Zumdahl Ch 10 Liquids, Solids powerpoint

(1)

CHEMISTRY

(2)

Chapter

10

(3)

10.1-2 Intermolecular forces

Objective: Students will be able to

describe the various types of

(4)

Why are some substances solids, while

others are liquids or gases at room

(5)

(6)

Copyright© by Houghton Mifflin Company. All rights reserved. 6



solid



liquid



gas



Similar density

~ 1000

for most; solids

times

generally more less

(7)

Review of intermolecular forces

1)

ion-dipole

2)

dipole – dipole (incl H bonding)

3)

dipole – induced dipole

4)

instantaneous dipole – induced

dipole

Ex:

Na

+

& H

2

O

mixture

H

2

O

H

2

O & Cl

2

mixture

I

2

Note: All of these are ~ 10 – 15 % of the strength of

bond energies.

“

London dispersion forces

”

(8)

(9)



Dipole-dipole interactions result in

higher mp and bp for similar sized

(10)

Hydrogen bonding

(11)

Hydrogen bonding

An attraction between a

hydrogen atom on one

(12)

… or sometimes on the same molecule,

as in DNA.

(13)

The boiling points of covalent hydrides

High bp due

to hydrogen

bonding

(14)

(15)

Copyright© by Houghton Mifflin Company. All rights reserved. 15



Think of e

-movement as a

swarm of bees. The

more e

-

s, the more

movement, and the

greater the

dispersion forces.



Can you explain why

F

2

and Cl

2

are gases

but Br

2

is a liquid

(16)

Figure 14.7: The heating/cooling curve for

water heated or cooled at a constant rate.

(17)

Figure 14.8: Both

liquid water and

gaseous water

contain H

2

O

molecules.

(18)

Figure 14.9: Microscopic view of a liquid

near its surface.

(19)

Figure 14.10: Behavior of a liquid in a closed

container.

(20)

Figure 14.11: (a) Measuring vapor of a liquid by using a

simple barometer.

(b) The water vapor pushed the mercury level down.

(c) Diethyl ether shows a higher vapor pressure than

water.

(21)

Figure 14.12: Water rapidly boiling on a stove.

(22)

Figure 14.13: Bubble expands as H

2

O molecules

enter.

(23)

Figure 14.14: The formation of the bubble is

opposed by atmospheric pressure.

(24)

End of lesson

(25)

10.3 – 4 Types of solids

Objective

: Students will be able to

recognize various types of solid crystal

structures, and carry out relevant

calculations.

Note

: This topic is omitted on the current

AP exam.

(26)

Crystalline solids - solids with a regular particle

arrangement.

Amorphous solids - solids with disordered structure.

Lattice - a 3-D system of points designating the positions

of the particles in a crystal.

Unit cell - the smallest repeating unit of a crystal lattice.

See fig. 10.9, p. 432

X-ray diffraction - common way of determining

crystalline structures

.

(27)

a. metallic - delocalized electrons – the e- sea

model.

b. network - atoms bond together with strong

covalent bonds, forming giant molecules.

c. group 8A - noble gas atoms attracted to each

other

(28)

Section 10.4 - Structure and bonding in

metals

Metallic properties such as high thermal and

electrical conductivity, malleability and

ductility, are explained by non-directional

covalent bonding.



Closest packing - a model which pictures a

metallic crystal as containing spherical atoms

packed together and bonded in all directions,

to most efficiently use the available space. See

p. 432.

(29)



A. Simple cubic



Like oranges stacked

up on top of each

other… lined up

perfectly. (Not a

closest-packing

arrangement)

(30)



B. Cubic closest packed (ccp/fcc) structure -

an abc arrangement (the atoms in the first

layer do not lie directly above the atoms in

the third layer) with a face-centered cubic

unit cell. See p. 437.

(31)



C. Hexagonal closest packed (hcp/bcc)

structure - an aba arrangement (the atoms

in the first layer lie directly above the atoms

in the third layer) has a hexagonal unit cell;

aka body-centered cubic. See p. 437.



3a

2

= (4r)

2

(32)

Summary

# of atoms/unit

cell



A) simple cubic

1



B) fcc/ccp/abca

4

(33)

Sample math problem



Sodium’s density is 0.971 g/cm

3

and it

crystallizes with a body-centered cubic

unit cell.



(a) What is the radius of a sodium atom?



(b) What is the edge length of the cell?

(34)

Solution:

1) Determine mass of two atoms in a bcc cell:

2 atoms 22.99 g mol = 7.635 x 10

-23

g

mol 6.022 x 10

23

atoms

2) Determine the volume of the unit cell using V = m/D



7.635 x 10

-23

g cm

3

. = 7.863 x 10

-23

cm

3



0.971 g

3) Determine the edge length -- the answer to (b):

(35)

4) Use the Pythagorean

theorem to find the

radius.



x

2

+ x

2

+ x

2

= (4r)

2



3x

2

= 16r

2



r

2

= 3(4.284 x 10

-8

)

2

16



r = 1.855 x 10

-8

cm



The radius of the Na

atom is 185.5 pm. The

edge length is 428.4 pm.

(36)

Sample problem 2



Gold is a face centered cubic unit cell. D =

19.32 g/cm

3

. Find the radius of the gold atom

(37)



Solution:

1) Determine mass of four atoms in an fcc cell:

4 atoms 196.97 g mol = 1.308 x 10

-21

g

mol 6.022 x 10

23

atoms

2) Find the volume of the unit cell using V = m/D

1.308 x 10

-21

g cm

3

= 6.772 x 10

-23

cm

3

19.32 g

3) Find the edge length

(38)



Bonding in most metals is both strong

and non directional. That is, although

it is difficult to separate metal atoms,

it is relatively easy to move them,

provided the atoms stay in contact

with each other.

(39)



Molecular orbital (MO) (or band model)

- in this model, the e-s travel around

the metal crystal in MOs formed from

the valence atomic orbitals of the metal

atoms. When many metal atoms

interact, the large number of resulting

molecular orbitals become more closely

spaced and finally form a virtual

continuum of levels, called bands.

See fig 10.19 and 10.20

(40)

End of lesson

(41)

10.5-6 Network solids;

molecular solids



(42)

Ex.

mp



H

fusion

Held

together by

metals

Al

varies

high

cations & e

-

s

nonpolar

molecular

solids

I

2

, CO

2

low

London

dispersion

forces

polar

molecular

solids

H

2

O

medium

medium dipole

interactions/

H bonding

network

solids

diamond,

graphite,

SiO

2

v. high

covalent

bonds

ionic solids NaCl

v. high

ionic bonds

(obviously)

amorphous

(43)

Network solids and amorphous solids.

Can you match them?

a.

silica (SiO

2

)

b.

diamond

c.

glass

d.

graphite

1.

2.

3.

(44)

Silica (SiO

2

)

Each Si is covalently bonded

to four O atoms. Net

formula = SiO

2

.

No



bonds; very different

from CO

2

(45)

silica

silicates

glass

similarities: all contain Si, O

network

solid

ionic

compounds

amorphous

SiO

2

O/Si ratio >

2:1

formed by

heating silica

Si

single-bonded to 4

O (shared)

(46)

Figure 14.15: Sodium and chloride ions.

(47)

Figure 14.18: A molecular solid.

(48)

Figure 14.19: The packing of Cl¯ and Na

+

ions in

solid sodium chloride.

(49)

Solid P

(50)



Alloy

- a mixture of

elements with metallic

properties.



1. Substitutional alloy -

some of the host metal

atoms are replaced by

other metal atoms of similar

size.



2. Interstitial alloy - formed

when some of the

interstices (holes) in the

closest packed metal

structure are occupied by

small atoms.

(51)

Brass… which type of alloy is this?

(52)

Steel… Which type of alloy is this?

(53)

As intermolecular forces

increase, what

happens to each of the following? Why?



boiling point



viscosity



surface tension



enthalpy of fusion



freezing point



vapor pressure



heat of vaporization

(54)

End of section

(55)

10.8 Vapor pressure and

changes of state



describe vapor pressure and carry

(56)

Vaporization



Vaporization – the process of a liquid

changing to a gas





H

vap

– heat of vaporization - the heat

required to vaporize 1 mole of a liquid at

1 atm pressure



Water has a very high



H

vap

(40.7 kJ/mol)

due to strong hydrogen bonding.

(57)

Vapor pressure

(58)

Behavior of a liquid in a closed container

(59)

Different liquids have

different vapor pressure.



Weak intermolecular attractions



higher v.p.



more volatile (evaporates

readily)

(60)

Vapor pressure

(61)

(62)



Notice as T



v.p.



.



The mathematical relationship is given by

the Clausius-Clapeyron equation:



l

n(P

vap

) = -



H

vap

1 + C

R T

(63)



l

n(P

vap

) = -



H

vap

1 + C

R T

This means a plot of

l

n(P

vap

) versus 1/T

(reciprocal of Kelvin temp) would give a

straight line with slope –



H

vap

/R.

(64)

8

7

6

5

4

3

2

1

0

l

n

P

(65)

slope =



y = 0 - 8.00

x (4.03 x 10

-3

K

-1

) - (2.75 x 10

-3

K

-1

)



-8.00 = -6.25 x 10

3

K

1.28 x 10

-3

K

-1



slope = -



H

R



H = -slope · R = (+6.25 x 10

3

K) · (8.3145

(66)

Clausius – Clapeyron Equation

P

vap

= vapor pressure

Δ

H

vap

= enthalpy of vaporization

R

= 8.3145 J/K·mol

(67)

End of lesson



See homework on board.

(68)

10.9 Phase diagrams



identify key points on a phase

diagram and use a phase diagram to

describe important characteristics of a

substance.



Note: This topic is omitted on the

(69)

Heating Curve for Water

(70)



phase diagram: a convenient way of representing

the phases of a substance as a function of

temperature and pressure. It shows:



triple point (s, l, g all exist at equilibrium)



critical point (T and P above which gas and liquid

phases are indistinguishable)

(71)

Phase

diagram

L

G

S

A typical phase

diagram. The

solid green line

applies to most

substances; the

dotted green line

gives the

anomalous

(72)

Phase

diagram for

water

(73)

Phase diagram for water

Normal

melting

point

Normal

boiling

point

Solid-liquid eq’m line

Liquid-gas eq’m line

(74)

Phase diagram for water

Notice:



S-L eq’m line has a negative

slope.



This means ice is less dense

than water.



When pressure is put on ice,

it melts.

AP Chem annual skating

(75)

(76)

Phase

diagram for

carbon

dioxide

(77)

Phase

diagram for

carbon

dioxide

Notice:



S-L eq’m line has a positive slope.



This means dry ice is more dense than liquid

CO

2

.

(78)

End of lesson



See homework on board.

(79)



Extra stuff from the publishers

…

(80)

Chapter 10

(81)

Hydrogen Bonding in Water



Blue dotted lines are the

(82)

Concept Check

Which are stronger, intramolecular bonds or

intermolecular forces?

(83)

Phase Changes



When a substance changes from solid to liquid to

gas, the molecules remain intact.



The changes in state are due to changes in the forces

(84)

Schematic Representations of the Three States of Matter

(85)

Phase Changes



Solid to Liquid



As energy is added, the motions of the molecules

increase, and they eventually achieve the greater

movement and disorder characteristic of a liquid.



Liquid to Gas



As more energy is added, the gaseous state is

(86)

Densities of the Three States of Water

(87)

Dipole-Dipole Forces



Dipole moment – molecules with polar bonds often

behave in an electric field as if they had a center of

positive charge and a center of negative charge.



Molecules with dipole moments can attract each other

electrostatically. They line up so that the positive and

negative ends are close to each other.

(88)

Dipole-Dipole Forces

(89)

Hydrogen Bonding

(90)

Hydrogen Bonding



Strong dipole-dipole forces.



Hydrogen is bound to a highly electronegative atom –

(91)

London Dispersion Forces

(92)

London Dispersion Forces



Instantaneous dipole that occurs accidentally in a

given atom induces a similar dipole in a neighboring

atom.



Significant in large atoms/molecules.

(93)

Melting and Boiling Points



In general, the stronger the intermolecular forces, the

(94)

The Boiling Points of the Covalent Hydrides of the Elements in Groups

4A, 5A, 6A, and 7A

(95)

Concept Check

Which molecule is capable of forming stronger

intermolecular forces?

N

2

H

2

O

(96)

Concept Check

Draw two Lewis structures for the formula C

2

H

6

O

and

compare

the boiling points of the two molecules.

C

H

C

H

O H

C

(97)

Concept Check

Which gas would behave

more ideally

at the same

conditions of P and T?

CO or

N

2

(98)

Liquids



Low compressibility, lack of rigidity, and high

density compared with gases.



Surface tension – resistance of a liquid to an increase

in its surface area:



Liquids with large intermolecular forces tend to

(99)

Liquids



Capillary action – spontaneous rising of a liquid in a

narrow tube:



Cohesive forces – intermolecular forces among

the molecules of the liquid.



Adhesive forces – forces between the liquid

(100)

Convex Meniscus Formed by Nonpolar Liquid Mercury



Which force dominates alongside the glass tube –

cohesive or adhesive forces?

(101)

Concave Meniscus Formed by Polar Water



Which force dominates alongside the glass tube –

cohesive or adhesive forces?

(102)

Liquids



Viscosity – measure of a liquid

’

s resistance to flow:



Liquids with large intermolecular forces or

(103)

Solids



Amorphous Solids:



Disorder in the structures



Glass



Crystalline Solids:

(104)

Three Cubic Unit Cells

and the Corresponding Lattices

(105)

Bragg Equation



Used to determine the interatomic spacings.

n = integer

= wavelength of the X rays

d = distance between the atoms

= angle of incidence and reflection

= 2 sin





n

d



(106)

Bragg Equation

= 2 sin





(107)

Types of Crystalline Solids



Ionic Solids – ions at the points of the lattice that

describes the structure of the solid.



Molecular Solids – discrete covalently bonded

molecules at each of its lattice points.



Atomic Solids – atoms at the lattice points that

(108)

Examples of Three Types of Crystalline Solids

(109)

Classification of Solids

(110)

Closest Packing Model



Closest Packing:



Assumes that metal atoms are uniform, hard

spheres.

(111)

The Closest Packing Arrangement of Uniform Spheres



aba

packing – the 2

nd

layer is like the 1

st

but it is displaced so

that each sphere in the 2

nd

layer occupies a dimple in the 1

st

layer.

(112)

The Closest Packing Arrangement of Uniform Spheres



abc

packing – the spheres in the 3

rd

layer occupy dimples in the

2

nd

layer so that no spheres in the 3

rd

layer lie above any in the

1

st

layer.

(113)

Hexagonal Closest Packing

(114)

Cubic Closest Packing

(115)

The Indicated Sphere Has 12 Nearest Neighbors



Each sphere in both

(116)

The Net Number of Spheres in a Face-Centered Cubic Unit Cell

(117)

Concept Check

Determine the number of metal atoms in a unit cell if the

packing is:

a)

Simple cubic

b)

Cubic closest packing

(118)

Concept Check

A metal crystallizes in a face-centered cubic

structure.

Determine the relationship between the radius of

the metal atom and the length of an edge of the

unit cell.

(119)

Concept Check

Silver metal crystallizes in a

cubic closest packed

structure

. The face centered cubic unit cell edge is

409 pm.

Calculate the

density

of the silver metal.

(120)

Bonding Models for Metals



Electron Sea Model

(121)

The Electron Sea Model



A regular array of cations in a

“

sea

”

of mobile

(122)

Molecular Orbital Energy Levels Produced When Various Numbers of

Atomic Orbitals Interact

(123)

The Band Model for Magnesium

(124)

Metal Alloys



Substitutional Alloy – some of the host metal atoms

are replaced by other metal atoms of similar size.



Interstitial Alloy – some of the holes in the closest

(125)

Two Types of Alloys



Brass is a

substitutional

alloy.



Steel is an

(126)

The Structures of Diamond and Graphite

(127)

Partial Representation of the Molecular Orbital Energies in

a) Diamond b) a Typical Metal

(128)

The

(129)

Ceramics



Typically made from clays (which contain silicates)

and hardened by firing at high temperatures.



Nonmetallic materials that are strong, brittle, and

(130)

Semiconductors



n-type semiconductor – substance whose

conductivity is increased by doping it with atoms

having more valence electrons than the atoms in the

host crystal.



p-type semiconductor – substance whose

conductivity is increased by doping it with atoms

(131)

Energy Level Diagrams for

(132)

Silicon Crystal Doped with

(a) Arsenic and

(b) Boron

(133)

Three Types of Holes in Closest Packed Structures

1)

Trigonal holes are formed by three spheres in the

(134)

2)

Tetrahedral holes are formed when a sphere sits in

(135)

Three Types of Holes in Closest Packed Structures

3)

Octahedral holes are formed between two sets of

(136)



For spheres of a given diameter, the holes increase in

size in the order:

(137)

Types and Properties of Solids

(138)

The Rates of Condensation and Evaporation

(139)

Vapor Pressure



Pressure of the vapor present at equilibrium.



The system is at equilibrium when no net change

(140)

Concept Check

What is the vapor pressure of water at 100°C? How

do you know?

(141)

Vapor Pressure

(142)

Vapor Pressure



Liquids in which the intermolecular forces are strong

have relatively low vapor pressures.



Vapor pressure increases significantly with

(143)

Vapor Pressure

(144)

Concept Check

The vapor pressure of water at 25°C is 23.8 torr, and

the heat of vaporization of water at 25°C is 43.9 kJ/

mol. Calculate the vapor pressure of water at 65°C.

(145)

Concept Check

Which would you predict should be

larger

for a

given substance:



H

vap

or



H

fus

?

(146)

Concept Check

As intermolecular forces

increase

, what happens to each

of the following? Why?



Boiling point



Viscosity



Surface tension



Enthalpy of fusion



Freezing point



Vapor pressure