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Intermolecular Forces. (a) Particles in solid (b) Particles in liquid (c) Particles in gas. Intermolecular Forces. Objectives. Distinguish various properties of liquids and solids. Relate different intermolecular forces to explain observations in lab and nature.

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intermolecular forces
Intermolecular Forces

(a) Particles in solid (b) Particles in liquid (c) Particles in gas

slide2

Intermolecular Forces

Objectives

  • Distinguish various properties of liquids and solids.
  • Relate different intermolecular forces to explain observations in lab and nature.
  • Use heating curves to calculate energy transfers of substances across various phase changes (melting, freezing, boiling, condensing, sublimation, deposition).
  • Use phase diagrams to relate pressure and temperature changes to physical states.
properties of solids liquids and gases
Properties of Solids, Liquids, and Gases

Property Solid Liquid Gas

Shape Has definite shape Takes the shape of Takes the shape

the container of its container

Volume Has a definite volume Has a definite volume Fills the volume of

High densities High densities the container

Low densities

Bonding Ionic, Metallic, Covalent Covalent Covalent

Arrangement of Fixed, very close Random, close Random, far apart,

Particles Crystalline or amorphous Collisions

Interactions BetweenVery strong forces: Strong forces: Essentially none

Particles (i.e. Melting point, (i.e. Boiling point,

malleability, ductility, Surface Tension,

conductivity…) Viscosity, Vapor

pressure…)

liquid
Liquid

Surface Tension

  • molecules minimize
  • their surface area (“skin”)
  • molecules at surface
  • interact only with molecules
  • in the interior of liquid
  • surface molecules
  • subjected to inward force,
  • so surface is under tension
  • surface tension increases
  • with increasing intermolecular
  • forces

H2O(l) Water

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 31

liquid5
Liquid

Viscosity

  • resistance of a liquid to flow
  • greatest in substances with
  • strong intermolecular forces,
  • which hinder flow
  • longer molecules higher
  • viscosity than shorter ones

H2O(l) Water

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 31

intermolecular forces dispersion london force

-

+

+

-

-

+

-

+

-

+

+

-

-

+

+

-

Repulsion

Attraction

Intermolecular ForcesDispersion (London Force)
  • (a) Interaction of any two

atoms or molecules.

Electrons unevenly

distributed. Creates

instantaneous (temporary

dipole). Polarization

increases with size.

  • (b) interaction of many

dipoles. WEAKEST forces!

- +

- +

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 442

intermolecular forces dipole dipole

-

+

+

-

-

+

-

+

-

+

+

-

-

+

+

-

Repulsion

Attraction

Intermolecular Forces Dipole-Dipole
  • (a) Interaction of two

polar molecules. Polar

molecules have permanent

dipoles from electronegativity

difference. Higher melting and

boiling points due to

stronger IM forces.

  • (b) interaction of many

dipoles in a liquid.

- +

- +

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 442

intermolecular forces hydrogen bonding

H

H

O

O

H

H

Intermolecular Forces Hydrogen Bonding
  • Strong intermolecular forces of attraction between molecules containing fluorine, oxygen, or nitrogen bonded to hydrogen
  • Results from large

electronegativity

difference and small

atomic size of hydrogen

Hydrogen

bonds

Chemical

Bonds

Chemical

Bonds

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 442

molecular structure of ice
Molecular Structure of Ice

Hydrogen

bonding

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 455

hot vs cold tea
Hot vs. Cold Tea

Low temperature

(iced tea)

Many molecules have an

intermediate kinetic energy

High temperature

(hot tea)

Few molecules have a

very high kinetic energy

Percent of molecules

Kinetic energy

slide17

Evaporation (Vaporization)

  • Molecules must have sufficient energy to break IM bonds.
  • Molecules at the surface break away and become gas (“volatility”).
  • Only molecules with enough KE escape.
  • Breaking IM bonds absorbs energy.

Evaporation is endothermic.

  • Rate of evaporation increases with

increasing surface area, increasing

temperature, and weaker IM forces

condensation
Condensation
  • Forming IM bonds from gas to liquid
  • Condensation is exothermic because energy is released.
  • Dynamic equilibrium: rate of vaporization

equals rate of condensation (gas molecules above liquid becomes constant).

  • Vapor pressure: partial pressure

of gas in dynamic equilibrium with liquid

  • Vapor pressure increases

with increasing temperature

and weaker IM forces

boiling point temperature at which the vapor pressure of a liquid is equal to the pressure above it
Boiling point: temperatureat which the vapor pressure of a liquid is equal to the pressure above it

Microscopic view of a liquid near its surface

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 446

formation of a bubble is opposed by the pressure of the atmosphere
Formation of a bubble is opposed by the pressure of the atmosphere

Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 452

heating curves
Heating Curves
  • Temperature Change
    • change in KE (molecular motion)
    • depends on heat capacity
  • Phase Change
    • change in PE (molecular arrangement)
    • temperature remains constant
  • Heat Capacity
    • energy required to raise the temp of 1 gram of a substance by 1°C (q = mC∆T)

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

heating curve for water
Heating Curve for Water

vaporization

E

gas

D

100

condensation

C

liquid

melting

Temperature (oC)

B

0

A

freezing

solid

Heat added

LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

heating curves24

Gas - KE

Boiling - PE 

Liquid - KE 

Melting - PE 

Solid - KE 

Heating Curves

140

120

100

80

60

40

Temperature (oC)

20

0

-20

-40

-60

-80

-100

Energy

heating curves25
Heating Curves
  • Heat of Fusion (Hfus)
    • energy required to melt 1 gram of a substance at its melting point. Breaking intermolecular forces in the solid. Water = 6.02 kJ/mol
  • Heat of Vaporization (Hvap)
    • energy required to vaporize 1 gram of a substance at its boiling point.
    • usually larger than Hfus (requires complete separation of molecules)
    • higher temperatures = lower (Hvap)

Water = 40.7 kJ/mol (@ B.P.)

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

calculating energy changes heating curve for water
Calculating Energy Changes - Heating Curve for Water

140

DH = mol xDHvap

DH = mol xDHfus

120

100

80

Heat = mass xDt x Cp, gas

60

40

Temperature (oC)

20

Heat = mass xDt x Cp, liquid

0

-20

-40

-60

Heat = mass xDt x Cp, solid

-80

-100

Energy

specific heat capacities
Specific Heat Capacities

Tro's "Introductory Chemistry", Chapter 3

heating curves29
Heating Curves
  • Heat of Vaporization (Hvap)
    • energy required to vaporize 1 gram of a substance at its boiling point.
    • usually larger than Hfus (requires complete separation of molecules)
    • higher temperatures = lower (Hvap)
  • EX: sweating, steam burns, the drinking bird

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

energy changes accompanying phase changes
Energy Changes Accompanying Phase Changes

Gas

Vaporization

Condensation

Sublimation

Deposition

Energy of system

Liquid

Melting

Freezing

Solid

Brown, LeMay, Bursten, Chemistry2000, page 405

phase diagrams
Phase Diagrams
  • Show the phases of a substance at different temps and pressures.

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

slide32

Intermolecular Forces

  • Determine the kinds of IM forces that are present in the following: Kr, N2, CO, HF, NCl3, NH3, SiH4, CCl4
  • Which of the following has the highest boiling point. Why? CH4, CH3CH3, CH3CH2CH3, CH3CH2CH2CH3
  • A flask containing a mixture of NH3(g) and CH4(g) is cooled. At -33.30C a liquid begins to form in the flask. What is the liquid?
boiling points of covalent hydrides

GeH4

Boiling Points of Covalent Hydrides

H2O

100

0

H2Te

Boiling point (oC)

H2Se

SnH4

H2S

-100

SiH4

CH4

-200

50

100

150

Molecular mass

melting points
Melting Points

H

-259.2

He

-269.7

Mg

650

1

1

Symbol

Melting point oC

Li

180.5

Be

1283

B

2027

C

4100

N

-210.1

O

-218.8

F

-219.6

Ne

-248.6

2

2

> 3000 oC

2000 - 3000 oC

Na

98

Mg

650

Al

660

Si

1423

P

44.2

S

119

Cl

-101

Ar

-189.6

3

3

K

63.2

Ca

850

Sc

1423

Ti

1677

V

1917

Cr

1900

Mn

1244

Fe

1539

Co

1495

Ni

1455

Cu

1083

Zn

420

Ga

29.78

Ge

960

As

817

Se

217.4

Br

-7.2

Kr

-157.2

4

4

Rb

38.8

Sr

770

Y

1500

Zr

1852

Nb

2487

Mo

2610

Tc

2127

Ru

2427

Rh

1966

Pd

1550

Ag

961

Cd

321

In

156.2

Sn

231.9

Sb

630.5

Te

450

I

113.6

Xe

-111.9

5

5

Cs

28.6

Ba

710

La

920

Hf

2222

Ta

2997

W

3380

Re

3180

Os

2727

Ir

2454

Pt

1769

Au

1063

Hg

-38.9

Tl

303.6

Pb

327.4

Bi

271.3

Po

254

At

Rn

-71

6

6

Ralph A. Burns, Fundamentals of Chemistry , 1999, page 1999

densities of elements
Densities of Elements

H

0.071

He

0.126

1

1

Li

0.53

Be

1.8

B

2.5

C

2.26

N

0.81

O

1.14

F

1.11

Ne

1.204

2

2

Na

0.97

Mg

1.74

Al

2.70

Si

2.4

P

1.82w

S

2.07

Cl

1.557

Ar

1.402

3

3

K

0.86

Ca

1.55

Sc

(2.5)

Ti

4.5

V

5.96

Cr

7.1

Mn

7.4

Fe

7.86

Co

8.9

Ni

8.90

Cu

8.92

Zn

7.14

Ga

5.91

Ge

5.36

As

5,7

Se

4.7

Br

3.119

Kr

2.6

4

4

Rb

1.53

Sr

2.6

Y

5.51

Zr

6.4

Nb

8.4

Mo

10.2

Tc

11.5

Ru

12.5

Rh

12.5

Pd

12.0

Ag

10.5

Cd

8.6

In

7.3

Sn

7.3

Sb

6.7

Te

6.1

I

4.93

Xe

3.06

5

5

Cs

1.90

Ba

3.5

La

6.7

Hf

13.1

Ta

16.6

W

19.3

Re

21.4

Os

22.48

Ir

22.4

Pt

21.45

Au

19.3

Hg

13.55

Tl

11.85

Pb

11.34

Bi

9.8

Po

9.4

At

---

Rn

4.4

6

6

8.0 – 11.9 g/cm3

12.0 – 17.9 g/cm3

> 18.0 g/cm3

Mg

1.74

Symbol

Density in g/cm3C, for gases, in g/L

W