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

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the intermolecular forces when boiling water than when melting ice. Boiling

causes all of the remaining hydrogen bonds keeping the water as a liquid to be

broken. The heat needed to boil 10.0 kg of water is 24,400 kJ.

q = mass ×



vap

10,000 g ×

2.44 kJ

g water

= 24,400 kJ

This amount of heat can be supplied by burning about 440 g of methane, over 3

times the total of the previous changes! The heat required from start to ﬁnish is

equal to the sum of the amounts of heat required for all the change along the way.

Total heat = q

ice warming

+ q

ice melting

+ q

water warming

+ q

water boiling

= 205 kJ + 3340 kJ + 4180 kJ + 24,400 kJ

= 32,100 kJ

It is possible to heat the water vapor to quite a high temperature, and the amount

of heat needed is related to the speciﬁc heat of water vapor, 1.84 J/g

·°C. This is

shown in Figure 11.16.

EXERCISE 11.5 Energy Changes Upon Heating the Water

Determine the heat required to raise the temperature of a 2.50 × 10

g block of ice

(equivalent to about a cup of water) in a 425 g aluminum pot from 0.0°C to room

temperature, 20.0°C. Roughly how many grams of methane must be burned in air

to supply this much heat? The heat of combustion of methane is –890 kJ/mol. Solid

aluminum has a speciﬁc heat of 0.880 J/g·°C. Assume that all heat goes to the ice and

the aluminum pot.

First Thoughts

What processes occur? The input of heat will melt the ice at constant temperature

(change of state) and then warm the ice from 0.0°C to 20.0°C. The aluminum pot

will be warmed from 0.0°C to 20.0°C.

Solution

Total heat = q

melting ice

+ q

warming liquid

+ q

warming aluminum



334 J

gice

× (2.50 ×10

gice)





4.184 J

gice·

◦

× (2.50 ×10

gice)

×(20.0°C −0°C)





0.880 J

gAl·

◦

× 425 g Al ×(20.0°C −0°C)



= 83,500 J + 20,920 J + 7480 J = 111,900 J = 112 kJ

Grams of CH

required =112 kJ

1 mol CH

890 kJ

16.0gCH

1 mol CH

= 2.0gCH

Further Insights

Even though the speciﬁc heat of the aluminum pan is a lot less than that of water, it

is a heavy pan, and a lot of heat is required to raise its temperature. A more realistic

picture would need to take into account the heat loss, and lots of it, to the air.

PRACTICE 11.5

How much energy does it take to convert 130.0 g of ice at –40.0°C to steam at

160.0°C?

See Problems 31, 32, 37, and 38.

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

FIGURE 11.17

The moons of Jupiter as taken by NASA’s Galileo space probe.

11.4 Phase Diagrams

We now understand the changes that occur in water as we raise its temperature

from below freezing to the boiling point at normal pressure. However, much of

the known universe is nowhere close to our own “normal” pressure or tempera-

ture, and that affects the form water takes beyond our world. Figure 11.17 shows

recent photographs of the moons of Jupiter taken by NASA’s Galileo space probe.

Why do the moons look so different from one another, from their host planet,

and from anything else we’ve seen in the solar system? One reason apparently has

to do with ice—not the type of ice we ﬁnd in the refrigerator, but ice that exists

under pressures of tens of thousands of atmospheres. Under these immense pres-

sures, the molecules of ice that form large parts of the interiors of the “Jovian”

moons are highly distorted. Hydrogen bonds are squeezed and shifted. Atoms are

brought closer together and bond lengths shortened so that the ice becomes un-

usually dense. Temperature changes deep inside these moons also lead to changes

in the nature of the ice. The consequences of all these changes are that the moons

contract and expand, causing huge cracks and grooves. We add to our under-

standing of the solar system by looking at the changes in ice as temperature and

pressure change.

A graph showing the way the phase of a substance, or a mixture of substances,

is related to temperature and pressure is called a

phase diagram. The heating

curves we discussed previously describe how water changes with temperature at

a constant pressure. However, as with the changes in the moons of Jupiter, a great

many processes do not happen at constant pressure, and a phase diagram gives us

a more comprehensive picture of these changes. In the chemical industry, phase

diagrams have many uses, including determining the conditions required for

manufacturing ceramics from clays (see Chapter 13). The composition of homo-

geneous mixtures of metals, called alloys, can also be changed on the basis of the

information in phase diagrams.

The phase diagram that includes the various types of ice in the core of plan-

ets is fairly complex, so let’s work ﬁrst with a small portion of it, that for water at

more Earth-like temperatures and pressures, shown in Figure 11.18. Note in the

11.4 Phase Diagrams 459

Tutorial: Phase Diagrams

ﬁgure the three normal phases: solid, liquid, and gas (water vapor).

There are also bold lines, called

phase boundaries, at which two phases

exist in equilibrium. How does water change state as we heat it from

well below 0°C and at 1 atm (the dotted line on Figure 11.18)? Keep in

mind that these are the common changes that we experience every day

as we interact with water by boiling it to make tea or coffee or cause it

to melt when we add it as ice to a soda.

Our water sample begins as ice. We raise the temperature to 0°C,

where it arrives at a phase boundary at which ice and liquid water exist

in equilibrium. This is the normal melting (and freezing) point of

water. As we continue to raise the temperature, the water remains liq-

uid until 100°C, the boiling point, which is at the liquid/gas phase

boundary. Above 100°C, the water will always be a gas if the pressure is

kept constant at 1 atm. This is the same information we gleaned from

the heating curve of water, but a phase diagram gives us so much

more. We can tell, for example, that there is a

triple point of water (Fig-

ure 11.18) where all three phases exist in equilibrium. For water, this

occurs at 0.01°C and 4.6 torr. The triple points of elements and com-

pounds are useful because they serve as reference points with which temperature

scales for thermometers are deﬁned, and against which the highest-quality ther-

mometers are calibrated. If we follow the liquid/gas phase boundary (the curved

line segment AB, the boiling point) toward increased temperature and pressure,

we get to a point above which the substance can no longer be liqueﬁed and the

gas and liquid have the same density—that is, we have a

supercritical ﬂuid. This

point (B) is called the

critical point, and it is associated with a critical temperature.

The pressure at the critical temperature is the

critical pressure.

EXERCISE 11.6 Reading Phase Diagrams

What does the line segment AD on the phase diagram of water represent? What

does line AC indicate?

Solution

Line segment AD represents the point of equilibrium between solid and liquid—the

melting (freezing) point at which the solid and liquid exist in equilibrium. Line seg-

ment AC is point of equilibrium between solid and gas for water; sublimation and

deposition occur at equilibrium anywhere along this line.

PRACTICE 11.6

What is the approximate boiling point of water when the external pres-

sure is 0.70 atm? Does this agree with our previous discussion about boil-

ing points at high altitudes?

See Problems 39–48.

The phase diagram of CO

, shown in Figure 11.19, looks similar to that

of water, with a triple point, critical temperature and pressure, and sev-

eral phase boundaries. At 1 atm, CO

changes directly between the solid

and gas phases (sublimation/deposition). You might have seen dry ice,

which does not melt but, rather, sublimes directly from a solid into a

gas. At pressures greater than about 5 atm, CO

will liquefy from the

solid phase. This liquid is used for dry-cleaning clothing to replace

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

218

1.00

0.0060

0 0.0098 100

374

Solid

Liquid

Gas

Temperature (

°C)

Pressure (atm)

FIGURE 11.18

Phase diagram of water.

72.8

1.00

5.1

Solid

Liquid

Gas

Temperature (°C)

Pressure (atm)

–78 –56.6 31

FIGURE 11.19

Phase diagram of carbon dioxide.

Video Lesson: Phase Diagrams

Tetrachloroethylene

nonpolar, but environmentally hazardous, solvents such as tetrachloroethylene

). These solvents rid clothing of nonpolar grease and dirt.

Liquid CO

is an excellent and relatively benign reusable alternative. Clothes

are cleaned in CO

at 15°C and 50 atm pressure, which is along the liquid/gas

boundary in the CO

phase diagram. The clothes are kept close to room temper-

ature, and about 98% of the CO

is recycled. The critical temperature and pres-

sure of CO

are 304.2 K and 70.8 atm. Supercritical carbon dioxide, typically in

the range of 310 to 400 K and 75 to 100 atm, can be used to clean industrial

machinery.

11.4 Phase Diagrams 461

The moons of the outer planets make

most elegant chemical laboratories,

because conditions on these worlds are so very dif-

ferent from those here at home. When we look at

photos of Jupiter’s moons, Europa and Ganymede

(Figure 11.17), we see some vexing geological fea-

tures. How do we know what caused these features?

We can’t visit the moons, at least not directly. How-

ever, our space probes can gather various types of

electromagnetic radiation—including light, ultra-

violet radiation, and X-rays—that give us data

from which to draw conclusions.

The probes’ data seem to show that deep within

the surface of these moons, there are layers of ice,

each in a different phase, mixed with rocks. The phase

diagram of water that we showed as Figure 11.18

is inadequate to account for the low temperatures

and massive pressures found within these moons.

A more comprehensive phase diagram for water,

emphasizing its solid phases, is given as Figure 11.20.

The ice that is made in our kitchen freezer or found

in an ice storm—what we might call “normal ice”—

is technically called Ice Ih. Its structure is shown in

Figure 11.21. There are 11 other forms

of ice that have been made in the laboratory or

simulated by computer.

The phase diagram shows that as the

temperature and pressure inside the moons

change, different types of ice are formed.

Each ice has its own density. As layers

form versions of ice of different

densities, they form the geological

features, such as cracking and

grooves, that we see on the surface

of these incredible satellites. Part

of the task of scientists is to

develop phase diagrams that

tell us about the conditions at

which each type of ice exists.

How do we know?

The moons of Jupiter

0 100 200 300 400 500 600 700 800

Solid

IhIc

III

VII

VIII

Liquid

Vapor

Temperature (K)

Pressure (Pa)

FIGURE 11.20

This is an expanded phase diagram for water, taking into account

the phase changes among the various solid (ice) forms. These forms

of ice are thought to occur in layers beneath the surface of some

of the moons of the outer planets.

FIGURE 11.21

The arrangement of water

molecules within a crystal

of “normal” ice (Ice Ih).

Application

HEMICAL ENCOUNTERS:

as a Dry Cleaning Solvent

FIGURE 11.22

A water strider (Gerris remigis) on the surface

of a pond. Note that the water indents as it

supports the weight of the water strider.

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

11.5 Impact of Intermolecular Forces on

the Physical Properties of Water, II

Viscosity

Making chocolate at home is difﬁcult, because so much can go wrong, including

unintended air bubbles, cracking, and a greasy look and feel to the ﬁnal product.

Having these things occur on an industrial manufacturing scale can be especially

damaging. One of the most important properties that cause these and other in-

consistencies in chocolate is called

viscosity, which is a measure of resistance to

ﬂow. In the chocolate process industries, a viscometer is used to constantly moni-

tor the viscosity of fully melted chocolate so that it ﬂows into the molds and

produces a ﬁnal product that is consistent from batch to batch. Viscosity mea-

surements are also important to the fuel industry, where the ability of fuels to

ﬂow is important to consider when storing, pumping and, within engines, inject-

ing the fuel. Viscosity is reported in units of millipascal seconds (mPa·s). A highly

viscous liquid, ethylene glycol (C

), is combined with water to form com-

mercial antifreeze and has a viscosity of 16.1 mPa·s at 25°C and 1 atm. The water

with which it is combined has a viscosity of 0.890 mPa·s at the same conditions.

Why does the viscosity differ among liquids?

As with so many things related to chemistry, we gain understanding by look-

ing at behavior at the molecular level. The molecules that constitute viscous liq-

uids do not move well relative to each other.

Why is this so? Part of the answer has

to do with intermolecular forces. More such interactions tend to increase vis-

cosity. Size is also important. It is hard to move large molecules.And, as with boil-

ing point, molecular shape is relevant because molecules must be able to move

well among each other in order to ﬂow easily. Octane (C

) has a viscosity of

0.508 mPa·s at 25°C and 1 atm, whereas the viscosity of pentane (C

) mea-

sures 0.224 mPa·s at these conditions. Temperature is also a key factor because, as

we have seen, molecules with high average energy move relatively quickly and can

overcome intermolecular forces, leading to low viscosity. For example, the viscos-

ity of water changes from 0.890 mPa·s at 25°C to 0.378 mPa·s at 75°C. That’s why

chocolate, as well as oil in your car’s engine, ﬂows more smoothly at higher than

at lower temperatures.

Surface Tension

The water strider glides from place to place along the water’s surface, as shown in

Figure 11.22. Why can the insect travel on the water’s surface? The key is a prop-

erty of liquids called

surface tension. This property is also responsible for the

Motor oil viscosity is rated for use on

the basis of the type of engine and the

outside temperature.

beads of water that run down your window in the rain or across the surface of

your car after it has been waxed. Surface tension is a measure of the energy per

unit area on the surface of a liquid. To better understand what this measure

means, we note that hydrogen bonding in water makes it more energetically

stable. Water molecules on the interior of the solution can hydrogen-bond in all

directions, whereas surface molecules can form these bonds only toward the

solution itself. In order to maximize the number of hydrogen bonds (and, there-

fore, the energetic stability of the liquid), water forms a shape with the minimum

surface area, a sphere. This is why small droplets of water form beads and water

striders can glide on water.

The chemical industries pay close attention to surface tension in the design of

many consumer products. For example, soaps and detergents are designed to

have very low surface tension in order to interact well with clothing and dishes.

Surface tension is also considered in the design of pharmaceuticals and plays a

large role in how tablets and gelcaps dissolve after they are swallowed.

Capillary Action

A buret, shown in Figure 11.23, is an essential piece of glassware for measuring

volume when performing laboratory analyses with liquids. Water has been added

to the buret in Figure 11.24. Why does the water seem to rise up the sides of the

glass? And given this curvature, what do we say is the volume of water in the

buret? The answer to the ﬁrst question again lies in one of our chemical themes,

the impact of hydrogen bonding on the structure of water. The forces that keep

the large group of water molecules together are called

cohesive forces. Then there

is the glass in which the water sits. The glass is largely silicon dioxide (SiO

) but

with a surface containing many polar oxygen and hydrogen atoms. As is shown in

Figure 11.25, the hydrogen atoms on water interact with the oxygen atoms on the

surface of the glass. The net result is that some molecules of water adhere to

the glass, a process called,

adhesion. The effect in which water seems to crawl up

the sides of a thin tube such as a buret is called

capillary action. In this case, the ad-

hesive forces between water and the glass are stronger than the cohesive forces

within water itself. The weight of the water prevents it from rising up even higher

11.5 Impact of Intermolecular Forces on the Physical Properties of Water, II 463

FIGURE 11.23

A buret is essential labware

in the analytical laboratory.

FIGURE 11.24

The meniscus of water within

a buret is readily observable.

Water beads on a freshly waxed car.

FIGURE 11.25

The interaction of water with the surface of glass

results in adhesion. Hydrogen bonds between the

water molecules and the glass surface are illustrated.

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

Hexane in a buret.

within the buret. Mercury in a glass tube (such as in a thermometer) shows the

opposite behavior, bulging upward within the glass, in response to the domi-

nance of cohesive over adhesive forces.

Adhesive forces lead to the formation of a

meniscus, the concave shape at the

surface of the water in the buret. By convention, the volume of liquid delivered

from the buret is read at the bottom of the meniscus, as shown in Figure 11.24, so

our reading of the buret is 40.03 mL. Because burets are graduated from the top

(0 mL) down, there are 9.97 mL of water remaining in our 50.00-mL buret.

The balance between forces is not always easy to assess. For example, hexane

) shows a concave meniscus similar to that of water, though not quite so

pronounced.

Would we have predicted this fact? The surface of the glass can, pre-

sumably, induce a dipole in hexane, so we can explain the behavior. However,

making predictions when several effects are important is sometimes quite

challenging.

HERE’S WHAT WE KNOW SO FAR

■

A liquid ﬁlls its container from the bottom up and ﬂows easily.

■

Intermolecular forces of attraction determine the properties of a substance.

The number and strength of these forces contribute to the attractions between

adjacent molecules.

■

London forces are typically the weakest forces between molecules. Because

every molecule has the ability to become polarized, every molecule possesses

London forces. In very large molecules, these forces can be signiﬁcant.

■

Dipole–dipole interactions are much stronger than London forces. They exist

in molecules that contain a nonzero dipole moment, and their strength is

related to the strength of the dipole moment.

■

Hydrogen bonds are the strongest of the van der Waals forces. They exist pri-

marily between the interaction of a hydrogen atom and two atoms of F, O,

and/or N.

■

The strength of the intermolecular forces of attraction determines the shape of

a meniscus, the surface tension of the liquid, its vapor pressure, and other

physical properties associated with the liquid.

Buret

Meniscu

Water

Glass

11.6 Water: The Universal Solvent

In the introduction to this chapter, we raised three questions related to the pro-

jected water shortage in Tampa Bay: What is it about water that makes it unique

among molecules? How can we use our understanding of the nature of water to

explain its properties? And how can we use these properties to help in the search

for clean water in places like Tampa Bay? We have answered the ﬁrst question, fo-

cusing especially on the polar nature of water and its ability to form the special

interaction called a hydrogen bond. This enabled us to explain a variety of

properties of water, including its unusually high boiling point, its vapor pressure,

and its phase diagram, giving us the understanding to answer much of the second

question.

Implicit in the third question about the search for clean water is recognition

that the water found in nature is not pure. It is ﬁlled with all manner of

solutes,

substances that dissolve in it. Water is often called the

universal solvent because of

its ability to dissolve so many chemicals to form homogeneous mixtures called

solutions. Aqueous solutions (water-based solutions) can have solutes that are

gases, liquids, or solids. A special class of large-molar-mass molecules that ﬁnely

disperse in solution are called colloids.

Why Do So Many Substances Dissolve in Water?

We can answer the question that opens this section by looking at seawater, in

which dissolved sodium chloride is one of the primary solutes. Chemically, we

can understand what happens when salt dissolves in water by considering the

process of dissolving as a series of discrete steps, as shown in Figure 11.26. Al-

though the process does not, in fact, occur in steps, we may use this procedure to

enhance our understanding of the dissolution process. At each of the steps along

our imaginary solution-forming pathway, ask yourself, “Is this likely to be an

endothermic or an exothermic process?”

■

Solute separation: In the ﬁrst step, the salt must separate into sodium ions

and chloride ions. The electrostatic forces that keep the ions together must

be overcome, so energy is required. This part of the process is endothermic

(H =+).

11.6 Water: The Universal Solvent 465

This water contains many dissolved

salts, compounds, and molecules. It

also contains particles too big to

dissolve; these particles remain sus-

pended and make the water cloudy.

∆H

+ ∆H

∆H

soln

Solute

Separated solute Separated solvent

Solvent Solution

FIGURE 11.26

The processes necessary, along with

their energy exchanges, for the disso-

lution of a salt, such as NaCl, in water.

The overall change in energy can be +

or −, depending on the ion–dipole

interaction.

■

Solvent separation: Energy is required to break hydrogen bonds holding water

molecules together in order to accommodate the incoming ions from the

solute. This part of the process is endothermic (H =+).

■

Electrostatic interaction between the solvent and solute ions: Electrostatic at-

traction is the great stabilizer. As shown in Figure 11.27, sodium cations are

attracted to the negatively charged dipole at oxygen on each water molecule.

Conversely, chloride anions are attracted to the positively charged hydrogen

poles. Although there is constant ion movement in the solution, at any one

time it is typical to have up to six water molecules surrounding each ion. This

ion–dipole interaction contributes to the solution process in two ways. First,

for the process involving sodium chloride and water, the interaction is highly

exothermic (H =“–”). In order to have favorable formation of a solution,

the value of the exothermic interaction of the solute and solvent is typically

larger than the endothermic processes in the ﬁrst two steps. Second, dissolv-

ing salt ions in water results in dispersal of the cations and anions through-

out the solution. This ion–solvent interaction is called

solvation, and when

the solvent is water, it is known as

hydration. The heat of solution (

sol

H) can

be viewed as the sum of the heats of solute separation, solvent separation,

and ion–dipole interaction.



sol

H = 

solute

H + 

solvent

H + 

solvation

An overall exothermic heat of solution (–

sol

H) favors solution formation

but is not by itself an indicator of solubility. The additional dispersal of the ions

in the solvent (see Chapter 14) is also important. For example, NaCl has a value

of 

sol

H =+3 kJ/mol, yet it dissolves in water. So does potassium hydroxide,

KOH (

sol

H =–57.6 kJ/mol). In fact, the heat of solution of KOH is so large that

adding 100 g to a liter of water will raise its temperature to nearly boiling, with



sol

H = –103 kJ for the 100 g.

As a rule of thumb, like dissolves like. Polar molecules typically dissolve many

salts via ion–dipole interactions. Similarly, polar liquids mix as a consequence of

the stability gained from dipole–dipole and, when possible, hydrogen-bonding

interactions. Nonpolar substances, such as the industrial solvents benzene (C

)

and toluene (C

) are often miscible (mix in all proportions) with other nonpo-

lar substances. Part of the reason lies with London forces and part with the sta-

bility gained by dispersion as the liquids mix.

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

FIGURE 11.27

The interaction of the sodium ion with the

negative end of the dipole on water and

that of the chloride ion with the positive

end of the dipole on water contribute to the

energetic stability that permits NaCl to dis-

solve in water. Note in the electrostatic

potential maps how the water molecule’s

electron potential changes as it interacts

with the ions.

EXERCISE 11.7 Predicting Solubility

Which of the following substances is (are) soluble in water: ethanol (C

OH),

ethylene glycol (C

), and cyclohexane (C

First Thoughts

The ability to form electrostatic attractions to water, including hydrogen bonds, dic-

tates the solubility of a substance in water. Which of these compounds can have

such attractions?

Solution

Ethanol and ethylene glycol are inﬁnitely soluble in water. In fact, both can act as a

solvent or solute with water. Cyclohexane is not appreciably soluble in water.

Further Insights

11.6 Water: The Universal Solvent 467

Ethanol

COH

Ethylene glycol

COH

Cyclohexane

Both benzene (C

) and toluene (C

) are excellent solvents for cyclohexane.

Industrial manufacturing of polymer-based plastics often involves dissolving non-

polar solutes in nonpolar solvents.

PRACTICE 11.7

Which of the following substances is (are) soluble in hexane: water, carbon tetra-

chloride (CCl

), and iodine crystals (I

See Problems 58–60 and 62.

The maximum amount of a solute that can dissolve in a solvent is its solu-

bility

, often given in grams of solute per 100 mL of solvent (

g solute

100 mL solvent

). If we

added any more solute, it would not dissolve. At that point, we say that the solu-

tion is

saturated (compare this with the use of the term saturated for the vapor of

the pure liquid). Sometimes the solubility of a substance in water depends on the

relative size of the polar and nonpolar parts of a molecule of the substance.

Table 11.3 shows that the series of alcohols gets progressively less soluble in water

as the nonpolar end of each becomes larger. However, even substances that we

Solubility of Some Alcohol Compounds in Water

Compound Name Solubility (g/L)

OH Methyl alcohol (methanol) Inﬁnite

OH Ethyl alcohol (ethanol) Inﬁnite

OH Propyl alcohol (1-propanol) Inﬁnite

OH Butyl alcohol (1-butanol) 79

OH Pentyl alcohol (1-pentanol) 27

OH Hexyl alcohol (1-hexanol) 5.9

OH Heptyl alcohol (1-heptanol) 0.9

OH Decyl alcohol (1-decanol) <0.04

Source: Kelter, Carr, and Scott. Chemistry: A World of Choices; McGraw-Hill: New York, 2003.

TABLE 11.3

Methanol

Ethanol

n-Propanol

Visualization: Saturated

Solution Equilibrium

Video Lesson: Types of Solutions