEnergyofSystem
Gas
Solid
Liquid
Melting
or Fusion
Vaporization Condensation
Freezing
Sublimation
Deposition
↓ Exothermic, releases heat
↑ Endothermic, absorbs heat

Energy Changes Accompanying
Phase Changes
 All phase changes are possible under the
right conditions
 Following sequence is endothermic
heat solid → melt → heat liquid → boil → heat gas
 Following sequence is exothermic
cool gas → condense → cool liquid → freeze → cool
solid

∆H Accompanying Phase
ChangesEndothermic Phase Changes
1. Must Add Heat
2. Energy entering system (+)
Sublimation: ∆Hsub > 0
Vaporization: ∆Hvap > 0
Melting or Fusion: ∆Hfus > 0
Exothermic Phase Changes
1. Must Give Off Heat
2. Energy leaving system (–)
Deposition: ∆H < 0 = − ∆Hsub
Condensation: ∆H < 0 = − ∆Hvap
Freezing: ∆H < 0 = − ∆Hfus

Phase Changes
 As T changes, matter undergoes phase
changes
 Phase Change
 Transformation from one phase to another
 Liquid-Vapor Equilibrium
 Molecules in liquid
 Not in rigid lattice
 In constant motion
 Denser than gas, so more collisions
 Some have enough kinetic energy to
escape, some don’t

Liquid-Vapor Equilibrium
 At any given T,
 Average Kinetic
Energy of molecules
is constant
 But have distribution
KEs of particles
 Certain number of
molecules have
enough KE to
escape surface
 Process is Evaporation or Vaporization
 As T ↑, Avg. KE ↑ and number
molecules with enough KE to escape ↑
Kinetic Energy
Fractionofmolecules

Vapor Pressure (VP)
 Pressure molecules exert when they evaporate or
escape into gas (vapor) phase
 Pressure of gas when liquid or solid is at
equilibrium with its gas phase
 Increasing temperature increases vapor pressure
because conversion is endothermic
 liquid + heat of vaporization ↔ gas
Equilibrium Vapor Pressure
 VP once dynamic equilibrium reached
 Usually referred to as simply vapor pressure

Measuring Vapor Pressure
To measure pressures inside vessels, a manometer is
used.

Vapor Pressure Diagram
 Variation of
vapor pressure
with T
 Ether
 Volatile
 High vapor
pressure near RT
 Propylene
glycol
 Non-volatile
 Low vapor
pressure near
RT
RT = 25 °C

Effect of Volume on VP
A. Initial V
 Liquid – vapor
equilibrium exists
Α. ↑ V
 P ↓
 Rate of
condensation ↓
A. More liquid
evaporates
 New equilibrium
established

Measuring ∆Hvap
 Clausis-Clayperon Equation
 Measure P at various Ts, then plot
 Two point form of Clausis-Clayperon Equation
 Measure P at two T’s and solve equation
C
TR
H
P
vap
+






 ∆
−=
1
ln






−
∆
=
122
1 11
ln
TTR
H
P
P vap

Learning Check
The vapor pressure of diethyl ether is 401 mmHg at
18°C, and its molar heat of vaporization is 26 kJ/mol.
Calculate its vapor pressure at 32°C.
4928.0
15.291
1
15.305
1
)/(314.8
/106.2
ln
4
2
1
−=





−
⋅
×
=
KKmolKJ
molJ
P
P






−
∆
=
122
1 11
ln
TTR
H
P
P vap
6109.04928.0
2
1
== −
e
P
P
2
1
6109.0
P
P
=
mmHg
mmHg
P 2
2 106.6
6109.0
401
×==
T1 = 273.15 + 18 = 291.15K
T2 = 273.15 + 32 = 305.15K

Your Turn!
Determine the enthalpy of vaporization, in
kJ/mol, for benzene, using the following vapor
pressure data. T = 60.6 C; P = 400 torr
 T = 80.1 C; P = 760 torr
A. 32.2 kJ/mol
B. 14.0 kJ/mol
C. -32.4 kJ/mol
D. 0.32 kJ/mol
E. -14.0 kJ/mol
74

Your Turn! - Solution
75
1
2 2 1
1 1
ln
400 Hg 1 1
ln
J760 Hg 353.1 K 333.6 K
8.314
K mol
32,235 J/mol or 32.2 kJ/mol
vap
vap
vap
HP
P R T T
Hmm
mm
H
∆  
= − ÷
 
∆  
= − ÷
 
∆ =

Do Solids Have Vapor Pressures?
 Yes
 At given T
 Some solid particles have enough KE so escape
into vapor phase
 When vapor particles collide with surface
 they can be captured
 Equilibrium vapor pressure of solid
 P of vapor in equilibrium with solid

Boiling Point (bp)
 T at which vapor pressure of liquid =
atmospheric pressure.
 Bp ↑ as strength of IMF ↑
Normal boiling point
 T at which vapor pressure = 1 atm

Effects of Hydrogen Bonding
 Boiling points of H
compounds of
elements of Groups
IVA, VA, VIA, and VIIA.
 Boiling points of
molecules with H
bonding are higher than
expected
 Don’t follow rule that
BP ↑ as MM ↑ (London
forces ↑)

Your Turn!
Which of the following will affect the boiling
point of a substance?
A.Molecular mass of the material
B.Intermolecular attractions
C.The external pressure on the material
D.All of these
E.None of these

Heating Curve
 Heat added at constant rate
 Diagonal lines
Heating of solid, liquid or gas
 Horizontal lines
Phase changes
Melting point
Boiling point

Cooling Curve
 Heat removed at constant rate
 Diagonal lines
Cooling of solid,
liquid or gas
 Horizontal lines
Phase changes
Melting point
Boiling point
 Supercooling
T of liquid dips below its freezing point

Your Turn!
 How much heat, in J, is required to convert
10.00 g of ice at -10.00 o
C to water at
50.00 o
C ?
 Specific heat (J/g K): ice, 2.108, water, 1.487
 Enthalpy of fusion = 6.010 kJ/mol54
A. 5483 J
B. 5643 J
C. 2304 J
D. 2364 J
E. 62,400 J
82

Energies of Phase Changes
 Expressed per mole
 Molar heat of fusion (∆Hfus)
 heat absorbed by 1 mole of solid when it melts to give
liquid at same T and P
 Molar heat of vaporization (∆Hvap )
 heat absorbed when 1 mole of liquid is changed to 1 mole
of vapor at constant T and P
 Molar heat of sublimation (∆Hsub )
 Heat absorbed by 1 mole of solid when it sublimes to give 1
mole of vapor at constant T and P
 All of these quantities tend to ↑ with ↑ing
intermolecular forces

Le Chatelier’s Principle
 Equilibria are often disturbed or upset
 When dynamic equilibrium of system is upset
by a disturbance
 System responds in direction that tends to
counteract disturbance and, if possible, restore
equilibrium
 Position of equilibrium
 used to refer to relative amounts of substance on
each side of double (equilibrium) arrows

Liquid Vapor Equilibrium
Liquid + Heat  Vapor
 ↑ing T
 ↑s amount of vapor
 ↓s amount of liquid
 Equilibrium has shifted
 = right shift
 more vapor is produced at expense of liquid
 Temperature-pressure relationships can be
represented using a phase diagram

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 = temperature
 Y axis = pressure
 As P ↑ (T const), solid
most likely
 More compact
 As T ↑(P const), gas
most likely
 Higher energy
 Each point = T and P
 B =
 E =
 F =
E
0.01°C, 4.58 torr
100°C, 760 torr
–10°C, 2.15 torr
F

Phase Diagram of Water
 AB = vapor pressure
curve for ice
 BD = vapor pressure
curve for liquid water
 BC = melting point line
 B = triple point = T
where all 3 phases in
equilibrium
 D = critical point
 T and P above which
liquid does not exist

Case Study: An Ice Necklace
 A cube of ice may be
suspended on a string simply
by pressing the string into the
ice cube. As the string is
pressed onto the surface, it
becomes embedded into the
ice.
 Why does this happen?

Phase Diagram – CO2
 Now line
between solid
and liquid
slants to right
 More typical
 Where is Triple
point?
 Where is
Critical point?

Supercritical Fluid
 Substance with temperature above its critical
temperature (TC) and density near its liquid
density
 Have unique properties that make them
excellent solvents
 Values of TC tend to ↑ with increased
intermolecular attractions between particles

Your Turn!
At 89 °C and 760 mmHg,
what physical state is
present?
A.Solid
B.Liquid
C.Gas
D.Supercritical fluid
E.Not enough
information is given

Types of Solids
 Crystalline Solids
 Solids with highly regular arrangements of
components
 Amorphous Solids
 Solids with considerable disorder in their
structures

Crystalline Solids
 Unit Cell
 Smallest
segment that
repeats
regularly
 Smallest
repeating unit of
lattice
 2-Dimensional
Unit Cells

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

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 r22a =
3
r4
a =

Close Packing of Spheres
1st
layer 2nd
layer
 Most efficient arrangement of spheres in 2-D
 Each sphere has 6 nearest neighbors
 Square lattice: 2-dimensional arrays

Two Ways to Put on 3rd
Layer
1. Directly above
spheres in 1st
layer
2. Above holes in 1st
layer
 Remaining holes not
covered by 2nd
layer
Cubic lattice: 3-dimensional arrays

3-D Simple Cubic Lattice
Portion of lattice—
open view
Unit Cell
Space filling
model

Other Cubic Lattices
100
Face Centered
Cubic
Body Centered
Cubic

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 Solids
Rock salt or NaCl
 Face centered cubic lattice of Cl ions (green)
 Na+
ions (blue) in all octahedral holes
102

Other Common Ionic Solids
Cesium
Chloride,
CsCl
Zinc Sulfide,
ZnS
Calcium
Fluoride,
CaF2

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

Counting Atoms per Unit Cell
 4 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 cells
Body 1 1
Face 1/2 2
Edge 1/4 4
Corner 1/8 8

Example: NaCl
Site # of Na+
# of Cl−
Body 1 0
Face 0
Edge 0
Corner 0
Total 4 4
( ) 36 2
1 =×
( ) 312 4
1 =×
( ) 18 8
1 =×
Face
Edge Corner
Center

Learning Check:
1:1
CsCl
Determine the number of each type of ion in
the unit cell.
4:4
ZnS
4:8
CaF2

Some Factors Affecting
Crystalline Structure
 Size of atoms or ions involved
 Stoichiometry of salt
 Materials involved
 Some substances do not form crystalline solids

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

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

X-ray Diffraction
Experimental Setup Diffraction Pattern

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

Interpreting Diffraction Data
Bragg Equation
 nλ=2dsinθ
 n = integer (1, 2, …)
 λ = wavelength of
X–rays
 d = interplane
spacing in crystal
 θ = angle of
incidence and angle
of reflectance of
X–rays to various
crystal planes

Ex. 2 Diffraction Data
The diffraction pattern of copper metal was
measured with x-ray radiation of wavelength of
1.315 Å. The first order (n=1) Bragg diffraction
peak was found at an angle theta of 50.5
degrees. Calculate the spacing between the
diffracting planes in the copper metal.
1(1.315 Ǻ)=2×d×sin(50.5°)
nλ = 2dsinθ
d = 2.83 Ǻ

Ex 3. Using Diffraction data
X-ray diffraction measurements reveal that
copper crystallizes with a face-centered cubic
lattice in which the unit cell length is 3.62 Å.
What is the radius of a copper atom expressed
in angstroms and in picometers?
This is basically a geometry problem.

Ex.3 (cont)
A12.5)A62.3(2diagonal =×=
diagonal = 4 × rCu = 5.12 Å
rCu = 1.28 Å
Now convert to pm
Recall 1 Å = 1 × 10−10
m and 1 pm = 1 × 10−12
m
pm
m
pm
A
m
A 128
101
1101
28.1 12
10
=
×
×
×
× −
−
Pythagorean theorem: a2
+ b2
= c2
Where a = b = 3.62 Å sides and c = diagonal
2a2
= c2
and aac 22 2
==

Learning Check
Silver packs together in a faced center cubic
fashion. The interplanar distance, d,
corresponds to the length of a side of the unit
cell, and is 4.07 angstroms. What is the radius
of a silver atom?
ra 22=
r22A07.4 =
r = 0.536 Å
a

Ionic Crystals (ex. 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

Sample Homework Problem
 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
6.97º. From this, calculate the density
(g/cm3
).
119

Your Turn!
Yitterbium crystallizes with a face centered
cubic lattice. The atomic radius of Yitterbium
is 175 pm. Determine the unit cell length.
A. 495 pm
B. 700 pm
C. 350 pm
D. 990 pm
E. 247 pm
120

Your Turn! - Solution
121
diagonal of cube = 4 where = atomic radius
diagonal of cube = 2 a where a = side of cube
4 4 x 175 pm
a = 495 pm
2 2
r r
r
= =

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

Ex. Covalent (Network) Solid
 Diamond (all C)
 shown
 SiO2 silicon oxide
 Alternating Si and O
 Basis of glass and quartz
 Silicon carbide (SiC)

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

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

126
Learning Check:
Substance ionic molecular covalent metallic
X: pulverizes when struck;
non-conductive of heat
and electricity
Y: White crystalline solid
that conducts electrical
current when molten or
dissolved
Z: shiny, conductive,
malleable with high
melting temperature
Classify the following in terms of most likely type of solid.




Your Turn!
Molecular crystals can contain all of the listed
attraction forces except:
A. Dipole-dipole attractions
B. Electrostatic forces
C. London forces
D. Hydrogen bonding
127

Gaya Antar Molekul

Empfohlen

Empfohlen

Weitere ähnliche Inhalte

Was ist angesagt?

Was ist angesagt? (20)

Ähnlich wie Gaya Antar Molekul

Ähnlich wie Gaya Antar Molekul (20)

Kürzlich hochgeladen

Kürzlich hochgeladen (20)

Gaya Antar Molekul

Hinweis der Redaktion