Chapter 11 Jan12

88
1 Chapter 11 Chapter 11 Intermolecular Intermolecular Forces, Liquids, and Forces, Liquids, and Solids Solids

Transcript of Chapter 11 Jan12

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Chapter 11Chapter 11

Intermolecular Forces, Intermolecular Forces, Liquids, and SolidsLiquids, and Solids

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CONTENTS11.1 Introduction: A Molecular Comparison

11.2 Types of Intermolecular Forces :– 11.2.1 Ion-dipole forces– 11.2.2 Dipole-dipole forces– 11.2.3 Hydrogen bonding– 11.2.4 London dispersion forces– 11.2.5 Comparison between forces

11.3 Properties of Liquids– 11.3.1 Viscosity and surface tension

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11.4 Phase Change– 11.4.1 Energy changes accompanying

phase change– 11.4.2 Cooling curve– 11.4.3 Critical temperature and pressure

11.5 Vapour pressure

11.6 Phase diagram– 11.6.1 Phase diagram of CO2

– 11.6.2 Phase diagram of water

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11.7 Structure of solids– 11.7.1 Crystalline and amorphous

solids– 11.7.2 Unit cell - PC, FCC & BCC

structure

11.8 Bonding in Solids– 11.8.1 Molecular solids– 11.8.2 Covalent-network solids– 11.8.3 Ionic and metallic solids

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Learning Outcomes

Able to differentiate the 4 main intermolecular forces (IMF)

Able to relate IMF of a compound to the boiling point, surface tension, viscosity and vapor pressure

Able to use phase diagram in problem solvingAble to use knowledge on unit cell of an element

to calculate molar mass, density, radius etc.

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11.1 Introduction: A Molecular Comparison11.1 Introduction: A Molecular ComparisonSolid - Liquid - GasSolid - Liquid - Gas

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Solid - Liquid - GasSolid - Liquid - GasState Gas Liquid Solid

Volume Assumes vol. & shape of container

Assumes shape of the portion of

container but not volume

Retain its own shape & volume

Compressibility

Compressible Virtually incompressible

Virtually incompressible

Diffusion Occurs rapidly Occurs slowly

Extremely low

Flow Readily Readily

Does not flow

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11.2 Intermolecular Forces11.2 Intermolecular Forces

The attraction between molecules is an intermolecular force.

Intermolecular forces are much weaker than ionic or covalent bonds.

When a substance melts or boils, intermolecular forces are broken.

When a substance condenses, intermolecular forces are formed.

Boiling points/melting points indicate the strength of intermolecular forces.

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11.2 Types of Intermolecular Forces11.2 Types of Intermolecular Forces

Ion-dipole force

Dipole-dipole force

Hydrogen bonding

London-Dispersion force

+ +- -

+-+

A - H :B -

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11.2.1 Ion-Dipole Forces11.2.1 Ion-Dipole ForcesExists between ions and the partial charge

on the end of polar molecules/dipoleImportant for solutions of ionic compounds

in polar liquidsE.g.: NaCl in water Magnitude of attraction increases:

– charge of the ion increases– magnitude of dipole moment

increaseIonic bonding > Ion-dipoleIon-dipole > Dipole-dipole

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Ion-Dipole ForcesIon-Dipole Forces

+ -+

++

+ +

+ +

+-

- -

-

-

- -

-

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11.2.2 Dipole-dipole Forces11.2.2 Dipole-dipole ForcesExists between neutral polar moleculesThe partially positive end of one molecule attracts the partially negative end of another.

Weaker than ion-dipole forces

E.g.: CO---CO

CO

C C

C

C O

OO

O

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Dipole-dipole ForcesDipole-dipole Forces

+

+

+

+

+

+

++

-

- -

-

-

- -

-

Attractions are greater than repulsion, so the molecules feel a net attraction to each other

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Dipole-dipole ForcesDipole-dipole Forces

For molecules of approximately equal mass and size, strength of attraction increases with increasing polarity Mass

(amu) Dipole

moment () Boiling point

(K) CH3CH2CH3 44 0.1 231 CH3OCH3 46 1.3 248 CH3CHO 44 3.9 294

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Dipole-dipole ForcesDipole-dipole Forces

For molecules of comparable polarity, those with smaller molecular volumes generally experience higher dipole-dipole attractive forces.

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11.2.3 Hydrogen Bonding11.2.3 Hydrogen Bonding

Special type of intermolecular attraction.Hydrogen bonding is a special case of

dipole-dipole interactions.H-bonding requires:

– H bonded to a small electronegative element (most important for compounds of F, O, and N)

– An unshared electron pair on a nearby small electronegative ion or atoms (usually F, O, or N on another molecule)

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11.2.3 Hydrogen Bonding11.2.3 Hydrogen BondingE.g.:

Ion-dipole > Hydrogen bonding > Dipole-dipole

OH

H

OH

H

NH

H

OH

H

H

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Hydrogen BondingHydrogen Bonding

Hydrogen bond

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Hydrogen BondingHydrogen BondingHydrogen bonding is important in :Stabilising the structure of protein

– Folding of protein moleculesSurvival of aquatic in frozen lake

– Ice is less dense than water– Ice floats so forms an insulating layer on top

of lakes or river. Therefore, aquatic life can survive in winter.

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11.2.4 London Dispersion Forces11.2.4 London Dispersion Forces

What intermolecular forces exist between nonpolar molecules in liquid and solid state ?•Dipole-dipole attractions cannot exist between non-polar atoms or molecules•Non-polar molecules do not have permanent dipoles but all nonpolar substances can be liquefied•Therefore, there must exist some kind of attractive interactions between the particles

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Two schematic presentations of the instantaneous dipoles on two adjacent helium atoms, showing the electrostatic attraction between them

London Dispersion ForcesLondon Dispersion Forces

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London Dispersion ForcesLondon Dispersion Forces

In a collection of helium atoms, the average distribution of electrons about each nucleus is spherically symmetrical

The atoms are nonpolar and posses no permanent dipole moment

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London Dispersion ForcesLondon Dispersion ForcesIn any atom or molecule, electrons

constantly movingThe nucleus of one molecule (or atom)

attracts the electrons of the adjacent molecule (or atom).

For an instant, the electron clouds become distorted.

In that instant a dipole id formed (called an instantaneous dipole)

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London Dispersion ForcesLondon Dispersion ForcesOne instantaneous dipole can induce

another instantaneous dipole in ad adjacent molecule (or atom).

These two temporary dipoles attract each other.

The attraction is called London-dispersion force (LDF) , or simply a dispersion force.

LDF is the weakest of all intermolecular forces.LDF exist between all molecules (polar or

nonpolar)

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London Dispersion Forces (LDF)London Dispersion Forces (LDF)

What affects the strength of a dispersion force?– Molecules must be very close together for

these attractive forces to occur.– POLARIZABILITY is the ease with which an

electron distribution can be deformed.– The larger the molecule, the more polarize it is.– LDF increase as molecular weight increases.– LDF depend on the shape of the molecule.

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London Dispersion ForcesLondon Dispersion Forces

Factors affecting London dispersion force :1. Molecular weight

– increase in mol. wt. (increase in atomic radii) results in increase number of electrons

– the larger the molecule, the farther its electron from the nuclei, the greater its polarizability

– greater polarizability, the greater the strength of London dispersion force

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London Dispersion ForcesLondon Dispersion Forces

E.g. : Noble gases . As go down the periodic table, mol. wt. increase (atomic radii increase) , resulting greater boiling point

Substance Mol. Wt (amu) Boiling point (K)He 4.0 4.6Ne 20.2 27.3Ar 39.9 87.5Kr 83.8 120.9Xe 131.3 166.1

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London Dispersion ForcesLondon Dispersion ForcesFactors affecting London dispersion

force2. Shape of molecule

– greater the surface area available for contact, greater the dispersion force– e.g.: straight chain molecule >branched chain molecule

Neopentane(bp=282.7K)

N-pentane(bp=309.4K)

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11.2.5 Comparison between Forces11.2.5 Comparison between Forces

•LDF are found in all substances.•Their strength depends on molecular

shapes and molecular weights.•Dipole-dipole forces add to the effect of LDF.

•They are found only in polar molecules.•Ion-dipole interactions are stronger than

H-bonding.

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11.2.5 Comparison between Forces11.2.5 Comparison between Forces

•H-bonding is a special case of dipole-dipole interactions.•It is the strongest of intermolecular

forces involving neutral species.•H-bonding is most important for H

compounds of N, O, and F.

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•If ions are involved, ion-dipole (if a dipole is present) and ionic bonding are possible.•Ion-dipole interactions are stronger than

H-bonding.•Ordinary ionic or covalent bonds are

much stronger than these interactions!

11.2.5 Comparison between Forces11.2.5 Comparison between Forces

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11.3 Properties of Liquids11.3 Properties of Liquids

11.3.1 VISCOSITY– Viscosity - resistance of liquid to

flow.– Liquid flows by sliding molecules

over each other.– The stronger the intermolecular

forces, the higher the viscosity.– Viscosity usually decreases with an

increase in temperature.

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11.3.2 SURFACE TENSION– bulk molecules (those in liquid) are

equally attracted to their neighbours– surface molecules are only attracted

inwards to the bulk molecules– therefore surface molecules are

packed more closely than bulk molecules

Surface tension - amount of energy required to increase the surface area of a liquid by a unit amount.

Properties of LiquidsProperties of Liquids

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Explaining Surface Tension

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11_16

Surface

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SURFACE TENSION (CONT..)Strong intermolecular forces cause higher

surface tension.– Water has a high surface tension (H-

bonding)– Hg(l) has even higher surface tension

(there are very strong metallic bonds between Hg atoms)

Properties of LiquidsProperties of Liquids

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SURFACE TENSION (CONT..)Cohesive forces

– Intermolecular forces that bind molecules to one another

Adhesive forces– Intermolecular forces that bind

molecules to a surface

Properties of LiquidsProperties of Liquids

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Properties of LiquidsProperties of Liquids To illustrate this: Meniscus - the shape of the liquid surface

– if adhesive forces > than cohesive forces - •the liquid is attracted to its container, •meniscus is U-shaped (e.g. water in glass)

– if cohesive forces > than adhesive forces - •meniscus curved downwards (e.g. Hg in glass)

Capillary action – the rise of liquids up very narrow tubes.

The liquid climbs until adhesive and cohesive forces are balanced by gravity.

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Liquid Levels in Capillaries: Capillary Rise

Copyright © Houghton Mifflin Company. All rights reserved 11-19A

11_19a

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11.4 Phase Changes11.4 Phase Changes

A phase change occurs when the matter is transformed from one physical state to another

During phase change temperature remains constant– Sublimation : Solid Gas– Vaporization : Liquid Gas– Melting or Fusion : Solid Liquid– Deposition : Gas Solid– Condensation : Gas Liquid– Freezing : Liquid Solid

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During phase change, the energy change of the system are as given below– Sublimation HSUB > 0 (Endothermic)– Vaporisation HVAP > 0 (Endothermic)– Melting OR Fusion HFUS > 0 (Endothermic)– Deposition HDEP < 0 (Exothermic)– Condensation HCON < 0 (Exothermic)– Freezing HFREEZ < 0 (Exothermic)

11.4.1 Energy changes accompanying phase changes

Phase ChangesPhase Changes

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The amount of energy required to cause a phase change increases as the strength of the intermolecular forces increases

Generally heat of fusion is less than heat of vaporisation– Less energy is needed to allow particles to move

past one another than to separate them totally

Steam can cause severe burns, when it comes in contact with skin because condensation (exothermic)

11.4.1 Energy changes accompanying phase changes

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+

- -

11.4.2 Cooling curve11.4.2 Cooling curve

Cooling curve for the conversion of gaseous water to ice.

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+

- -

11.4.2 Cooling curve11.4.2 Cooling curve

Supercooling: When a liquid is cooled below its freezing point and it still remains a liquid.

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11.4.3 Critical Temperature and 11.4.3 Critical Temperature and PressurePressure

Gas can be liquefied by compressing them at a suitable T– As temperature increases, gas become more difficult to

liquefy because of the increasing kinetic energy

For every substance, there exist a temperature above which the gas cannot be liquefied, regardless of the pressure

CRITICAL TEMPERATURE (TC) - the highest temperature at which a substance can exist as a liquid (using pressure)

CRITICAL PRESSURE (PC) - the pressure required to bring about liquefaction at TC

The greater the intermolecular forces, the easier it is to liquefy a substance, thus the higher the TC of the substance.

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11.5 Vapour Pressure11.5 Vapour Pressure

What is vapor pressure ?– We place a quantity of ethanol ( b.p. 78.3°C @ 1 atm) in

an evacuated container

– some of the molecules on the surface of the liquid have enough energy to escape the attraction of the bulk liquid

– these molecules moves into the gaseous phase– as the number of molecules in the gas phase increases,

some of the gas phase molecules struck the surface and returns to the liquid

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Measurement of the Vapour Pressure of Water

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• It will reach a condition where the rate at which molecules return to the liquid is equal to the rate they escape from the liquid.

• At this condition ;– the number of molecules in the gas phase reaches a

steady state– pressure of the vapor remains constant– dynamic equilibrium is attained between the rate of

condensation and rate of evaporation– pressure of the vapor remains constant (vapor

pressure)

Vapour PressureVapour Pressure

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Vapour pressure of a liquid is defined as the pressure exerted by its vapor when the liquid and the vapor state are in dynamic equilibrium (Dynamic equilibrium is a condition in which two opposing processes occur simultaneously at equal rates.)

Vapour PressureVapour Pressure

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Liquids boil when the external pressure at the liquid surface equals the vapour pressure.

The normal boiling point is the boiling point at 760 mmHg (1 atm).

The temperature of the boiling point increases as the external pressure increases.

Two ways to get a liquid to boil:– Increase temperature or decrease pressure.

Vapour Pressure and Vapour Pressure and Boiling PointBoiling Point

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Under appropriate condition, equilibrium can exist between:– liquid - gas– solid - liquid– solid - gas

We can summarise these conditions as phase diagram

Phase diagram - a plot of pressure vs. temperature summarizing all equilibria between phases.

11.6 Phase Diagram11.6 Phase Diagram

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(Explanation for Figure) A-B = vapor pressure curve, represent equilibrium

between liquid and gas phase Vapor pressure curve ends at the critical point A = triple point (temp. & pressure at which all three

phases are in equilibrium A-D = change in melting point with increasing pressure The solid is denser than liquid. An increase in pressure

will result in more compact solid phase i.e. higher temperature are required to melt a solid in high pressure

Phase DiagramPhase Diagram

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11.6.1 Phase Diagrams for Carbon 11.6.1 Phase Diagrams for Carbon DioxideDioxide• Notice that the solid-

liquid equilibrium (melting point) line of CO2 follows the normal behaviour i.e. its melting point increases with increasing temperature.

• For solid CO2 to exist as liquid the pressure must exceed 5.11 atm

• CO2 does not melt, but sublimes at 1 atm

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11.6.2 Phase Diagram for Water11.6.2 Phase Diagram for Water

• The melting point of water decreases with increasing pressure

• Liquid form is more compact than its solid form

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Objectives:Distinguish between crystalline and

amorphous solidsDetermine the net contents in a cubic unit

cellcalculate the density of a substancePredict the types of solids

– Molecular solid– Covalent-Network solid– Ionic or metallic solid

11.7 Structure of Solids11.7 Structure of Solids

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11.7.1CRYSTALLINE SOLIDSWell-ordered ,definite arrangement of

molecules, atoms or ionsHave highly regular shape, repeating in

three dimensional patternsTend to melt at specific temperatureHave narrow range of intermolecular forcesE.g. : NaCl, Cu

Structure of SolidsStructure of Solids

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11.7.1AMORPHOUSMolecules, atoms or ions do not have

orderly arrangement– Lack of well-defined faces and shape

Tend to melt over a range of temperatureAmorphous solids have variable

intermolecular forcesE.g. : rubber, glass

Structure of SolidsStructure of Solids

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11.7.2 UNIT CELLSCrystalline solids have ordered, repeating

structure.The smallest repeating unit in a crystal is

called a UNIT CELL.Unit cell - the smallest unit with all the

symmetry of the entire crystal.The 3-D stacking of unit cells is the

crystal lattice CRYSTAL LATTICE.

Structure of SolidsStructure of Solids

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CubicTetragonalOrthorhombicRhombohedralMonoclinic

Triclinic

Hexagonal

Types of unit cellsTypes of unit cells

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There are 3 types of a cubic unit cellPrimitive cubic (PC)

– atoms are at the corners of a simple cube I.e. each atom shared by 8 unit cells

Cubic CellCubic Cell

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Cubic CellCubic CellBody-centred cubic (BCC)

– atoms at the corners of a cube plus one in the centre of the cube body

– corner atoms shared by 8 unit cells, centre atom completely enclosed in one unit cell

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Face-centred cubic (FCC) atoms at the corners of a cube plus one

atom in the centre of each cube face corner atoms shared by 8 unit cells, face

atoms shared by 2 unit cells

Cubic CellCubic Cell

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Fraction of an Atom at Fraction of an Atom at Various Positions in the Unit Various Positions in the Unit CellCell

Position Fraction

Centre 1

Face 1/2

Corner 1/8

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EXERCISE 1 - Determine the number of atoms per unit cell

How many atoms are there in the face-centred cubic unit cell?

SolutionFCC has 6 faces and 8 corners– 6 faces 1/2 atom = 3

atoms face

– 8 corners 1/8 atom = 1 atom cornerTotal = 4 atoms

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To determine the net content in a cubic unit cell – 1. calculate the atomic radius– 2. density

The relationship between the edge length (a) and radius (r) of atoms in – Simple/primitive cubic cell a = 2r– Body-centred cubic cell

– Face-centred cubic cell

Net content in a cubic unit Net content in a cubic unit cellcell

r34a

r 8r24a

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Exercise 2Exercise 2

Find the density (g/cm3) for Fe given that it has BCC (Body-centred Cubic) unit cell and the following data:

Edge of unit cell = 2.86 ÅFe molar mass = 55.85 g/molNA = 6.022 x 1023 atoms/mol

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Volume of Unit Cell = (2.86 Å )3 (1 cm/108 Å )3

= 2.34 10-23 cm3

mass/unit cell = 55.85g/mole 2 atoms

6.022 1023 atoms/mole = 1.85 10-22 g

Density = Mass/Volume = 7.91 g/cm3

Answer to Exercise 2Answer to Exercise 2

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Answer:Answer:

Density = 7.95 g/cm3

Volume = (2.86 10-10 m)3 = 2.34 10-29m3

Mass = density volume= 7.95 g 2.34 x 10-29m3 (100cm)3

cm3 1m3

= 1.86 10-22 g per unit cell

Exercise 3 Exercise 3 Using Fe as BCC, calculate the mass of one Fe atom given density 7.95 g/cm3 and edge, a is 2.86 Å.

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Because Fe is a body-centred cubic cell, there are two Fe atoms in the cell, therefore:

Mass of one Fe atom = (1.86 x 10-22g) / 2 = 9.30 x 10-23 g

Using the molar mass to calculate the mass of oneFe atom, you find the agreement is good:

Answer to Exercise 3 (cont)Answer to Exercise 3 (cont)

atomFegatomsFe

FemolFemolFeg /10277.9

1002.61

185.55 23

23

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Gold (Au) crystallizes in a cubic close-packed (CCP or FCC) and has a density of 19.3 g/cm3. Calculate the atomic radius of gold.

Answer

Step1each unit cell has 8 corners and 6 faces = 4 atomsm = 4atom 197.0 g Au 1mol of Au I unit cell 1 mol Au 6.022 X 1023

atom = 1.31 X 10-21 g/unit cell

Exercise 4Exercise 4

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Step2Volume can be calculated from density:V = mass/density = 1.31 10-21 g = 6.79 10-23

cm3

19.3 g/cm3 Convert to picometer, pm unit

AnswerAnswer

37

3

12

32323

1079.6

1011

11011079.6

pmV

mpm

cmmcmV

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Step3 Step 4a = V 1/3 a = 8 1/2 r

= (6.79 X 107 pm3)1/3 r = a/ 8 1/2

= 408 pm = 408 / 8 1/2

= 144 pm

AnswerAnswer

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Close Packing of Close Packing of SpheresSpheres

Molecular crystals are formed by close packing molecules

A crystal is built up by placing close packed layers of spheres on top of each other

The most efficient arrangement is a layer of equal sized spheres, each surrounded by 6 others in the layer

Solids have maximum intermolecular forces due to efficient packing

There is only one place for the second layers and two choices for the third layers

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There are 2 types of close packing (Based on third layer) Hexagonal Close Packing (HCP)

ABABHexagonal

Close Packing of SpheresClose Packing of Spheres

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Close Packing of SpheresClose Packing of Spheres

•Cubical Close Packing (CCP)

•ABCABC•Or also known as Face-centred Cubic (FCC)

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The number of particles immediately surrounding a particle in the crystal is known as Co-ordinate number (CN)

In the case of CCP and HCP, sphere has 12 equidistant nearest neighbor (6 in one plane and 3 each in above and below the plane. Therefore the CN is 12 for HCP and CCP)

Close Packing of SpheresClose Packing of Spheres

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The melting point and hardness depends on

arrangement of particlesattraction forces between them

We can classify solids according to types of forces (4 types)

11.8 Bonding in Solids11.8 Bonding in Solids

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Bonding in SolidsBonding in Solids Molecular (formed from atoms or molecules)

– usually soft with low melting point and conductivity

Covalent network (formed from atoms)– very hard with very high melting point and poor

conductor

Ionic (formed from ions)– hard, brittle, high melting point and poor conductivity

Metallic (formed from metal atoms)– soft or hard with high melting point, good

conductivity and ductile)

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Intermolecular forces

Dipole-dipole, London dispersion and H-bonds.

Weak intermolecular forces give rise to low melting point.

Efficient packing of molecules is important (since they are not regular spheres)

E.g. : Benzene (m.p 5°C & b.p 80°C) Toluene (m.p -95°C & b.p 111°C)

11.8.1 Molecular 11.8.1 Molecular SolidsSolids

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E.g. : Benzene VS Toluene

Benzene m.p 5°C, b.p 80°C Symmetrical planar

molecule packed efficiently in solid form - high melting point

IMF in liquid state lower than toluene

Toluene m.p -95°C, b.p 111°C Lower symmetry,

prevents from packing efficiently

Intermolecular forces depend on close contact - not effective

IMF in liquid state larger than benzene

Molecular SolidsMolecular Solids

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Atoms are held together in large networks or chain by covalent bonds

Covalent bonds are stronger than intermolecular forces E.g. Diamond & Graphite

– Diamond •Each carbon atom is bonded to 4 other carbon

atom– Graphite

•The carbon atoms are arranged in layers of inter-connected hexagonal rings

•The layers are held by dispersion forces•Readily slide past one another

11.8.2 Covalent-Network 11.8.2 Covalent-Network SolidsSolids

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Structures of Diamond and Graphite

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Ions are held together by ionic bonds The strength of ionic solids depends on the charges of

the ions – NaCl ---> Na+ Cl- ---> has m.p. of 801°C– MgO ----> Mg 2+ O 2- ----> has m.p. of 2852°C

The structure adopted by ionic solids depends largely on the charges and relative sizes of the ions

•E.g.: NaCl (Na+ has CN of 6 and surrounded by 6 Cl) whereas CsCl (Cs is surrounded by 8 Cl- atoms)

– increase in CN results in increase in structure

11.8.3 Ionic Solids11.8.3 Ionic Solids

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Metallic solids have metal atoms in the HCP and FCC or BCC arrangements

The bonding is due to valence electrons that are delocalized throughout the entire solid– the metal nuclei floats in a sea of

electrons

11.8.4 Metallic Solids11.8.4 Metallic Solids

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Metallic SolidsMetallic Solids Hardness and melting point depends on strength

of bonding – Bonding strength increases with increase in

bonding electrons•e.g. Na (one valence electron- m.pt= 97.5°C)•Cr (6 valence electrons - m.pt = 1890°c)

– Mobility of electron explains why electrons are good conductor of heat and electricity

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END of CHAPTER