The Liquid State...The Liquid State A liquid forms when a gas condenses. • This occurs at low T...

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© 2008 Brooks/Cole 1

Chapter 11: Liquids, Solids and Materials


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.


The Liquid State


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).

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


The Liquid State


The Liquid State


low T

velocity or energy

num

ber o

f mol

ecul

es

Threshold for

escape

high T

Volatility increases with increased T.

Increase T = increase Ek = more can escape.

Vapor Pressure


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

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Boiling Point Boiling points (bp) vary with pressure.

Nor

mal

bp

= 78

.5°C

ethanol Water boils at 100°C

at sea level

1 atm

Nor

mal

bp

-40 -20 0 20 40 60 80 100 120 Temperature, °C

Vapo

r pre

ssur

e, m

mH

g

1000

500

0

H 2O

diethyl ether

bp =

34.

6°C

Salt 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 ΔHvap

R 1

T2 1

T1

Ethanol

Plot ln P versus 1/T

slope = – ΔHvap R


Clausius-Clapeyron Equation The normal bp of water = 100°C and ΔHvap = 40.71 kJ/mol. At what T will water boil if P = 500 Torr?

ln = – P2 P1

1 T2

1 T1

-ΔHvap R

ln = – 500 760

1 373

1 T2

- 40710 J g-1 8.315 J K-1mol-1

- 0.4187 = - 4896.3 - 2.681 x 10-3 1 T2

T2 = 361 K = 88 °C © 2008 Brooks/Cole 16

Phase Changes: Solids, Liquids & Gases

low T

velocity or energy

num

ber o

f mol

ecul

es

Threshold for

escape

high T


Phase Changes

For H2O(l) → H2O(g) ΔH° = ΔH°vap = +40.7 kJ/mol

For H2O(g) → H2O(l) ΔH° = ΔH°cond = -40.7 kJ/mol


Melting and Freezing

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Phase Changes

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

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

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


Phase Changes


Phase Changes

endothermic process

(heat added)

exothermic process

(heat released)


Phase Changes

Substance mp(°C) ΔHfus(kJ/mol) bp(°C) ΔHvap(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) ΔHfus(kJ/mol)* bp(°C) ΔHvap(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) ΔHfus(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. ΔH = mcΔT = (100g)(2.06 Jg-1°C-1)(0-[-20]°C) = 4.1 kJ

Convert the ice to water ΔH = nΔHfus = (100g/18.02g mol-1)(6.020 kJ/mol) = 33.4 kJ

Heat the water from 0°C to 100°C. ΔH = mcΔT = (100g)(4.184 Jg-1°C-1)(100 - 0°C) = 41.8 kJ

Convert water to steam (at its normal bp). ΔH = nΔHvap = (100g/18.02 g mol-1)(40.7 kJ/mol) = 225.9 kJ

Heat steam from 100 to 120°C. ΔH = mcΔT = (100g)(1.84 Jg-1°C-1)(120 - 100°C) = 3.7 kJ

Heating Curve


Pre

ssur

e (a

tm)

Temperature (°C)

Phase Diagrams

Solid Liquid

Gas

critical point

triple point

Melting point curve

Vapor-pressure curve

Supercritical fluid


Phase Diagrams

Water T = 0.01°C, P = 4.58 mmHg CO2 T = -57°C, P = 5.2 atm


Supercritical CO2 is an important solvent. It is used:

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

ssur

e (m

m H

g)

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

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

These facts allow aquatic life to survive at low T.





Types of Solids


Types of Solids Solids can be divided into:


Crystalline Solids

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Crystalline Solids

•  The unit cell is a square. •  Each corner contributes ¼ 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

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

1

3

2 4 5 6 at rear

Each atom has 6 equivalent neighbors:


Cubic Unit Cells

68 % of the space is occupied.


Cubic Unit Cells

74% of the space is filled by atoms


Closest Packing of Spheres Metal crystals have equal-sized atoms (spheres).

Single layer

A large percentage of the space is occupied.

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Closest Packing of Spheres


Closest Packing of Spheres Start with “ab” (“b” layer in “a” holes)

Add another “a” layer (Green directly over green).


Closest Packing of Spheres


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...)


Ionic Crystal Structures Many ionic compounds have: •  sc or fcc negative-ion lattices. •  positive ions occupy “holes”.

CsCl unit cell

Cs+

Cl-

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Ionic Crystal Structures 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 ¼ Na+/edge = 3 Na+ 1 center x 1 Na+/center = 1 Na+

The NaCl unit cell contains 4 Cl- ions:

Ionic Crystal Structures

8 corners x 1/8 Cl- / corner = 1 Cl- 6 faces x ½ Cl- / face = 3 Cl-


Network Solids Some non-metals link together in extended networks.

Diamond Graphite

335 pm

141 pm


Materials Science Study of the relationships between the structure and the chemical and physical properties of materials.

The Liquid State...The Liquid State A liquid forms when a gas condenses. • This occurs at low T...

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