Chapter 12: Intermolecular Attractions and the Properties of Liquids and Solids Chemistry: The...

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

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


3

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


4

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 decreases, kinetic energy of molecules decreases At certain T, molecules don’t have

enough energy to break away from one another’s attraction

Intermolecular Attractions


5

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)


6

Electronegativity ReviewElectronegativity: Measure of

attractive force that one atom in a covalent bond has for electrons of the bond


7

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


8

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


9

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


10

Three Important Types of Intermolecular Forces

1. London dispersion forces2. Dipole-dipole forces

Hydrogen bonds

3. Ion-dipole forces Ion-induced dipole forces


11

London Forces When atoms near one

another, their valence electrons interact

Repulsion causes electron clouds in each to distort and polarize

Instantaneous dipoles result from this distortion Effect enhanced with increased

volume of electron cloud size Effect diminished by increased

distance between particles and compact arrangement of atoms


12

London Forces Instantaneous dipole-induced dipole

attractions London Forces Dispersion forces

Operate between all molecules Neutral or net charged Nonpolar or polar


13

London Dispersion Forces

Ease with which dipole moments can be induced and thus London Forces depend on

1. Polarizability of electron cloud2. Points of attraction

Number atoms Molecular shape (compact or

elongated)


14

Polarizability

Ease with which the electron cloud can be distorted

Larger molecules often more polarizable Larger number of less tightly

held electrons Magnitude of resulting partial

charge is larger Larger electron cloud


15

Table 12.1 Boiling Points of Halogens and Noble Gases

Larger molecules have stronger London forces and thus higher boiling points.


16

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, CCH4 –161.5 C5H12 36.1

C2H6 –88.6 C6H14 68.7

C3H8 –42.1 : :

C4H10 –0.5 C22H46 327


Hexane, C6H14

BP 68.7 °C More sites (marked with *) along its chain where attraction to other molecules can occur

17

How Intermolecular Forces Determine Physical Properties

Propane, C3H8

BP –42.1 °C


18

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

Hydrogen atoms not as free to interact with hydrogen atoms on other molecules

Less compact molecules Hydrogen atoms have more chance to

interact with hydrogen atoms on other molecules


19

Physical Origin of Shape Effect Small area for

interaction Larger area

for interaction

More compact – lower BP Less compact – higher BP


20

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 Positive to negative

As dipole moment increases, intermolecular force increases

+ +

+ +

+ +


21

Dipole-Dipole Attractions Tumbling molecules

Mixture of attractive and repulsive dipole-dipole forces

Attractions (- -) are maintained longer than repulsions(- -)

Get net attraction ~1–4% of covalent

bond


22

Dipole-Dipole Attractions

Interactions between net dipoles in polar molecules

About 1–4% as strong as a covalent bond Decrease as molecular distance

increases Dipole-dipole forces increase with

increasing polarity


23

Hydrogen Bonds Special type of dipole-dipole Interaction

Very strong dipole-dipole attraction ~10% of a covalent bond

Occurs between H and highly electronegative atom (O, N, or F) H—F, H—O, and H—N bonds very polar

Electrons are drawn away from H, so high partial charges

H only has one electron, so +H presents almost bare

proton –

X almost full –1 charge Element’s small size, means high charge density Positive end of one can get very close to negative end

of another


24

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


25

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


Your Turn!List all intermolecular forces for

CH3CH2OH.

A. Hydrogen-bondsB. Hydrogen-bonds, dipole-dipole

attractions, London dispersion forcesC. Dipole-dipole attractionsD. London dispersion forcesE. London dispersion forces, dipole-dipole

attractions

26


Your Turn!In the liquid state, which species has the strongest intermolecular forces, CH4, Cl2, O2 or HF?

A. CH4

B. Cl2C. O2

D. HF

27


28

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


29

Ex. Ion-Dipole Attractions: AlCl3·6H2O

Positive charge of Al3+ ion attracts partial negative charges – 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


30

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 is constant, unlike instantaneous dipoles of ordinary London forces

E.g., I– and Benzene


31

Summary of Intermolecular Attractions

Dipole-dipole Occur between neutral molecules with

permanent dipoles About 1–4% of covalent bond Mid range in terms of intermolecular forcesHydrogen bonding

Special type of dipole-dipole interaction Occur when molecules contain N—H,

H—F and O—H bonds About 10% of a covalent bond


32

Summary of Intermolecular Attractions

London dispersion Present in all substances Weakest intermolecular force Weak, but can add up 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


33

Using Intermolecular Forces Often can predict physical properties (like

BP, MP and many others) by comparing strengths of intermolecular attractions Ion-Dipole Hydrogen Bonding Dipole-Dipole London Forces

Larger, longer, and therefore heavier molecules often have stronger intermolecular forces

Smaller, more compact, lighter molecules have generally weaker intermolecular forces

Weakest

Strongest


34

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


35

Intermolecular Forces Determine Strength of Many Physical

Properties Retention of volume and shape Solids retain both volume and shape

Strongest intermolecular attractions Molecules closest

Liquids retain volume, but not shape Attractions intermediate

Gases, expand to fill their containers Weakest intermolecular attractions Molecules farthest apart


Intermolecular Forces and Temperature

Decrease with increasing temperature Increasing kinetic energy overcomes

attractive forces If allowed to expand, increasing

temperature increases distance between gas particles and decreases attractive forces

36


37

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


38

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 close to zero

at room temperature Will increase at high

temperature


39

Surface Tension

Inside body of liquid Intermolecular forces are

the same in all directions Molecules at surface

Potential energy increases when removing neighbors

Molecules move together to reduce surface area and potential energy

Why does H2O bead up on a freshly waxed car instead of forming a layer?


40

Surface Tension Causes a liquid to

take the shape (a sphere) that minimizes its surface area Molecules at

surface have higher potential energy than those in bulk of liquid and move to reduce the potential energy

Wax = nonpolar H2O = polar Water beads in order

to reduce potential energy by reducing surface area


41

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 increases as intermolecular forces increase

Surface tension decreases as temperature increases


42

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


43

WettingEx. H2O wets clean glass surface as it

forms H–bonds to SiO2 surface

Does not wet greasy glass, because grease is nonpolar and water is very polar Only London forces Forms beads instead

Surfactants Added to detergents to lower surface tension

of H2O Now water can spread out on greasy glass


44

Surfactants (Detergents) Substances that have both polar and non-polar

characteristics Long chain hydrocarbons with polar tail

OS

O

O Na+

O

O

O Na+

Nonpolar end dissolves in nonpolar grease Polar end dissolves in polar H2O Thus increasing solubility of grease in water


45

Viscosity Resistance to flow Measure of fluid’s

resistance to flow or changing form

Related to intermolecular attractive forces

Also called internal friction Depends on intermolecular attractions

www.chemistryexplained.com


46

Viscosity Viscosity decreases when temperature

increases Most people associate liquids with

viscosity Syrup more viscous than water

Gases have viscosity Respond almost instantly to form-changing

forces Solids, such as rocks and glass have

viscosity Normally respond very slowly to forces

acting to change their shape


47

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 has more

viscosity?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 )48


49

Solubility “Like dissolves like”

To dissolve polar substance, use polar solvent

To dissolve nonpolar substance, use nonpolar solvent

Compare relative polarity Similar polarity means greater ability to

dissolve in each other Differing polarity means that they don’t

dissolve, they are insoluble Surfactants

Both polar and non-polar characteristics Used to increase solubility


50

Your Turn!Which of the following are not expected to be soluble in water?A. HFB. CH4

C. CH3OH

D. All are soluble


51

Phase Changes

Changes of physical state Deal with motion of molecules

As temperature changes Matter will undergo phase changes

Liquid Gas Evaporation, vaporization As heat is added, H2O, forms steam or

water vapor Requires energy or source of heat


52

Phase Changes Solid Gas

Sublimation Ice cubes in freezer, leave in long enough

disappear Endothermic

Gas Liquid Condensation Dew is H2O vapor condensing onto cooler

ground Exothermic Often limits lower night time temperature


53

Rate of Evaporation Depends on

Temperature Surface area Strength of

intermolecular attractions

Molecules that escape from liquid have larger than minimum escape KE

When they leave Average KE of

remaining molecules is less and so T lower


54

Effect of Temperature on Evaporation Rate

For given liquid Rate of evaporation

per unit surface area increases as T increases

Why? At higher T, total

fraction of molecules with KE large enough to escape is larger

Result: rate of evaporation is larger


55

Kinetic Energy Distribution in Two Different Liquids

Smaller intermolecular forces

Lower KE required to escape liquid

A evaporates faster

Larger intermolecular forces

Higher KE required to escape liquid

B evaporates slower

A B


56

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 anothercondensedphase

separatedphase


57

Before System Reaches Equilibrium

Liquid is placed in empty, closed, container Begins to evaporate

Once in gas phase Molecules can

condense by Striking surface of liquid

and giving up some kinetic energy


58

System At Equilibrium

Rate of evaporation = rate of condensation

Occurs in closed systems where molecules cannot escape


59

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


60

Equilibria Reached in Sublimation

At equilibrium Molecules sublime

from solid at same rate as molecules condense from vapor


61

Phase ChangesEn

erg

y o

f S

yst

em

Gas

Solid

Liquid

Meltingor Fusion

Vaporization Condensation

Freezing

SublimationDeposition

Exothermic, releases heat Endothermic, absorbs heat


62

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


63

Enthalpy Of Phase ChangesEndothermic Phase Changes

1. Must add heat2. Energy entering system (+)

Sublimation: Hsub > 0

Vaporization: Hvap > 0

Melting or Fusion: Hfus > 0

Exothermic Phase Changes1. Must give off heat2. Energy leaving system (–)

Deposition: H < 0 = –Hsub

Condensation: H < 0 = –Hvap

Freezing: H < 0 = –Hfus


64

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


65

Liquid-Vapor Equilibrium At any given T,

Average kinetic energy of molecules is constant

But particles have a distribution of kinetic energies

Certain number of molecules have enough KE to escape surface

As T increases, average KE increases and number molecules with enough KE to escape increases

Kinetic EnergyF

ract

ion

of m

olec

ules


66

Vapor Pressure 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 vaporization is endothermic liquid + heat of vaporization ↔ gas

Equilibrium Vapor Pressure VP once dynamic equilibrium reached Usually referred to as simply vapor pressure


67

Measuring Vapor Pressure

To measure pressures inside vessels, a manometer is used.


68

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


69

Effect of Volume on VPA.Initial V

Liquid – vapor equilibrium exists

B. Increase V Pressure

decreases Rate of

condensation decreases

C. More liquid evaporates

New equilibrium established


70

Measuring Hvap

Clausius-Clapeyron equation Measure pressure at various temperatures,

then plot

Two point form of Clausius-Clapeyron equation

Measure pressure at two temperatures and solve equation

CTR

HP vap

1ln

122

1 11ln

TTR

H

PP vap


71

Learning Check 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.

122

1 11ln

TTR

H

PP vap

6109.04928.0

2

1 ePP

21

6109.0P

P

T1 = 273.15 + 18 = 291.15 KT2 = 273.15 + 32 = 305.15 K


Your Turn!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

72


Your Turn! - Solution

73


74

Do Solids Have Vapor Pressures? Yes At given temperature

Some solid particles have enough KE to escape into vapor phase

When vapor particles collide with surface They can be captured

Equilibrium vapor pressure of solid Pressure of vapor in equilibrium with solid


75

Boiling Point (bp)

T at which vapor pressure of liquid = atmospheric pressure.

Bp increases as strength of intermolecular forces increase

Normal Boiling Point T at which vapor pressure of liquid = 1

atm


76

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


77

Your Turn!Which of the following will affect the boiling point of a substance?A.PolarizabilityB.Intermolecular attractionsC.The external pressure on the materialD.All of theseE.None of these


78

Heating Curve Heat added at constant rate

Diagonal lines Heating of solid, liquid or gas

Horizontal linesPhase changesMelting pointBoiling point


79

Cooling Curve Heat removed at constant rate

Diagonal lines Cooling of solid,

liquid or gas

Horizontal linesPhase changesMelting pointBoiling point

Supercooling Temperature 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 °C to water at 50.00 °C?Specific heat (J/g K): ice, 2.108, water, 4.184Enthalpy of fusion = 6.010 kJ/molA. 5483 JB. 5643 JC. 2304 JD. 2364 JE. 62,400 J

80


81

Energies of Phase Changes Expressed per mole Molar 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

All of these quantities tend to increase with increasing intermolecular forces


82

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


83

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


84

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


85

Phase Diagram

X axis – temperature

Y axis – pressure As P increases

(T constant), solid most likely More compact

As T increases (P constant), 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


86

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 and

P where all three phases are in equilibrium

D = critical point T and P above which

liquid does not exist


87

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?


88

Phase Diagram – CO2

Now line between solid and liquid slants to right

More typical Where is triple

point? Where is

critical point?


89

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 increase with increased intermolecular attractions between particles


90

Your Turn!

At 89 °C and 760 mmHg, what physical state is present?A.SolidB.LiquidC.GasD.Supercritical fluidE.Not enough information is given


91


92

Types of Solids Crystalline Solids

Solids with highly regular arrangements of components

Amorphous Solids Solids with considerable disorder in their

structures


93

Crystalline Solids

Unit Cell Smallest

segment that repeats regularly

Smallest repeating unit of lattice

Two-dimensional unit cells


94

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


95

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

layer96

Close Packing of Spheres

1st layer 2nd layer


97

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


98

3-D Simple Cubic Lattice

Portion of lattice—open view

Unit Cell

Space filling model


Other Cubic Lattices

99

Face Centered

Cubic

Body Centered

Cubic


100

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

101


102

Other Common Ionic Solids

Cesium Chloride,

CsCl

Zinc Sulfide,

ZnS

Calcium Fluoride,

CaF2


103

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


104

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 cells

Body 1 1

Face 1/2 2

Edge 1/4 4

Corner 1/8 8


105

Example: NaCl

Site # of Na+ # of Cl–

Body 1 0

Face 0

Edge 0

Corner 0

Total 4 4

36 21

312 41

18 81

FaceEdge Corner

Center


106

Learning Check:

1:1CsCl

Determine the number of each type of ion in the unit cell.

4:4ZnS

4:8CaF2


107

Some Factors Affecting Crystalline Structure

Size of atoms or ions involved Stoichiometry of salt Materials involved

Some substances do not form crystalline solids


108

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


109

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


110

X-ray Diffraction

Experimental Setup Diffraction Pattern


111

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


112

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


113

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


114

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.


115

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


116

Learning CheckSilver 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 407 pm. What is the radius of a silver atom?

ra 22

r = 53.6 pm

a


117

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


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 6.97°. From this, calculate the density of potassium chloride in g/cm3.

118


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 pmB. 700 pmC. 350 pmD. 990 pmE. 247 pm

119


Your Turn! - Solution

120

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


121

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


122

Ex. Covalent (Network) Solid

Diamond (all C) Shown

SiO2 silicon oxide Alternating Si and O Basis of glass and quartz

Silicon carbide (SiC)


123

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


124

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


125

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 attractionsB. Electrostatic forcesC. London forcesD. Hydrogen bonding

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Chapter 12: Intermolecular Attractions and the Properties of Liquids and Solids Chemistry: The...

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