The Liquid State © 2008 Brooks/Cole 1 Chapter 11: Liquids, Solids and Materials Chemistry: The Molecular Science Moore, Stanitski and Jurs © 2008 Brooks/Cole 2Published in: American

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© 2008 Brooks/Cole 1

Chapter 11: Liquids, Solids and Materials

Chemistry: The Molecular ScienceMoore, Stanitski and Jurs


The Liquid State

A liquid forms when a gas condenses.

• This occurs at low T (and/or high P).

• Average intermolecular attraction > average Ek.

Molecules are close together and in constant motion

They have a distribution of speeds (and Ek)

A few have enough Ek to slide past each other.


Liquids are very hard to compress

• Molecules are close together and resist movingcloser.

• Used as hydraulic fluids to transmit pressure.

The Liquid State

Viscosity = Resistance to flow.

• High viscosity = slow flow.

low viscosity: water, ethanol, gasoline…

high viscosity: honey, corn syrup …


High viscosity is caused by:

• Large intermolecular attractions.

• Large entanglement of molecules.

Higher T = lower viscosity

• Ek > Ebarrier required to move past another molecule.

• Increase T = increase average Ek.

More molecules have Ek > Ebarrier

The Liquid State


Liquid molecules attract each other.

Surface molecules experience unbalanced forces.

They are not as well stabilized.

Smaller surface area higher stability.

The Liquid State


Raindrops are spherical (surface/V ratio minimized).

Surface tension = E required to form a flat surface.

Substance Formula Surface Tension

J/m2 at 20°C

Octane C8H18 2.16 x 10-2

Ethanol C2H5OH 2.23 x 10-2

Chloroform CHCl3 2.68 x 10-2

Benzene C6H6 2.85 x 10-2

Water H2O 7.29 x 10-2

Mercury Hg 46. x 10-2

The Liquid State

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Why don’t all liquids form spheres in all cases?

It’s a balancing act …

Gravity pulls down (giving a puddle).

Surface tension contracts (into a sphere).

The “winner” depends on the amount of liquid and

any surface/liquid attractions. A drop of water:

puddles on a dirty surface

“beads” on a waxy surface

The Liquid State


Capillary action

A competition between:

• Adhesion: forces between a surface and a liquid.

• Cohesion: forces between liquid molecules.

The shape of the meniscus illustrates the relative

strength of the adhesive and cohesive forces

The Liquid State


The Liquid State


Volatility = tendency of a liquid to vaporize.

low T

velocity or energy

num

ber

of m

ole

cule

s

Threshold

for

escape

high T

Volatility increases with increased T.

Increase T = increase Ek = more can escape.

Vapor Pressure


If a liquid container is:

Open – the liquid can evaporate completely.

Closed – the liquid evaporates but cannot disperse.

Vaporized molecules can condense.

• more vapor = faster condensation

At some time:

evaporation rate = condensation rate.

P increases but reaches a maximum.

Vapor Pressure


Vapor Pressure

Equilibrium vapor pressure is independent of the

volume or surface area of the liquid

A liquid boils when the vapor P = atmospheric P

• Vaporization occurs throughout the liquid.

• Bubbles of pure vapor form and rise to the surface.

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The normal bp occurs at 1 atm = 760 mmHg.

Boiling Point

Boiling points (bp) vary with pressure.

Norm

al bp =

78.5

°C

ethanol Water boils

at 100°C

at sea level

1 atm

Norm

al bp

-40 -20 0 20 40 60 80 100 120

Temperature, °C

Va

po

r p

ressu

re,

mm

Hg

1000

500

0

H 2O

diethyl ether

bp =

34.6

°CSalt Lake City (4400 ft) P=650 mmHg

Water boils at 95°C

in Salt Lake City


Clausius-Clapeyron Equation

The relationship between vapor pressure and T:

ln = – –P2

P1

Hvap

R

1

T2

1

T1

Ethanol

NOTE: Absolute T must be used.

Plot ln P versus 1/T

slope = – Hvap

R


Clausius-Clapeyron Equation

The normal bp of water = 100°C and Hvap = 40.71

kJ/mol. At what T will water boil if P = 500 Torr?

ln = – P2

P1

1

T2

1

T1

- Hvap

R

ln = – 500

760

1

3731

T2

- 40710 J g-1

8.315 J K-1mol-1

- 0.4187 = - 4896.3 - 2.681 x 10-31

T2

T2 = 361 K = 88 °C


Phase Changes: Solids, Liquids & Gases

Vaporizationlow T

velocity or energy

num

ber

of m

ole

cule

s

Threshold

for

escape

high T

• High Ek molecules leave.

• The average Ek of the liquid drops.

• The liquid cools.

Vaporization is endothermic

Condensation

The reverse process is exothermic


Phase Changes

H°vap = – H°cond

Molar heat of vaporization, H°vap

Heat required to vaporize 1 mol (P = 1 bar).

Molar heat of condensation, H°cond

Heat released when 1 mol condenses (P = 1 bar).

For H2O(l) H2O(g) H° = H°vap = +40.7

kJ/molFor H2O(g) H2O(l) H° = H°cond = -40.7 kJ/mol


Melting and FreezingThe molecules in a solid are closely packed and in

constant motion. They

vibrate about fixed positions

do not have enough E to slide past each other.

A solid melts when

• Ek > forces holding the particles in their fixed

positions.

Melting is endothermic.

Freezing (solidification, crystallization) is exothermic.

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Phase Changes

H°fus = – H°cryst

Molar heat of crystallization, H°cryst

Heat released when 1 mol crystallizes (P = 1 bar).

For H2O(s) H2O(l) H° = H°fus = +6.02 kJ/mol

For H2O(l) H2O(s) H° = H°cryst = -6.02 kJ/mol

Normal melting or freezing occurs at P = 1 atm

Molar heat of fusion, H°fus

Heat required to melt 1 mol (P = 1 bar).


Phase Changes

Sublimation – the direct conversion of a solid to a gas

Examples

Solid CO2 (dry ice) sublimes.

Ice (v.p. = 4.6 mmHg @ 0°C) in a frost-free ‘fridge is removed

by blowing dry air over it…

Deposition is the reverse process: (g) (s)

Example: plating Al onto the surface of a CD

H°sub = – H°dep


Phase Changes

endothermic

process

(heat added)

exothermic

process

(heat released)


Phase Changes

Substance mp(°C) Hfus(kJ/mol) bp(°C) Hvap(kJ/mol)

O2 (16 e-) –248 0.445 –183.0 6.8

F2 (18 e-) –220 1.020 –188.1 6.54

Cl2 (34 e-) –103 6.406 –34.6 20.39

Br2 (70 e-) –7.2 10.794 59.6 29.54

Nonpolar molecules

London forces increase as the number of e- increase.

The data support this trend…


Phase Changes

Substance mp(°C) Hfus(kJ/mol)* bp(°C) Hvap(kJ/mol)*

SO2 (32 e-) –76 7.4 –10.0 24.9

HCl (18 e-) –115 2.0 –85.1 16.2

HBr (36 e-) –87 2.4 –66.8 17.6

H2O (10 e-) 0 6.0 +100.0 40.7

HF (10 e-) –83 4.6 +19.5 7.5

NH3 (10 e-) –78 5.6 –33.5 23.4

Polar molecules

More difficult to explore trends.

Values depend on London, dipole forces and H-bonding.

*At the normal phase-transition temperature (Air Liquide Gas Encyclopedia)


Phase Changes

Substance mp(°C) Hfus(kJ/mol)

NaCl 800 30.21

NaBr 747 25.69

NaI 662 21.95

Ionic solids

All Na+ ions / halide ion (-1 charge) compounds.

They only differ in halide size: I- > Br- > Cl-

Larger ion further apart weaker attraction

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Heating Curve

Convert 100 g of ice at -20°C into vapor at 120°C.


TOTAL heat required = 308.9 kJ

Heat the ice to 0°C.

H = mc T = (100g)(2.06 Jg-1°C-1)(0-[-20]°C) = 4.1 kJ

Convert the ice to water

H = n Hfus = (100g/18.02g mol-1)(6.020 kJ/mol) = 33.4 kJ

Heat the water from 0°C to 100°C.

H = mc T = (100g)(4.184 Jg-1°C-1)(100 - 0°C) = 41.8 kJ

Convert water to steam (at its normal bp).

H = n Hvap = (100g/18.02 g mol-1)(40.7 kJ/mol) = 225.9 kJ

Heat steam from 100 to 120°C.

H = mc T = (100g)(1.84 Jg-1°C-1)(120 - 100°C) = 3.7 kJ

Heating Curve


Pre

ssu

re (

atm

)

Temperature (°C)

Phase Diagrams

Solid Liquid

Gas

critical point

triple pointsublimation

deposition

freezing

melting

condensation

vaporization

Melting point

curve

Vapor-pressure

curve

Supercritical

fluid


Phase Diagrams

Triple point

Three phases in equilibrium. T and P are fixed.

Water T = 0.01°C, P = 4.58 mmHg

CO2 T = -57°C, P = 5.2 atm

Critical point

The end of the liquid/gas equilibrium line. It occurs at

the critical temperature, Tc and critical pressure, Pc

Water Tc = 374°C, Pc = 218 atm

CO2 Tc = 31°C, Pc = 73 atm

Above Tc a gas cannot be liquefied (at any P)


Supercritical CO2 is an important solvent. It is used:

Above Tc and Pc the substance is neither a liquid nor

a gas. It is a supercritical fluid.

It has:

• to extract caffeine from coffee beans

• as a dry-cleaning fluid

• a density characteristic of a liquid.

• flow properties of a gas.

Critical Temperature and Pressure


Water Phase Diagram

For most materials, the solid/liquid line has a positive slope.

Water is unusual. Ice can be melted by increased P !

4.58

Solid Liquid Gas

Pre

ssu

re (

mm

Hg

)

Temperature (°C)0 0.01 100

760

triple point

normal fp normal bp

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Physical Comparison with Importance in

Property Other Substances and Biological Environment

Specific heat capacity Highest of all liquids and Moderates T in the environment and in

(4.18 J g-1 °C-1) solids except NH3 organisms; climate affected by

movement of water (e.g., Gulf Stream)

Heat of fusion Highest of all molecular Freezing water releases large quantity

(333 J/g) solids except NH3 of thermal E; used to save crops from

freezing by spraying them with liq. water

Heat of vaporization Highest of all molecular Condensation of water vapor in clouds

(2250 J/g) substances releases large quantities of thermal E

fueling storms

Surface tension Highest of all molecular Contributes to capillary action in plants

(7.3 x 10-2 J/m2) liquids causes formation of spherical droplets;

supports insects on water surfaces

Thermal conductivity Highest of all molecular Aids heat transfer in organisms;

(0.6 J s-1 m-1 °C-1) liquids rapidly cools organisms immersed in

cold water, causing hypothermia

Table 11.4 Unusual Properties of Water

Water: A Liquid with Unusual Properties


Most liquids: Lower T = higher density.

Most materials: dsolid > dliquid

Not water! In winter:

• Surface water cools & sinks until T=4°C.

• Water (T < 4°C) floats on the 4°C layer.

• The top layer cools to 0°C.

• Water freezes “top-down”.

• Ice floats on water, insulating the water below it.

These facts allow aquatic life to survive at low T.





Types of Solids

Ionic (ionic bonding)

• Hard, brittle, high mp. Some are water soluble.

Metallic (metallic bonding)

• Malleable, ductile, good heat/electrical conduction, wide

range of hardness and mp (Pt, Fe, Pb, Hg …)

Molecular (London forces, dipole/dipole forces, and H-bonds)

• Low to moderate mp/bp, soft, poor conductors.

Network (covalent bonds; atoms in infinite arrays)

• Wide range of properties (diamond, graphite, mica …)


Types of Solids

Solids can be divided into:

Amorphous solids

• no regular repeating units.

• melt over a range of T.

Examples: glass, plastics, ceramics…

Crystalline solids

• long-range order.

• sharp melting points.

Examples: ice, table salt, gems…


Crystalline Solids

Crystals have planar faces and sharp angles.

The angles are characteristic of the substance.

– they don’t depend on crystal size

– they reflect the shape of the crystal lattice

The smallest part of a lattice that can be used

to construct the full lattice is the unit cell.

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Crystalline Solids

2d unit cell example:

• The unit cell is a square.

• Each corner contributes of a

circle to the unit cell.

• Net of 1 circle/unit cell.

• The entire lattice can be formed

by adding unit cells to each

face


Cubic Unit Cells

Crystals have 3d unit cells. There are 7 types.

The simplest are the cubic cells with 3 subtypes:

– simple (or primitive) cubic (sc).

– body-centered cubic (bcc).

– face-centered cubic (fcc).

Faces meet at 90° with equal-length sides.


Cubic Unit Cells

Primitive (simple) cubic

• Atoms sit at the corners of squares.

• Layers stack atom-on-atom.

• One atom/cell (8 x 1/8).

• Atoms occupy 52% of the total V.

1

3

24 56

at rear

Each atom has6 equivalent

neighbors:


Cubic Unit Cells

Body-centered cubic

• An extra (unshared) atom sits inthe center

• (8 x 1/8 ) + 1 = 2 atoms/cell

68 % of the space is occupied.


Cubic Unit Cells

Face-centered cubic

1 atom at the center of each face.

Face atoms are shared by 2 cells.

(6 x + 8 x 1/8 ) = 4 atoms/cell.

74% of the space is filled by atoms


Closest Packing of Spheres

Metal crystals have equal-sized atoms (spheres).

Single

layer

In closest packing, each atom has 6 neighbors/layer

A large percentage of the space is occupied.

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Layers stack with the atoms in one

layer resting in “holes” of another.

hexagonal close packing

An abab structure



Start with “ab”

(“b” layer in “a” holes)

Add another “a” layer

(Green directly over green).


Cubic Close Packing

has an abcabc structure

(the unit cell is face centered cubic)



Cubic Close Packed

Start with “ab”.

(“b” layer in “a” holes)

Add “c” (gold) above holes in

the original green “a” layer.


Ionic Crystal Structures

Ionic crystal structures are more complex.

The ions making up the crystal:

• are not identical to each other.

• may be of very different sizes

• may not have the same charge magnitude.

• may not be “spheres” (polyatomic ions...)



Many ionic compounds have:

• sc or fcc negative-ion lattices.

• positive ions occupy “holes”.

CsCl unit cell

Cs+

Cl-

CsCl

Simple cubic Cl- (Cs+ in the central hole).

• This is not bcc

• Only use “bcc” if all atoms are identical

• In a unit cell: 8 x 1/8 = 1 Cl- and 1 Cs+

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NaCl has an fcc Cl- lattice; Na+ in octahedral holes.

• Each Na+ is surrounded by 6 Cl-.

• Each Cl- is surrounded by 6 Na+.


It also contains 4 Na+:

12 edges x Na+/edge = 3 Na+

1 center x 1 Na+/center = 1 Na+

The NaCl unit cell contains 4 Cl- ions:


8 corners x 1/8 Cl- / corner = 1 Cl-

6 faces x Cl- / face = 3 Cl-


Network Solids

Some non-metals link together in extended networks.

DiamondGraphite

335 pm

141 pm


Materials Science

Study of the relationships between the structure and

the chemical and physical properties of materials.

Four major classes of materials:

Metals

Polymers

• Natural (cotton, rubber…) and synthetic (teflon,

polyethylene…). Thermal and electrical insulators.Nonmagnetic.


Materials Science

Ceramics

• Nonmetallic. Often clay based. Porcelain,

cement… Poor thermal conductors. Some are

magnetic. Usually crystalline. Include

semiconductors and glasses.

Composites

• Combinations of the 3 other materials. Properties

that exceed those of the individual components.


Metals

Opaque, shiny, usually crystalline. Form alloys.

• High electrical conductivity

• High thermal conductivity

• Ductile and malleable

• Mostly nonmagnetic (Fe, Ni, Co are exceptions)

• Have luster

Polished metals reflect light. Most metals reflect all

wavelengths equally (appear silvery-white).

• Insoluble in water and other common solvents

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Metals

Metal Hfus mp

(kJ/mol) (°C)

Hg 2.3 -38.8

Ga 7.5 29.8

Na 2.6 97.9

Li 3.0 180.5

Al 10.7 660.4

U 12.6 1132.1

Fe 13.8 1535.0

Ti 20.9 1660.1

Cr 16.9 1857.0

W 35.2 3410.1

Hfus (and mp) vary widely:

Hfus


Metallic bonding = Nondirectional attraction between

M+ and a sea of mobile valence e-.

Electrons in Metals, Semiconductors, and Insulators

Valence e- spread throughout the M+ lattice, holding

the M+ ions together.

An electric field can pull valence e- to a “+” electrode

(electrical conduction).

M+ can move without destroying M+ ~ e- attractions

(ductility and malleability).


Band Theory

Every atom has E-levels (s, p, d…).1

Number of

atoms

e- inatoms

e- in latticeenergy levels

With many atoms, the E-levels get

so close together they form a band

3

2

12

As multiple atoms come togetherorbitals interact.

many

filled

empty

Electrons in Metals, Semiconductors, and Insulators


The valence band contains the valence e-.

Higher-E orbitals form the conduction band

In a metal, the valence and conduction band overlap.

e- can move freely between bands.

Metals, Semiconductors & Insulators


band gap

insulator semi-conductorconductor

MetalBands overlap.

Lots of open levels.

e- move freely.

InsulatorLarge band gap

e- cannot jump the gap.

No conduction.

SemiconductorNarrow band gap. Some e- can jump the gap…

Higher T = more energetic e- = better conductor.

Metals, Semiconductors & Insulators


Metals are better conductors at low T.

typical metal

Re

sis

tan

ce

Temperature (K)

superconductor

Tc

critical T

Superconductors have zero resistance (are perfectconductors) at some (low) T.

0

Superconductors

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Metal Tc (K)

Aluminum 1.15

Gallium 1.10

Tin 3.72

Mercury 4.15

Lanthanum 4.9

Lead 7.2

Several metals are superconductors at low T

Nb3Sn alloy 18.1

YBa2Cu3O7 90.

LaBa2Cu3Ox 35.

He(l) boils at 4.2 K

Many He(l)-cooled

magnets use Nb-alloys

1st ceramic

superconductor (1986).

N2(l) boils at 77 K

“Y123” found 4 months later.

Highest Tc so far…Hg0.8Tl0.2Ba2Ca2Cu3O8.23 138.

Some alloys are better…

Superconductors


Silicon and the Chip

Ultra-pure Si is prepared by zone refining. Impurities are more

soluble in Si(l) than in Si(s). A molten zone is moved through a

sample. Impurities move along in the liquefied portion.



Doping is the controlled addition of some other element.



A p-n junction. Charge carriers: holes (p-type) and e- (n-type)



A photovoltaic cell. Light drives e- around an external circuit


Cement, Ceramics and Glasses

Amorphous solids:

Concrete = a mix of cement, sand and gravel.

Cement

Tiny metal-oxide particles (CaO, SiO2, Al2O3 & SiO2).

They crystallize irregularly when hydrated.

Made by heating limestone, sand, clay and iron oxide.

High compressive, but low tensile strength.

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Ceramics – made from clay or other natural earths.

Permanently hardened by heat.

Silicate ceramics – from clays. China clay (kaolin) is

white and used for table china. Red clay contains Fe.

Oxide ceramics – from metal oxides at high P and T.

Used in high-T applications. Excellent insulators. Canhave good mechanical strength.

Nonoxide ceramics – include Si3N2, SiC, and BN.

High P, T preparation. Hard, strong, but brittle.



Glasses – optically transparent amorphous solids.

Simplest: vitreous silica = SiO2 in a 3d-lattice.

No long range order. (mp = 1800°C)

Soda lime glass, a SiO2, Na2O, CaO and MgO mix.

More inert and easier to make (mp ~800°C).

It can be colored (blue = add CoO, violet = MnO2…)

Documents

The Liquid State © 2008 Brooks/Cole 1 Chapter 11: Liquids, Solids and Materials Chemistry: The Molecular Science Moore, Stanitski and Jurs © 2008 Brooks/Cole 2Published in: American