Intermolecular Forces, Then we explained how atoms combine to form - - PDF document

SLIDE 1

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Intermolecular Forces, Liquids, and Solids

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Intermolecular forces are the piece we need to add to the puzzle to explain the world around us. We first explained atoms, and how to build up the periodic table from quantum numbers. Then we explained how atoms combine to form molecules: the most common way we find most atoms in nature. Now, we're going to use intermolecular forces to combine molecules to create the common states of matter. Without intermolecular forces, we wouldn't have tables, lakes, wall...or even our bodies. Intermolecular forces shape our world.

Intermolecular Forces Slide 3 / 92

While there are many states of matter, the three common states that dominate our world are gases, liquids and solids. These are the states of matter we'll be studying. We won't be discussing more exotic states such as plasma, nuclear matter, etc.

States of Matter Slide 4 / 92 States of Matter

The fundamental differences between states of matter is: · the distance between particles · the particles' freedom to move

cool or increase pressure heat or decrease pressure cool heat Particles are far apart, total freedom, much of empty space, total disorder disorder, freedom, free to move relative to each other, close together

• rderd arrangement,

particles are in fixed positions, close together Gas Liquid Crystalline solid

Slide 5 / 92 Characteristics of the States of Matter

Gas Assumes the shape of its container Expands to the volume of its container Is compressible Flows easily Diffusion within a gas is rapid Liquid Assumes the shape of the part of a container it occupies Does not expand to the volume of its container Is virtually incompressible Flows easily Diffusion within a liquid is slow Solid Retains its own shape, regardless of container Does not expand to the volume of its container Is virtually incompressible Does not flow Diffusion within a solid is very very slow

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In the solid and liquid states particles are closer together, we refer to them as condensed phases.

Condensed Phases

cool or increase pressure heat or decrease pressure cool heat Particles are far apart, total freedom, much of empty space, total disorder disorder, freedom, free to move relative to each other, close together

• rderd arrangement,

particles are in fixed positions, close together Gas Liquid Crystalline solid

SLIDE 2

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1

Which of the following is not a type of solid?

A

ionic

B

molecular

C

covalent-network

D

supercritical

E

metallic

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2 Which of the below is a characteristic of a gas?

A It fills only a portion of its container. B Its molecules are in relatively rigid positions. C It takes on the shape of its container. D It is not compressible. E

Diffusion is very slow within it.

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3 Which of the below is a characteristic of a liquid?

A It fills only a portion of its container. B Its molecules are in relatively rigid positions. C It takes on the shape of its container.

D It is compressible. E Diffusion is very rapid within it.

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4 Which of the below is a characteristic of a solid?

A It fills all of its container. B Its molecules are in relatively rigid positions.

C It takes on the shape of its container.

D It is compressible. E Diffusion is very rapid within it.

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5

Together, liquids and solids, constitute __________ phases of matter.

A

the compressible

B

the fluid

C

the condensed

D

all of the above

E

the disordered

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The state of a substance at a particular temperature and pressure depends on two opposing properties: · Intermolecular Forces: the strength of the attractions between the particles, which pulls them together · the kinetic energy of the particles, which pulls them apart Without intermolecular forces, all molecules would be ideal gases...there would be no liquids or solids.

States of Matter

SLIDE 3

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Boiling represents a transition from a liquid to a gas. To make that transition, molecules in the liquid must break free of the intermolecular forces that bind them. The kinetic energy of the molecules is proportional to the temperature: as temperature rises, so does kinetic energy. The temperature where the molecules' energy

• vercomes intermolecular forces is called the boiling

point. The boiling point is a measure of the strength of the intermolecular forces: the higher the boiling point - the stronger the intermolecular forces.

Intermolecular Forces & Boiling Points Slide 14 / 92

Intermolecular Forces

The attractions between molecules, intermolecular forces, are not nearly as strong as the intramolecular attractions that hold compounds together. They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

H Cl H Cl Intermolecular attraction ( week) Covalent bond (strong)

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There are three types of Intermolecular Forces: they are sometimes called van der Waals Forces · Dipole-dipole interactions · Hydrogen bonding · London dispersion forces

Intermolecular Forces

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Dipole-Dipole Interactions

· Molecules that have permanent dipoles are attracted to each other. · The positive end of one is attracted to the negative end

• f the other and vice-versa.

· These forces are only important when the molecules are close to each other.

+

• +
• +
• +
• +
• +
• The interaction between any

two opposite charges is attractive ( red) The interaction between any two like charges is repulsive (black)

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The polarity of a molecule is measured by its dipole moment, m. The more polar the molecule, the greater its dipole moment. The more polar the molecule, the higher its boiling point. That's because the attraction between the dipoles holds the molecules together, not letting them boil away.

Dipole-Dipole Interactions

Substance Molecular Dipole Boiling Weight (amu) Moment u( D) Poi nt ( K) Acetonitrile, CH3CN 41 3.9 355 Acetaldehyde, CH3CHO 44 2.7 294 Methyl chloride, CH3Cl 50 1.9 249 Dimethyl ether, CH3OCH3 46 1.3 248 Propane, CH3CH2CH3 44 0.1 231

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6 Which of the below molecules will have the highest boiling point?

A CH3CH2CH3 B

CH3OCH3

Substance Molecular Wt. Dipole Moment CH3CH2CH3 44 0.1 CH3OCH3 46 1.3 CH3Cl 50 1.9 CH3CHO 44 2.7 CH3CN 41 3.9

C CH3Cl D CH3CHO E CH3CN

SLIDE 4

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7 Which of the below molecules will have the lowest boiling point?

A CH3CH2CH3 B

CH3OCH3

Substance Molecular Wt. Dipole Moment CH3CH2CH3 44 0.1 CH3OCH3 46 1.3 CH3Cl 50 1.9 CH3CHO 44 2.7 CH3CN 41 3.9

C CH3Cl D CH3CHO E CH3CN

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London Dispersion Forces

London Dispersion Forces occur between all molecules. They result from the fact that electrons are in constant motion and sometimes are the same side of the molecule. When they are on one side, the molecule is polarized: one side is negative and the other is positive; the molecule acts like a dipole. That creates an electric field that oppositely polarizes nearby molecules...leading to an attraction. Let's see how that works using Helium as an example.

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London Dispersion Forces

While the electrons in helium repel each other, they

• ccasionally wind up on the same side of the atom.

At that instant, the helium atom is polar, with an excess

• f electrons on one side and a shortage on the other.

Another helium atom nearby, has a dipole induced in it, as the electrons on the left side of the first atom repel the electrons in the second. London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

2+ e- e-

Helium atom

δ

+

δ

• 2+

e- e- 2+ e- e-

Helium atom 1 Helium atom 2 electrostatic attraction

δ- δ+ δ

+

δ

• Slide 22 / 92

·

These forces are present in all molecules, whether they are polar or nonpolar.

·

The tendency of an electron cloud to distort in this way is called polarizability.

·

The larger the molecule, the more polarizable it is...and the stronger the London Dispersion Force.

·

That means, the higher the molecular weight of a molecule, the more London Dispersion Force it experiences.

London Dispersion Forces

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· The strength of dispersion forces tends to increase with increased molecular weight. · Larger atoms and molecules have larger electron clouds which are easier to polarize.

London Dispersion Forces

Halogen Molecular Weight ( amu) Boiling Point (K) Noble gas Molecula rWeight (amu) Boiling point (K) F2 38.0 85.1 He 4.0 4.6 Cl2 71.0 238.6 Ne 20.2 27.3 Br2 159.8 332.0 Ar 39.9 87.5 I2 253.8 457.6 Kr 83.8 120.9 Xe 131.3 166.1

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8 Which of the molecules below will have the highest boiling point?

A F2 B

Cl2

C Br2

D I2

SLIDE 5

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9

What intermolecular force is responsible for ice being less dense than liquid water?

A

London Dispersion Forces

B

Dipole-Dipole Forces

C

Ion-Dipole Forces

D

Hydrogen Bonding

E

Ionic Bonding

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10 Intermolecular force(s) responsible for the fact that

CH4 has the lowest boiling point in the set CH4, SiH4, GeH4, SnH4 is/are __________.

A

Hydrogen Bonding

B

Dipole-Dipole Interactions

C

London Dispersion Forces

D

Mainly H2 Bonding with some dipole-dipole interactions

E

mainly London Dispersion Forces with dipole-dipole interactions

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11 Which of the below molecules will have the lowest boiling point?

A F2 B

Cl2

C Br2

D I2

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12 Which of the below molecules will have the highest boiling point?

A He B Ne C

Ar

D Kr

E Xe

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13 Which of the below molecules will have the lowest boiling point?

A He B Ne C

Ar

D Kr

E Xe

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Which Have a Greater Effect?

Dipole-Dipole Interactions or Dispersion Forces

· If two polar molecules are of comparable size, dipole-dipole interactions are the dominating force. · If one molecule is much larger than another, dispersion forces will likely determine its physical properties. · If molecules are nonpolar, dispersion forces will dominate, since all molecules experience dispersion forces.

SLIDE 6

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Hydrogen Bonding

In these graphs, boiling points increase with the mass of the molecules...as we'd expect: larger dispersion forces due to larger masses. · The nonpolar series (CH4 to SnH4 ) follow the expected trend. · The polar series follows the trend from H2Te through H2S, · But water defies the trend. It is the lightest in its series, but has the highest boliing point.

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· The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. · We call these interactions hydrogen bonds. · These powerful bonds occur only when H is bonded directly to N, O or F.

Hydrogen Bonding

H H H H H N O

H H H H O O

H N O H H H H

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Hydrogen bonding arises in part from the high electronegativity and small radius of nitrogen,

• xygen, and fluorine.

When hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

Hydrogen Bonding

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Ice is the only solid that floats in its liquid form. If it didn't, life on Earth would be very different. For instance, lakes would freeze from the bottom and fish couldn't survive winters. Hydrogen bonding creates the space in ice that explains its low density.

Hydrogen Bonding

Water-Ice transition simulation.htm

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A

HF

B

HCl

C

HBr

D

HI E

All of the above.

14 Which of the following molecules has hydrogen bonding as one of its IM forces?

Slide 36 / 92 Ion-Dipole Interactions

· There is a fourth molecular force that will be important as we explore solutions later this year: Ion-dipole interactions are not considered a van der Waals force. · The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

+

• +
• +
• +
• _

Anion- dipole attractions

+

• +
• +
• +
• +

Cation -dipole attractions

SLIDE 7

Slide 37 / 92 Summarizing Intermolecular Forces

Interacting molecules or ions Are polar molecules involved? Are ions involved? Are polar molecules and ions both present?

Are hydrogen atoms bonded to N,O,or F atoms?

Ionic bonding NaCl, KI Ion- dipole forces NaCl in H2O Hydrogen bonding H2O, NH3 Dipole-dipole forces H2S, CH3Cl

Dispersion forces only (induced dipoles) Ar, I2 YES YES NO NO YES YES NO NO

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15 Which of the following molecules has London dispersion as its only IM force?

A

PH3

B

H2S

C

HCl

D

SiH4

E

None of the above.

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16 How many of these substances will have dipole-dipole interactions? (How many are polar molecules?) H2O CO2 CH4 NH3 A B

1

C 2

D 3 E 4

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17 Which of the following molecules will have the highest boiling point?

A

H2O

B

CO2

C

CH4

D

NH3

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18 Which liquid will have the lowest freezing point?

A

pure H2O

B

• aq. 0.60 M glucose

C

• aq. 0.60 M sucrose

D

• aq. 0.24 M FeI3

E

• aq. 0.50 M KF

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19 Of the following diatomic molecules, which has the highest boiling point?

A

N2 B

Br2

C

H2

D

Cl2

E

O2

SLIDE 8

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20 Of the following diatomic molecules, which has the lowest boiling point?

A

N2 B

Br2

C

H2

D

Cl2

E

O2

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21 Which one of the following derivatives of methane (CH4) has the lowest boiling point?

A

CBr4

B

CF4

C

CCl4

D

CI4

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22 Which one of the following derivatives of methane (CH4) has the highest boiling point?

A

CBr4

B

CF4

C

CCl4

D

CI4

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23 For an aqueous solution of a nonvolatile compound, the vapor pressure will be _______, the boiling point will be ________, and the freezing point will be _______ than pure water.

A

lower, lower, lower

B

lower, higher, lower

C

lower, higher, higher

D

higher, higher, lower

E

higher, lower, higher

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Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.

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Viscosity

· Resistance of a liquid to flow is called viscosity. · It is related to the ease with which molecules can move past each other. · Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Substance Formula Viscosity ( kg/m-s) Hexane CH3CH2CH2CH2CH2CH3 3.26 x 10-4 Heptane CH3CH2CH2CH2CH2CH2CH3 4.09 x 10-4 Octane CH3CH2CH2CH2CH2CH2CH2CH3 5.42 x 10-4 Nonane CH3CH2CH2CH2CH2CH2CH2CH2CH3 7.11 x 10-4 Decane CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3 1.42 x 10-3 http://upload.wikimedia.org/wikipedia/commons/b/b7/Viscosity.gif

SLIDE 9

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Surface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

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Phase Changes

Vaporization Condensation Sublimation Melting Freezing Deposition GAS SOLID

Energy of system

LIQUID

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Energy Changes Associated with Changes of State

Chemical and physical changes are usually accompanied by changes in energy. · When energy is put into the system, the process is called endothermic. · When energy is released by the system, the process is called exothermic.

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24 Which of the following is not a phase change?

A vaporization B

effusion C melting D sublimation

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25 Of the following, _____________ is an exothermic process?

A

melting

B

subliming

C

freezing

D

boiling

E

all are exothermic

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26 The first molecules to evaporate from a liquid

are__________.

A

Those with the lowest KE

B

Those farthest from the surface of the liquid

C

Those with the highest KE

SLIDE 10

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27 The direct change of a substance from a solid to a

gas is called ___________.

A

boiling

B

evaporation

C

sublimation

D

condensation

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Boiling vs. Evaporation

Boiling and evaporation are two ways in which a liquid can vaporize into a gas. However, there are important distinctions between these processes. Boiling Evaporation

• ccurs at a specific

temperature, the boiling point (B.P.)

• ccurs below the boiling point
• ccurs throughout the entire

liquid

• ccurs only at the surface of a

liquid achieved when atmospheric pressure equals vapor pressure (Patm = Pvap)

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Vapor Pressure

· At any temperature some molecules in a liquid have enough energy to escape. · As the temperature rises, the fraction of molecules that have enough energy to escape increases.

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Vapor Pressure

As more molecules escape the liquid, the pressure they exert increases. The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

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· The boiling point of a liquid is the temperature at which it’s vapor pressure equals atmospheric pressure. · The normal boiling point is the temperature at which its vapor pressure is 760 torr. (AKA 760 mm Hg = 1 atm)

Vapor Pressure

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28 When an aqueous salt solution is compared to

water, the salt solution will have

A

a higher boiling point, a lower freezing point, and a lower vapor pressure

B

a higher boiling point, a higher freezing point, and a lower vapor pressure.

C

a higher boiling point, a higher freezing point, and a higher vapor pressure

D

a lower boiling point, a lower freezing point, and a lower vapor pressure

E

a lower boiling point, a higher freezing point, and a higher vapor pressure

SLIDE 11

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29 What is the normal boiling point of ethanol?

A 45 B 55 C 35

D

78.3 E 80

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30 What is the boiling point ( in 0 C) of diethyl ether at 200 torr?

A

• 10

B

C 760

D 35

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31 What is the boiling point of water at 300 torr?

A 200 B

50

C 100 D

75

E

90

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Vapor Pressure

A liquid will boil when its vapor pressure equals atmospheric

• pressure. A pressure cooker works by increasing the

"atmospheric" pressure inside it, so water will not boil at 100oC; instead, it may be heated up to 120oC before turning to steam. Raising the cooking temperature cuts cooking time drastically.

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http://www.washingtonpost.com/wp-dyn/content/graphic/2007/02/20/GR2007022000797.html

Vapor Pressure

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Vapor Pressure

Recall that boiling occurs when Pvap = Patm. Since atmospheric pressure is so low at high altitudes, (e.g. top of Mount Everest) water will boil at a much lower temperature than here at Bergen Tech.

Patm = 33 kPa on Mt. Everest Patm = 101.3 kPa at sea level

SLIDE 12

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32 It will take longer to hard-boil an egg

A

• n top of Mt. Everest

B here at Bergen Tech C cooking times are equal

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33 A volatile liquid is one that _________.

A

is highly flammable

B

is highly viscous

C

is highly hydrogen-bonded

D

is highly cohesive

E

readily evaporates

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34 If a liquid is sealed in a container and kept at

constant temperature, how does its vapor pressure change over time?

A

it rises at first, then remains constant

B

it rises as first, then falls

C

it remains constant

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35 Based on this figure, the boiling point of diethyl ether under an external pressure of 1.32 atm is _______°C. A 10 B 20 C 30 D 40 E

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Phase Diagrams

Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

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Volatility

The more volatile a liquid: · The more quickly it evaporates · The higher its vapor pressure at a given temperature · The weaker its Intermolecular Forces

SLIDE 13

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· The circled line is the liquid-vapor interface. · It starts at the triple point (T), the point at which all three states are in equilibrium.

Phase Diagrams

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Phase Diagrams

It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

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Phase Diagrams

· The circled line in the diagram below is the interface between liquid and solid. · The melting point at each pressure can be found along this line.

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Phase Diagrams

· Below the triple point the substance cannot exist in the liquid state. · Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

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Phase Diagram of Water

· Note the high critical temperature and critical pressure. · These are due to the strong van der Waals forces between water molecules. · The slope of the solid-liquid line is negative. · This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

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Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures. The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (like caffeine).

SLIDE 14

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· The slope of the solid-liquid line is negative. · If you increase the pressure at a given temperature, near the freezing point, the ice will melt. · Water is the only known substance with this behavior. · The slope of the solid-liquid line is positive, as it is for most of the substances. · if you increase the pressure at a given temperature near -550C, the liquid will freeze. · Note that, since the Triple point is at a high pressure, you will see CO2 only subliming under normal atmospheric condition.

Phase Diagram of H2O and CO2

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36 The phase diagram of a substance is shown to the

• right. The region that corresponds to the solid phase

is _____.

A

w

B

x

C

y

D

x

E

x and y

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37 In this phase diagram, the substance is a

__________ at 25 °C and 1.0 atm

A

solid

B

liquid

C

gas

D

supercritical fluid

E

crystal

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38 On a phase diagram, the melting point is the same

as the __________

A

triple point

B

critical point

C

freezing point

D

boiling point

E

vapor-pressure curve

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39 On the phase diagram shown to the right, segment

__________ corresponds to the conditions of temperature and pressure under which the solid and the gas of the substance are in equilibrium

A

AB

B

AC

C

AD

D

CD

E

BC

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Solids

Crystalline, in which particles are in highly

• rdered arrangement.

We can think of solids as falling into two groups

Amorphous, in which there is no particular order in the arrangement of particles.

SLIDE 15

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Solids

Some examples of amorphous solids are: rubber, glass, paraffin wax and cotton candy. Crystalline solids include ionic compounds, metals and another group called covalent-network solids. Crystalline solids are categorized by bonding type as shown on the next slide.

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Types of Bonding in Crystalline Solids

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40 In a saturated solution of salt water, __________.

A

the rate of crystallization > the rate of solution

B

the rate of solution > the rate of crystallization

C

seed crystal addition may cause massive crystallization

D

rate of solution = rate of crystallization

E

addition of more water causes massive crystallization

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Covalent-Network Solids

· Diamonds are an example of a covalent-network solid, in which carbon atoms are covalently bonded to four other carbon atoms. · They tend to be hard and have high melting points.

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· Graphite is another example of a covalent-network

• solid. Each carbon atom is covalently bonded to 3
• thers in layers of interconnected hexagonal rings.

· The layers are held together by weak dispersion

• forces. The layers slide easily across one another, so

graphite is used as a lubricant as well as the "lead" in pencils.

Covalent-Network Solids

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Metallic Solids

· Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. · In metals, valence electrons are delocalized throughout the

• solid. This means that the

"sea" of electrons moves freely around all the nuclei.

SLIDE 16

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Properties of Metallic Solids

The delocalized nature

• f the electrons in

metals accounts for many physical properties. For example, metals are generally: · good conductors of heat and electricity · malleable · ductile, (i.e. may be drawn into wires)

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