Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

Подождите немного. Документ загружается.

FIGURE 11.8

Interaction of the oppositely charged

poles of water molecules makes water

a liquid at room temperature.

Solution

Comparing the 12-carbon compounds, we note that dodecane has a straight-chain

structure, whereas 2,2,4,6,6–pentamethylheptane is highly “branched,” with many

of the atoms hidden from other molecules, so dodecane will have the higher boiling

point of the two. Pentadecane is a longer straight-chain molecule than dodecane, so

its boiling point will be the highest among the group. The actual boiling points of

the compounds are (from lowest to highest)

177.8°C 2,2,4,6,6-Pentamethylheptane

216.3°C Dodecane

270.6°C Pentadecane

Further Insights

Other factors affect the boiling (and melting) points of organic chemicals. One is

the way in which molecules ﬁt together. Molecules that only have single bonds be-

tween carbon atoms ﬁt together nicely, leading to higher melting and boiling points

than those that have double or triple bonds. The presence of other types of atoms

can also substantially affect boiling points.

PRACTICE 11.1

Propane (C

) is a gas at room temperature, hexane (C

) is a liquid, and

dodecane (C

) is a solid. Explain why these three molecules have their speciﬁc

properties.

Propane C

Hexane C

Dodecane C

See Problems 11, 25, 26, and 27a.

Permanent Dipole–Dipole Forces

The forces that give strength to the interactions among water molecules are a re-

sult of the permanent dipoles that exist in water. Water molecules organize to

maximize the energetic stability gained by attractions and to minimize the re-

pulsions, as shown in Figure 11.8. Dipole–dipole interactions are relatively

weak compared to covalent bonds, but they become stronger in molecules

with relatively large dipole moments. However, even permanent dipoles in a

molecule may not lead to signiﬁcant intermolecular bonding at normal condi-

tions, as with the relatively small molecules HCl and H

S, which have boiling

points of only –85°C and –61°C.

Hydrogen Bonds

One very special type of interaction, the hydrogen bond, is of great importance.

Knowing about

hydrogen bonds, a term ﬁrst used in 1912, furthers our under-

standing of why water is a liquid at room temperature. This knowledge also helps

us explain some very elegant and important ideas about protein and DNA struc-

ture (Chapter 22) that add to our collective insight into human biology.

448 Chapter 11 The Chemistry of Water and the Nature of Liquids

–

Visualization: Intermolecular

Forces: Dipole–Dipole Forces

Visualization: Intermolecular

Forces: Hydrogen Bonding

Forces

11.2 A Closer Look at Intermolecular Forces 449

Hydrogen sulfide Hydrogen chloride

Dipole moments

–270

–250

–230

–210

–190

–170

–150

–130

–110

–90

–70

–50

–30

–10

110

Boiling point (˚C)

Molecular mass (amu)

10 30 50 70 90 110 130

Group VIIA

hydrides

Group VA

hydrides

Group VIA

hydrides

AsH

HCl

HBr

Group IVA

hydrides

SiH

GeH

SnH

FIGURE 11.9

Hydrogen bonds keep water liquid at

room temperature. In fact, hydrogen

bonds are breaking and re-forming at

an incredible rate.

Judging on the basis of its low molar mass alone, we would never have as-

sumed that water is a liquid at room temperature. However, water isn’t unique in

this regard. Other molecules, such as some of those shown in Figure 11.9, exhibit

boiling points higher than their molar masses would lead us to expect. Why are

their boiling points higher than expected? Each possesses hydrogen bonds.

A hydrogen bond is largely a dipole–dipole interaction of unusual strength

compared to other intermolecular forces. It is formed when a hydrogen nucleus

(a proton) is shared between two highly electronegative atoms of oxygen, nitro-

gen, or ﬂuorine. These electronegative atoms interact with the proton through

their available lone pair (or pairs) of electrons. In essence, the electronegative

atoms play tug-of-war with the proton, creating a very strong intermolecular

force of attraction. We can illustrate the hydrogen bond by drawing a series of

dots from one of the electronegative atoms to the hydrogen.

Recent data have shown that hydrogen bonds are at least partly—as much as

10%—covalent in nature. Why is a hydrogen bond so strong? The partial positive

charge on the tiny hydrogen atom gives it a relatively large charge-to-size ratio.

That helps it pack a large attractive punch. This permits a strong interaction with

HOHO

Water molecules play tug-of-war

with a proton.

Tutorial: Hydrogen Bonding

Ethane Ethanol Ethylene glycol

OH CH

OHHO

450 Chapter 11 The Chemistry of Water and the Nature of Liquids

Hydrogen

bond

FIGURE 11.10

Compounds such as water, ammonia,

and hydrogen ﬂuoride have much

higher boiling points than compounds

of similar size, as a consequence of

hydrogen bonding.

the highly electronegative ﬂuorine, oxygen, or nitrogen atoms. How strong are

hydrogen bonds compared to covalent bonds? The O···H hydrogen bond in liq-

uid water has a measured bond energy of about 23 kJ/mol, compared to the aver-

age O—H covalent bond energy of 470 kJ/mol, which means that this hydrogen

bond is only about 5% as strong as the intramolecular O—H bond, but stronger

than the 9 kJ required to vaporize a mole of methane.

What does our understanding of hydrogen bonding tell us about the properties

of water?

Hydrogen bonds between neighboring water molecules, as shown in

Figure 11.10, keep water in the liquid state at STP. Hydrogen bonds break and

re-form billions of times per second, but on average, enough hydrogen bonds

exist at any one time to keep water liquid. This process of bonds breaking and

re-forming means that a particular set of three atoms, HOH, will not stay to-

gether very long. Covalent bonds become hydrogen bonds, and vice versa. A

ﬂurry of activity is going on at the atomic level, even in a glass of water resting on

a tabletop.

EXERCISE 11.2 Comparing the Boiling Points

Ethane (C

) is an important starting material for the industrial production of

polyethylene plastics, used in items such as soft drink bottles. Ethanol (C

second only to water as an industrial solvent, is used in the synthesis of other

compounds and in some blends of gasoline. Ethylene glycol (C

) is the main

component in conventional automobile antifreeze. Judging on the basis of their

structures, arrange these compounds from lowest to highest boiling point.

First Thoughts

Hydrogen bonds have a substantial impact on boiling point, especially when the

molecule itself is small, with many hydrogen-bonding sites. Molecules that have

—OH groups are likely candidates for hydrogen-bonding interactions.

OCH

Ethylene glycol hydrogen bonding

HHCH

OCH

H H H H

Ethanol hydrogen bonding

OCH

Solution

Ethylene glycol has two polar sites at which hydrogen bonds can form, compared to

one on ethanol and none on ethane. Ethane therefore has the lowest boiling point,

–88.6°C, ethanol is next at 78.5°C, and ethylene glycol boils at 198°C.

Further Insights

An important idea in chemistry is that small changes in structure can lead to large

changes in properties. Having an alcohol group (—OH, see Chapter 12) substitute

for a hydrogen atom increases the boiling points of the compounds by over 100°C

for each such substitution. The effect is less in larger compounds. Why is this so?

Think about the types of intermolecular interactions that become possible.

11.3 Impact of Intermolecular Forces on the Physical Properties of Water, I 451

PRACTICE 11.2

Recent blends of “Green” (environmentally benign) antifreeze solutions use propy-

lene glycol (C

) which is far less toxic than ethylene glycol. Is its boiling point

higher or lower than that of ethylene glycol?

See Problems 12, 19–23, 27b, and 103.

11.3 Impact of Intermolecular Forces on

the Physical Properties of Water, I

A First Look at Phase Changes

A glass of pure water resting on a tabletop seems to be just that: at rest. But on a

molecular level, the system is deliriously active. The water molecules are moving

randomly and at great speed. Intermolecular hydrogen bonds are being broken

and re-formed at an incredible rate. Water molecules on the surface of the liquid

have enough energy to break free of their hydrogen bonds and go into the vapor

Propylene glycol

OHHO CH

Video Lesson: Properties

of Liquids

452 Chapter 11 The Chemistry of Water and the Nature of Liquids

Evaporation of water from the surface of the liquid.

FIGURE 11.11

How will the behavior of the liquid and

vapor change if we use a stoppered

(sealed) ﬂask?

Time Time Time

PPPP

FIGURE 11.12

In a sealed ﬂask, the system will come

to equilibrium, with the rate of evapora-

tion being equal to the rate of

condensation.

state. With higher temperature comes increased average kinetic energy of the sys-

tem. That means greater molecular motion, more breaking of hydrogen bonds,

and more molecules leaving the surface in the process we call

evaporation.Some

of the surrounding molecules return from the vapor to the liquid, attaching to

the surface via hydrogen bonding, in the process of

condensation. The corre-

sponding escape of molecules from a solid (such as ice) to the vapor is called

sublimation, and their return from the vapor to the solid is known as deposition.

Vapor Pressure

If we were to let our glass of pure water sit long enough, more molecules would

evaporate than would condense, and we would be left with an empty glass.

Let’s change the set-up by enclosing the water in a sealed ﬂask, as shown in

Figure 11.11. How will the behavior of the system change?

Initially, those water molecules with sufﬁcient energy will break free of their

intermolecular hydrogen bonds and escape from the surface. The resulting

vapor will collide with air molecules in the ﬂask and with the walls of the con-

tainer many times each second, exerting a force per unit area that we measure as

pressure. Recall from Chapter 10 that a pressure of 760 torr equals 1 atm.

In our sealed ﬂask, as illustrated in Figure 11.12, the more vapor that exists,

the more collisions per second with the surroundings and the greater the pres-

sure. As the concentration of vapor builds up in the ﬂask, some of it will con-

dense. Eventually, the air in the ﬂask will hold as much vapor as it possibly can,

and, as might occur on a sultry summer afternoon, the air will become

saturated

Video Lesson: CIA

Demonstration: Boiling Water

at Reduced Pressure

11.3 Impact of Intermolecular Forces on the Physical Properties of Water, I 453

with water vapor. At this point, the rate of evaporation from the

surface of the water will equal the rate of condensation back to

the surface. When the rates of the forward (evaporation) and reverse

(condensation) processes are equal, they are said to be in

equilibrium.

We indicate that both evaporation and condensation can occur by

using double arrows.

As with every process we study, bond breaking and bond forming

are always vigorously occurring, even though to our eyes all seems

quiet. We often indicate that the molecular-level equilibrium is dy-

namic, rather than static, by calling it

dynamic equilibrium. The pressure exerted

by a vapor in equilibrium with its liquid is called the

equilibrium vapor pressure,

or, commonly, the

vapor pressure, of the liquid at a given temperature. The vapor

pressure of hot water at 80°C is 355.1 torr; that of cool water at 20°C is 17.5 torr;

this means that more water molecules are in the gas phase at relatively higher

temperatures than at lower temperatures. The vapor pressure of a liquid increases

with temperature. Figure 11.13 and Table 11.2 show this dependence for water.

Does this make sense at the molecular level? A higher system temperature

means higher average kinetic energy. Water molecules are moving faster than at

lower temperature, so more molecules have enough energy to break their hydro-

gen bonds and evaporate.

Vapor Pressure of

Water at Selected

Temperatures

Temperature Vapor

(°C) Pressure (torr)

0 4.58

10 9.21

15 12.8

20 17.5

21 18.7

22 19.8

23 21.1

24 22.4

25 23.7

26 25.2

30 31.8

35 41.2

40 55.3

50 92.5

60 149.4

70 233.7

80 355.1

90 525.8

Source: Chemical Rubber Company.

Handbook of Chemistry and Physics, 66th

ed.; CRC Press: Boca Raton, FL, 1986.

Water vaporWater

760

1520

2280

3040

3800

50 100 150

Temperature (˚C)

Vapor pressure (torr)

FIGURE 11.13

The vapor pressure of liquids is related to temperature, as shown for water.

The line on the graph indicates the boiling point of water at 1 atm.

EXERCISE 11.3 Comparing Vapor Pressures

Which substance, water or methanol (CH

OH), would have a lower vapor pressure

at 50°C? Explain your choice.

Solution

The molecule that has more extensive hydrogen bonding will have the lower vapor

pressure at a given temperature. The entire structure of water encourages hydrogen

bonding. Methanol can also hydrogen-bond, but it also has nearly nonpolar C—H

bonds. The vapor pressure of water is 93 torr at 50°C, compared to about 400 torr

for methanol.

TABLE 11.2

EXERCISE 11.4 An Implication of Vapor Pressure

Acetone (propanone, C

O), is shown below. It is a polar compound with a vari-

ety of industrial uses as a solvent and in the manufacture of plastics and pharma-

ceuticals. It occurs naturally in plants and animals. If acetone and water were both

accidentally spilled on a lab bench, which would be likely to completely evaporate

ﬁrst at 20°C?

First Thoughts

To “completely evaporate” means to fully change from the liquid to the gaseous

state. Even though evaporation is not an equilibrium process (that is, molecules are

leaving the liquid’s surface more rapidly than they are condensing back), we can use

the vapor pressure as a guide. What contributes to a high vapor pressure? A low-

molar-mass molecule, even a polar one, that cannot form hydrogen bonds to itself

will probably have a higher vapor pressure than that of water.

Solution

Acetone at 20°C has a vapor pressure of about 185 torr, much greater than water’s

17.5 torr. Acetone (boiling point = 57°C) will completely evaporate ﬁrst.

Further Insights

Looking a bit more deeply, we note that there are hydrogen bond donors, such as

water, that can use a hydrogen atom to engage in hydrogen bonding with, for exam-

ple, another water molecule. Water is also a hydrogen bond acceptor, because its

oxygen can accept a hydrogen bond from another water or, for example, from

ethanol. On the other hand, acetone is only a hydrogen bond acceptor (as in the

454 Chapter 11 The Chemistry of Water and the Nature of Liquids

PRACTICE 11.3

Name a substance that would have a lower vapor pressure at 50°C than either water

or methanol. Explain your reasoning.

See Problems 33–35 and 36a.

We now know that the ability to form hydrogen bonds is one important fac-

tor leading to lower vapor pressure at a given temperature, but it is not the only

factor. Molar mass and structure are also important because, as we have already

learned, London forces can be signiﬁcant in relatively large molecules. In order to

develop a vapor pressure of 400 torr, octane (C

) must be heated to 104°C.

The branched molecule 2,3,3–trimethylpentane, shown in Figure 11.14, has this

vapor pressure at 92.7°C, and water reaches it at 83.0°C.

2,3,3-Trimethylpentane

CH C

Octane

FIGURE 11.14

The London forces are greater in octane than in

the highly branched 2,3,3-trimethylpentane. Which

would require a higher temperature to reach a

vapor pressure of 400 torr?

present example), so it cannot internally hydrogen-bond. However, it can form such

bonds with water and is soluble in it. There are millions of organic (carbon-based)

compounds that have a vast range of properties. Chapter 12 will introduce the com-

pounds and chemistry that make organic chemistry well worth knowing.

PRACTICE 11.4

Compare acetone and hexane (C

) using the same lab bench scenario as in

Exercise 11.4.

See Problems 29 and 30.

HERE’S WHAT WE KNOW SO FAR

■

The vapor pressure of a liquid increases with temperature.

■

More hydrogen bonding leads to lower vapor pressure.

■

All other things being equal, heavier molecules have lower vapor pressures

than lighter molecules.

■

Straight-chain molecules have lower vapor pressures than their branched

isomers.

■

When several of these factors come into play, it is hard to predict which will

dominate. We then run experiments or look up information in data tables.

Boiling Point

Let’s return to our glass of pure water. As we heat it, the vapor pressure of the liq-

uid increases along with the temperature. If we are at sea level on a day when the

surrounding pressure is 1 atm, the liquid will start to bubble from within as the

temperature approaches 100°C. As it does so, the vapor pressure of the liquid will

edge ever closer to the atmospheric pressure. When the temperature reaches

100°C, bubbles burst forth throughout the water in a familiar phenomenon we

call

boiling. We discussed boiling earlier in the chapter, and you understand its

general meaning from all the years you have been boiling water to make tea or

cook vegetables. Now we are ready to look at it in more depth. Boiling is not just

a surface process like evaporation, because it involves the entire liquid. We deﬁne

the

boiling point as the temperature at which the pressure of the liquid’s vapor

(rather than the vapor pressure, which is deﬁned for an equilibrium process) is

equal to the surrounding pressure. If that pressure is 1 atm, at or near which so

many of life’s normal activities take place, the temperature at which a liquid boils

is called its

normal boiling point.

A good portion of the world’s population does not live at sea level, and for

them,“normal”is anything but. In mile-high Denver, the atmosphere is less dense

than at sea level, so the atmospheric pressure is correspondingly lower, about

0.82 atm (620 torr). If our glass of water were heated in Denver, it would boil at

about 95°C. The difference in boiling point with pressure is even more dramatic

in Mexico City, at 2240 m (7340 ft), where the atmospheric pressure is only

0.76 atm (580 torr). The boiling point at that altitude is only about 90°C (194°F).

Figure 11.15 shows the decrease in the boiling point of water as the altitude

increases and atmospheric pressure consequently decreases. Food manufacturers

take advantage of the increase in boiling point with pressure when they process

foods by canning them at high pressure, allowing the food item to be heated to a

relatively high temperature, typically 107°C, for at least 3 minutes, killing any

bacteria within. A quick way to cook soup is to use a pressure cooker, which in-

creases the pressure within from 1 to 2 atm, raising the boiling point of the soup

11.3 Impact of Intermolecular Forces on the Physical Properties of Water, I 455

Application

Water boils at 100°C at 1.0 atm. At this

temperature, the vapor pressure equals

the atmospheric pressure.

Visualization: Boiling Water

with lce Water

Video Lesson: Vapor Pressure

and Boiling Point

to about 120°C (250°F). Commercial cake mixes often have two sets of instruc-

tions: one for cooking the batter at sea level (normal boiling point) and one for

higher altitudes.

Heating Curves

We opened this chapter by looking at drinking water in the Tampa Bay area. Now

let’s look north, perhaps above the Arctic Circle, to a group of hikers who want to

get drinking and washing water by melting some ice and then purifying it by boil-

ing. What happens as they heat a 10.0-kg block of the ice at –10.0°C? To fully un-

derstand the process that occurs, we must consider four changes: (1) warming the

ice; (2) melting the ice; (3) heating the water, and (4) boiling the water. As we pro-

ceed, think about what happens to the energy we add, the molecular motion, and

the resulting system temperature. We will assume that all of the heat goes into the

ice (the system), not into the surroundings.

Change 1: Warming the Ice

The temperature of our ice is –10.0°C. As we add heat, the molecules that are

fairly rigidly held in place begin to move a bit more. We still have ice, but the

average energy has increased, so the temperature rises. Figure 11.16 displays a

heating curve, a plot of the temperature of a compound versus time as it is heated.

The plot indicates the speciﬁc temperature ranges for solid, liquid, and gas

phases.

Theheatneededtowarm1gofasubstance enough so that its temperature

rises 1.0°C is its

speciﬁc heat (see Section 5.4), which for ice is equal to 2.05 J/g·°C.

Raising the temperature of our 10.0-kg block of ice from –10.0°C to its

melting point, at which it changes from a solid to a liquid, requires that 205 kJ of

heat, q, be added to the system.

q = speciﬁc heat × mass × change in temperature

q =

2.05 J

g·

◦

× 10,000 g × (0.0° C −

−

10.0°C)

= 205,000 J

= 205 KJ

This much heat would be supplied by, for example, burning about 4 g of methane

(natural gas is typically over 90% methane) in air.

Change 2: Melting the Ice

Our ice is now at its melting point of 0.0°C. As we add more heat, water molecules

move more freely and randomly, beginning the transformation from the solid

456 Chapter 11 The Chemistry of Water and the Nature of Liquids

Directions for baking a cake at high

elevations.

0 1000

Miami,

Florida

Denver,

Colorado

Mexico

City

La Paz,

Bolivia

2000 3000 4000 5000

100

105

Altitude (m)

Boiling point of water (˚C)

FIGURE 11.15

The boiling point of water goes down

as the altitude rises. This is because air

pressure is lower at higher altitude

(remember the deﬁnition of boiling

point).

Temperature (°C)

100

110

–10

Ice

Water

Water and steam

Steam

Ice

and

water

0 5,000 10,000 15,000 20,000 25,000 30,000

Heat added (kJ)

crystal to the liquid state. The temperature at this change of state, the melting

point, is constant (Figure 11.16), because the added heat is going into breaking

hydrogen bonds, rather than increasing the average kinetic energy of the water

molecules. The amount of heat needed to convert the solid to liquid at the melt-

ing point at constant pressure is called the

heat of fusion (

fus

H), which for water

is 334 J/g, or 6.01 kJ/mol. Melting is an endothermic process, which means that

freezing (the opposite process) is an exothermic process. Heat is released when a

liquid freezes.

How much heat is needed to melt 10.0 kg of ice? The calculation shown below

indicates that we need to add 3340 kJ to the system.

q = mass × ∆

fus

10,000 g ice ×

334 J

gice

= 3,340,000 J = 3340 kJ

This is about the amount of heat supplied by burning about 60 g of methane in

air.

Change 3: Heating the Liquid

The heat going into the (pure liquid) water increases molecular motion, raising

the temperature of the liquid (Figure 11.16). The heat needed to raise the tem-

perature of our 10.0-kg water sample from 0.0°C to 100.0°C can be calculated

using the speciﬁc heat of water, 4.184 J/g·°C.

q = speciﬁc heat × mass × change in temperature

4.184J

g·

◦

× 10,000 g × (100.0°C − 0.0°C)

= 4,180,000 J = 4180 kJ

The total of 4180 kJ can be supplied by about 75 g of methane burning in air.

Change 4: The Boiling Point

As we continue to heat the water, it boils as the liquid is converted to vapor at

constant temperature (Figure 11.16). Analogous to the heat of fusion for melting

is the

heat of vaporization (

vap

H) for converting the liquid water to vapor at the

normal boiling point. The value of

2.44 kJ

g water

is over 7 times that of the heat of

fusion



0.334 kJ

gice



. This indicates that a lot more energy is required to overcome

11.3 Impact of Intermolecular Forces on the Physical Properties of Water, I 457

FIGURE 11.16

As water is heated, it goes through changes from solid

to liquid to gas. Why are some regions sloped and oth-

ers ﬂat? Why are areas of constant temperature of

different lengths?

Visualization: Changes of State

Video Lesson: CIA

Demonstration: Boiling Water in

a Paper Cup