Chapter 11 Intermolecular Forces. 2 11.1: Intermolecular Forces (IMF) IMF < intramolecular forces...

Chapter 11 Intermolecular Forces

211.1: Intermolecular Forces (IMF)

IMF < intramolecular forces (covalent, metallic, ionic bonds)

IMF strength: solids > liquids > gases Boiling points and melting points are good

indicators of relative IMF strength.

311.2: Types of IMF1. Electrostatic forces: act over larger

distances in accordance with Coulomb’s law

a. Ion-ion forces: strongest; found in ionic crystals (i.e. lattice energy)

2d

QQF

4b. Ion-dipole: between an ion and a dipole (a neutral, polar molecule/has separated partial charges) Increase with increasing polarity of

molecule and increasing ion charge.

2d

QQF

Cl-

S2-<

Ex: Compare IMF in Cl- (aq) and S2- (aq).

5c. Dipole-dipole: weakest electrostatic force; exist between neutral polar molecules Increase with increasing

polarity (dipole moment) of molecule

Ex: What IMF exist in NaCl (aq)?

6

d. Hydrogen bonds (or H-bonds): H is unique among the elements because it has a

single e- that is also a valence e-.– When this e- is “hogged” by a highly EN atom (a

very polar covalent bond), the H nucleus is partially exposed and becomes attracted to an e--rich atom nearby.

7

H-bonds form with H-X•••X', where X and X' have high EN and X' possesses a lone pair of e-

X = F, O, N (since most EN elements) on two molecules:

F-H

O-H

N-H

:F

:O

:N

8 * There is no strict cutoff for the ability to form H-bonds (S forms a biologically important hydrogen bond in proteins).

* Hold DNA strands together in double-helix

Nucleotide pairs form H-bonds

DNA double helix

9 H-bonds explain why ice is less dense than water.

10

Ex: Boiling points of nonmetal hydrides

Boili

ng P

oin

ts (

ºC)

Conclusions:

Polar molecules have higher BP than nonpolar molecules

∴ Polar molecules have stronger IMF

BP increases with increasing MW

∴ Heavier molecules have stronger IMF

NH3, H2O, and HF have unusually high BP.

∴ H-bonds are stronger than dipole-dipole IMF

112. Inductive forces: Arise from distortion of the e- cloud

induced by the electrical field produced by another particle or molecule nearby.

London dispersion: between polar or nonpolar molecules or atoms– * Proposed by Fritz London in 1930– Must exist because nonpolar molecules

form liquids

Fritz London(1900-1954)

12How they form:1. Motion of e- creates an instantaneous

dipole moment, making it “temporarily polar”.

2. Instantaneous dipole moment induces a dipole in an adjacent atom• * Persist for about 10-14 or 10-15 second

Ex: two He atoms

13* Geckos!

Geckos’ feet make use of London dispersion forces to climb almost anything. A gecko can hang on a glass

surface using only one toe.

Researchers at Stanford University recently developed a gecko-like robot which uses synthetic setae to climb walls

http://en.wikipedia.org/wiki/Van_der_Waals%27_force

14London dispersion forces increase with: Increasing MW, # of e-, and # of atoms (increasing

# of e- orbitals to be distorted)Boiling points:

Effect of MW: Effect of # atoms:pentane 36ºC Ne –246°C hexane 69ºC CH4 –162°Cheptane 98ºC

??? effect:H2O 100°C

D2O 101.4°C

“Longer” shapes (more likely to interact with other molecules)

C5H12 isomers: 2,2-dimethylpropane 10°C pentane

36°C

Summary of IMF

Van der Waals forces

16Ex: Identify all IMF present in a pure sample of each substance, then explain the boiling points.

BP(⁰C) IMF Explanation

HF 20

HCl -85

HBr -67

HI -35

Lowest MW/weakest London, but most

polar/strongest dipole-dipole and has H-bonds

Low MW/weak London, moderate polarity/dipole-

dipole and no H-bonds

Medium MW/medium London, moderate

polarity/dipole-dipole and no H-bonds

Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no H-bonds

London, dipole-dipole, H-bonds

London, dipole-dipole



1711.3: Properties resulting from IMF

1. Viscosity: resistance of a liquid to flow

Viscosity depends on:-the attractive forces between molecules-the tendency of molecules to become entangled-the temperature

1811.3: Properties resulting from IMF

• Surface tension: energy required to increase the surface area of a liquid

193. Cohesion: attraction of molecules for other molecules of the same compound

4. Adhesion: attraction of molecules for a surface

20

5. Meniscus: curved upper surface of a liquid in a container; a relative measure of adhesive and cohesive forcesEx:

Hg H2O(cohesion rules) (adhesion rules)

21

Phase ChangesPhase Changes• Surface molecules are only attracted inwards towards

the bulk molecules.• Sublimation: solid gas.• Vaporization: liquid gas.• Melting or fusion: solid liquid.• Deposition: gas solid.• Condensation: gas liquid.• Freezing: liquid solid.

Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes• Energy change of the system for the above processes

are:

22

Intermolecular Forces Bulk and SurfaceIntermolecular Forces Bulk and Surface

23

Phase ChangesPhase Changes

Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes– Sublimation: Hsub > 0 (endothermic).

– Vaporization: Hvap > 0 (endothermic).

– Melting or Fusion: Hfus > 0 (endothermic).

– Deposition: Hdep < 0 (exothermic).

– Condensation: Hcon < 0 (exothermic).

– Freezing: Hfre < 0 (exothermic).

• Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization:– it takes more energy to completely separate molecules, than

partially separate them.

24


Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes• All phase changes are possible under the right

conditions (e.g. water sublimes when snow disappears without forming puddles).

• The sequence

heat solid melt heat liquid boil heat gas

is endothermic.• The sequence

cool gas condense cool liquid freeze cool solid

is exothermic.

25


Energy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes

26


Heating CurvesHeating Curves• Plot of temperature change versus heat added is a

heating curve.• During a phase change, adding heat causes no

temperature change.– These points are used to calculate Hfus and Hvap.

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

• Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

27


Heating CurvesHeating Curves

28

Heating Curve IllustratedHeating Curve Illustrated

29


Critical Temperature and PressureCritical Temperature and Pressure• Gases liquefied by increasing pressure at some

temperature.• Critical temperature: the minimum temperature for

liquefaction of a gas using pressure.• Critical pressure: pressure required for liquefaction.

30

Critical Temperature, TCritical Temperature, Tcc

31

Transition to Supercritical COTransition to Supercritical CO22

32

Supercritical COSupercritical CO22 Used to Decaffeinate Coffee Used to Decaffeinate Coffee

33

Vapor PressureVapor Pressure

Explaining Vapor Pressure on the Molecular Explaining Vapor Pressure on the Molecular LevelLevel

• Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.

• These molecules move into the gas phase.• As the number of molecules in the gas phase

increases, some of the gas phase molecules strike the surface and return to the liquid.

• After some time the pressure of the gas will be constant at the vapor pressure.

34

Gas-Liquid EquilibrationGas-Liquid Equilibration

35


Explaining Vapor Pressure Explaining Vapor Pressure on the Molecular Levelon the Molecular Level

• Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.

• Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

36


Volatility, Vapor Pressure, and TemperatureVolatility, Vapor Pressure, and Temperature• If equilibrium is never established then the liquid

evaporates.• Volatile substances evaporate rapidly.• The higher the temperature, the higher the average

kinetic energy, the faster the liquid evaporates.

37

Liquid Evaporates when no Equilibrium is EstablishedLiquid Evaporates when no Equilibrium is Established

38


Volatility, Vapor Pressure, and TemperatureVolatility, Vapor Pressure, and Temperature

39

Vapor PressureVapor PressureVapor Pressure and Boiling PointVapor Pressure and Boiling Point• Liquids boil when the external pressure equals the

vapor pressure.• Temperature of boiling point increases as pressure

increases.• Two ways to get a liquid to boil: increase temperature

or decrease pressure.– Pressure cookers operate at high pressure. At high

pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required.

• Normal boiling point is the boiling point at 760 mmHg (1 atm).

40

Phase DiagramsPhase Diagrams• Phase diagram: plot of pressure vs. Temperature

summarizing all equilibria between phases.• Given a temperature and pressure, phase diagrams

tell us which phase will exist.• Features of a phase diagram:

– Triple point: temperature and pressure at which all three phases are in equilibrium.

– Vapor-pressure curve: generally as pressure increases, temperature increases.

– Critical point: critical temperature and pressure for the gas.

– Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid.

– Normal melting point: melting point at 1 atm.

41

Phase DiagramsPhase Diagrams• Any temperature and pressure combination not on a

curve represents a single phase.

42

Phase DiagramsPhase DiagramsThe Phase Diagrams of HThe Phase Diagrams of H22O and COO and CO22

• Water:– The melting point curve slopes to the left because ice is less

dense than water.

– Triple point occurs at 0.0098C and 4.58 mmHg.

– Normal melting (freezing) point is 0C.

– Normal boiling point is 100C.

– Critical point is 374C and 218 atm.

• Carbon Dioxide:– Triple point occurs at -56.4C and 5.11 atm.

– Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.)

– Critical point occurs at 31.1C and 73 atm.

43

Phase DiagramsPhase DiagramsThe Phase Diagrams of HThe Phase Diagrams of H22O and COO and CO22

4411.4: Phase ChangesProcesses: Endothermic: melting,

vaporization, sublimation Exothermic:

condensation, freezing, deposition

I2 (s) and (g)

Microchip

45

Water: Enthalpy diagram or heating curve

J/g) 334(mQ

TmQ )CJ/g 4.18(

TmQ )CJ/g 87.1(

TmQ )CJ/g 06.2(

J/g) 2602(mQ

TmcQ mHQ

4611.5: Vapor pressure

A liquid will boil when the vapor pressure equals the atmospheric pressure, at any T above the triple point.

Pressure cooker ≈ 2 atm

Normal BP = 1 atm

10,000’ elev ≈ 0.7 atm

29,029’ elev (Mt. Everest) ≈ 0.3 atm

4711.6: Phase diagrams: CO2

Lines: 2 phases exist in equilibrium

Triple point: all 3 phases exist together in equilibrium (X on graph)

Critical point, or critical temperature & pressure: highest T and P at which a liquid can exist (Z on graph)

For most substances, inc P will cause a gas to condense (or deposit), a liquid to freeze, and a solid to become more dense (to a limit.)

Temp (ºC)

48Phase diagrams: H2O

• For H2O, inc P will cause ice to melt.

5111.7-8: Structures of solids

Amorphous: without orderly structureEx: rubber, glass

Crystalline: repeating structure; have many different stacking patterns based on chemical formula, atomic or ionic sizes, and bonding

52Cubic Unit Cells in Crystalline Solids

• Primitive-cubic shared atoms are located only at each of the corners. 1 atom per unit cell.

• Body-centered cubic 1 atom in center and the corner atoms give a net of 2 atoms per unit cell.

• Face-centered cubic corner atoms plus half-atoms in each face give 4 atoms per unit cell.

53

Common Lattice Structures

Types of Crystalline Solids

Type Particles ForcesNotable

propertiesExample

s

Atomic AtomsLondon

dispersion

Poor conductors

Very low MP

Ar (s),Kr (s)

A small (~2 cm long) piece of rapidly melting argon ice (the liquid is flowing off at the bottom) which has been frozen by allowing a slow stream of the gas to flow into a small graduated cylinder which was immersed into a cup of liquid nitrogen

http://upload.wikimedia.org/wikipedia/commons/0/0d/Argon_ice_1.jpg

Molecular

crystals

Molecules

(polar or non-

polar)

London dispersion,

dipole-dipole, H-

bonds

Poor conductors

Low to moderate MP

SO2(s)

CO2 (s),

C12H22O11,

H2O (s)

Sucrose (liq at room

T)

Carbon dioxide, dry ice(g at room T)

Ice(liq at room T)

Covalent (a.k.a.

covalent network)

Atoms bonded in a covalent network

Covalent bonds

Very hardVery high MP

Generally insoluble

Variable conductivity

C (diamond

& graphite)

SiO2

(quartz)

Ge, Si, SiC, BN

DiamondGraphite SiO2

Ionic

Anions and

cationsCrystal

s shatter

!

Ion-ion (ionic bonding)

High Lattice Energy

Hard & brittle

High MP,BPPoor conductors

Some solubility in H2O

NaCl,Ca(NO3)2

Metallic

Metal cations in a diffuse, delocalized e- cloud

Metallic bondsUsually face-

centered or body

centered

Excellent conductors

MalleableDuctileHigh but wide range of MP

Cu, Al, Fe (hard)Alloys

Pb, Au, Na (soft)

59Overall• Physical properties depend on these forces. The

stronger the forces between the particles,

• (a) the higher the melting point.• (b) the higher the boiling point.• (c) the lower the vapor pressure (partial pressure of

vapor in equilibrium with liquid or solid in a closed container at a fixed temperature).

• (d) the higher the viscosity (resistance to flow).• (e) the greater the surface tension (resistance to an

increase in surface area).

60Practice

• Determine the type of solid and the forces holding the particles together

• SiO2 Covalent NetworkCovalent Bonds

• NaNO3 Ionic Electrostatic Att.

• C2H6 Molecular Dispersion

• CH3OH Molecular Dispersion, Dipole-Dipole, H-

Bond

• C(diamond)Covalent NetworkCovalent Bonds• Al Metallic Metallic• Kr Atomic (Molecular) Dispersion• H2O Molecular Dispersion, Dipole-Dipole, H-

Bond

61Extra Material

• The following pages contain some additional material and review items

6262

ExamplesExamples

6363

Ionic Solids Ionic Solids

stable, high melting pointsstable, high melting points held together by strong electrostatic forces held together by strong electrostatic forces

between oppositely charged ionsbetween oppositely charged ions larger ions are arranged in closest packing larger ions are arranged in closest packing

arrangement arrangement smaller ions fit in the holes created by the smaller ions fit in the holes created by the

larger ionslarger ions

Chapter 11-64

8–64

Cubic Unit Cells in Crystalline Solids

• Primitive-cubic shared atoms are located only at each of the corners. 1 atom per unit cell.

• Body-centered cubic 1 atom in center and the corner atoms give a net of 2 atoms per unit cell.

• Face-centered cubic corner atoms plus half-atoms in each face give 4 atoms per unit cell.

6565

Common Lattice Structures

Chapter 11-66

8–66

Calculations involving the Unit Cell

• The density of a metal can be calculated if we know the length of the side of a unit cell.

• The radius of an metal atom can be determined if the unit cell type and the density of the metal known

– Relationship between length of side and radius of atom:

• Primitive 2r = l; FCC: BCC

E.g. Polonium crystallizes according to the primitive cubic structure. Determine its density if the atomic radius is 167 pm.

E.g.2 Calculate the radius of potassium if its density is 0.8560 g/cm3 and it has a BCC crystal structure.

l42

r l43

r

Chapter 11-67

8–67

Figure 11.31

• Length of sides a, b, and c as well as angles vary to give most of the unit cells. Return to unit cells

Chapter 11-68

8–68

Unit Cells in Crystalline Solids

• Metal crystals made up of atoms in regular arrays – the smallest of repeating array of atoms is called the unit cell.

• There are 14 different unit cells that are observed which vary in terms of the angles between atoms some are 90°, but others are not. Go to Figure 11.31

Chapter 11-69

8–69

Packing of Spheres and the Structures of Metals

• Arrays of atoms act as if they are spheres. Two or more layers produce 3-D structure.

• Angles between groups of atoms can be 90° or can be in a more compact arrangement such as the hexagonal closest pack (see below) where the spheres form hexagons.

• Two cubic arrays one directly on top of the other produces simple cubic (primitive) structure.

– Each atom has 6 nearest neighbors (coordination number of 6); nearest neighbor is where an atom touches another atom.

– 54% of the space in a cube is used.• Offset layers produces a-b-a-b arrangement since it takes two layers to

define arrangement of atoms. – BCC structure an example. – Coordination # is 8.

Chapter 11-70

8–70

Packing of Spheres and the Structures of Metals

• FCC structure has a-b-c-a-b-c stacking. It takes three layers to establish the repeating pattern and has 4 atoms per unit cell and the coordination number is 12.

7171

Metallic CrystalsMetallic Crystals

can be viewed as metals can be viewed as metals atoms (spheres) packed atoms (spheres) packed together in the closest together in the closest arrangement possiblearrangement possible

closest packing- when each closest packing- when each sphere has 12 neighborssphere has 12 neighbors 6 on the same plane6 on the same plane 3 in the plane above3 in the plane above 3 in the plane below3 in the plane below

7272

Bonding of MetalsBonding of Metals the highest energy level for most the highest energy level for most

metal atoms does not contain many metal atoms does not contain many electronselectrons

these vacant overlapping orbitals these vacant overlapping orbitals allow outer electrons to roam freely allow outer electrons to roam freely around the entire metalaround the entire metal

7373

Bonding of MetalsBonding of Metals

these roaming electrons these roaming electrons

form a sea of electrons form a sea of electrons

around the metal atomsaround the metal atoms malleability and ductilitymalleability and ductility

bonding is the same in every directionbonding is the same in every direction one layer of atoms can slide past another one layer of atoms can slide past another

without frictionwithout friction conductivityconductivity

from the freedom of electrons to move from the freedom of electrons to move around the atomsaround the atoms

7474

Metal AlloysMetal Alloys

substance that is a mixture of elements substance that is a mixture of elements and has metallic propertiesand has metallic properties

substitutional alloysubstitutional alloy host metal atoms are replaced by other metal host metal atoms are replaced by other metal

atoms atoms happens when they have similar sizeshappens when they have similar sizes

interstitial alloyinterstitial alloy metal atoms occupy spaces created between metal atoms occupy spaces created between

host metal atomshost metal atoms happens when metal atoms have large happens when metal atoms have large

difference in sizedifference in size

7575

ExamplesExamples BrassBrass

substitutionalsubstitutional 1/3 of Cu atoms 1/3 of Cu atoms

replaced by Zn replaced by Zn SteelSteel

interstitialinterstitial Fe with C atoms in Fe with C atoms in

betweenbetween makes harder and makes harder and

less malleableless malleable

Chapter 11-76

8–76

Chapter 11 Overview

• Changes of State– Phase transitions– Phase Diagrams

• Liquid State– Properties of Liquids; Surface tension and viscosity– Intermolecular forces; explaining liquid properties

• Solid State– Classification of Solids by Type of Attraction between Units– Crystalline solids; crystal lattices and unit cells– Structures of some crystalline solids– Calculations Involving Unit-Cell Dimensions– Determining the Crystal Structure by X-ray Diffraction

Exam on Friday

We will begin Chp 14 Thursday

Chapter 11-77

8–77

Comparison of Gases, Liquids and Solids

– Gases are compressible fluids. Their molecules are widely separated.

– Liquids are relatively incompressible fluids. Their molecules are more tightly packed.

– Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move.

Figure 11.2 States of Matter

Chapter 11-78

8–78

Phase Transitions

• Melting: change of a solid to a liquid.• Freezing: change a liquid to a solid.

• Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called sublimation.

• Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called deposition.

H2O(s) H2O(l)

H2O(l) H2O(s)

H2O(l) H2O(g) or

H2O(s) H2O(g)

H2O(g) H2O(l) orH2O(g) H2O(s)

Chapter 11-79

8–79

Vapor Pressure

• In a sealed container, some of a liquid evaporates to establish a pressure in the vapor phase.

• Vapor pressure: partial pressure of the vapor over the liquid measured at equilibrium and at some temperature.

• Dynamic equilibrium

Chapter 11-80

8–80

Temperature Dependence of Vapor Pressures

• The vapor pressure above the liquid varies exponentially with changes in the temperature.

• The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

CT

1

RT

HPln vap

Chapter 11-81

8–81

Clausius – Clapeyron Equation

• A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R.

• Clausius – Clapeyron equation is true for any two pairs of points.

• Write the equation for each and combine to get:

21

vap

1

2

T

1

T

1

RT

H

P

Pln

CTRT

HPln vap

1

Chapter 11-82

8–82

Using the Clausius – Clapeyron Equation

• Boiling point the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere.

• Normal boiling point the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm).

E.g. Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of 190.0 mmHg at 25.0°C.E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it has a vapor pressure of 100.0 mmHg. What is the heat of vaporization?

Chapter 11-83

8–83

Energy of Heat and Phase Change

• Heat of vaporization: heat needed for the vaporization of a liquid.

H2O(l) H2O(g) H = 40.7 kJ

• Heat of fusion: heat needed for the melting of a solid.

H2O(s) H2O(l) H = 6.01 kJ

• Temperature does not change during the change from one phase to another.

E.g. Start with a solution consisting of 50.0 g of H2O(s) and 50.0 g of H2O(l) at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water.

Chapter 11-84

8–84

Phase Diagrams

• Graph of pressure-temperature relationship; describes when 1,2,3 or more phases are present and/or in equilibrium with each other.

• Lines indicate equilibrium state two phases.

• Triple point- Temp. and press. where all three phases co-exist in equilibrium.

• Critical temp.- Temp. where substance must always be gas, no matter what pressure.

• Critical pressurevapor pressure at critical temp.• Critical point- point where system is at its critical pressure and

temp.

Chapter 11-85

8–85

Properties of Liquids

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

• Viscosity: a measure of a liquid’s resistance to flow.

• Surface tension: The net pull toward the interior of the liquid makes the surface tend to as small a surface area as possible and a substance does not penetrate it easily.

• Viscosity: Related to mobility of a molecule (proportional to the size and types of interactions in the liquid). – Viscosity decreases as the temperature increases since increased

temperatures tend to cause increased mobility of the molecule.

Chapter 11-86

8–86

Intermolecular Forces

• Intermolecular forces: attractions and repulsions between molecules that hold them together.

• Intermolecular forces (van der Waals forces) hold molecules together in liquid and solid phases.– Ion-dipole force: interaction between an ion and partial charges in

a polar molecule.

– Dipole-dipole force: attractive force between polar molecules with positive end of one molecule is aligned with negative side of other.

– London dispersion Forces: interactions between instantaneously formed electric dipoles on neighboring polar or nonpolar molecules.

– Polarizability: ease with which electron cloud of some substance can be distorted by presence of some electric field (such as another dipolar substance). Related to size of atom or molecule. Small atoms and molecules less easily polarized.

Chapter 11-87

8–87

Boiling Points vs. Molecular Weight

• Hydrogen bonds: the interaction between hydrogen bound to an electronegative element (N, O, or F) and an electron pair from another electronegative element. Hydrogen bonding is the dominate force holding the two DNA molecules together to form the double helix configuration of DNA.

Chapter 11-88

8–88

Comparisonof Energies for Intermolecular Forces

Interaction Forces :Approximate Energy

Intermolecular

London 1 – 10 kJ

Dipole-dipole 3 – 4 kJ

Ion-dipole 5 – 50 kJ

Hydrogen bonding 10– 40 kJ

Chemical bonding

Ionic 100 – 1000 kJ

Covalent 100 – 1000 kJ

Chapter 11-89

8–89

Structure of Solids

• Types of solids:– Crystalline – a well defined arrangement of atoms; this

arrangement is often seen on a macroscopic level.• Ionic solids – ionic bonds hold the solids in a regular

three dimensional arrangement.• Molecular solid – solids like ice that are held together

by intermolecular forces.• Covalent network – a solid consists of atoms held

together in large networks or chains by covalent networks.

• Metallic – similar to covalent network except with metals. Provides high conductivity.

– Amorphous – atoms are randomly arranged. No order exists in the solid. Example: glass

Chapter 11 Intermolecular Forces. 2 11.1: Intermolecular Forces (IMF) IMF < intramolecular forces...

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