3.3 Phase Changes

Reading StrategySummarizing Copy the diagram. As youread, complete the description of energy flowduring phase changes.

Key ConceptsWhat are six commonphase changes?

What happens to asubstance’s temperatureand a system’s energyduring a phase change?

How does thearrangement of watermolecules change duringmelting and freezing?

How are evaporation andboiling different?

Vocabulary◆ phase change◆ endothermic◆ heat of fusion◆ exothermic◆ vaporization◆ heat of vaporization◆ evaporation◆ vapor pressure◆ condensation◆ sublimation◆ deposition

Massive chunks of frozen water called icebergs are a common sightoff the continent of Antarctica. A large iceberg like the one in Figure 15contains enough fresh water to supply millions of people with waterfor a year. During the summer in southern Australia, fresh water is ascarce resource. People have proposed towing icebergs to Australiafrom Antarctica. The plan has not been implemented because the tripcould take months to complete and much of the iceberg would meltalong the way. In this section, you will find out what happens when asubstance, such as water, changes from one state to another.

Characteristics of Phase ChangesWhen at least two states of the same substance are present, scientistsdescribe each different state as a phase. For example, if an iceberg is float-ing in the ocean, there are two phases of water present—a solid phaseand a liquid phase. A phase change is the reversible physical change thatoccurs when a substance changes from one state of matter to another.

Endothermic Exothermic

LiquidSolid a. ? Solid

b. ? c. ? d. ?

Solid e. ? f. ?

Gas

Solid

84 Chapter 3

Figure 15 The solid and liquidphases of water are visible in thisphotograph of an iceberg in theAmundsen Sea near Antarctica.

84 Chapter 3

FOCUS

Objectives3.3.1 Describe phase changes.3.3.2 Explain how temperature can

be used to recognize a phasechange.

3.3.3 Explain what happens to themotion, arrangement, andaverage kinetic energy of watermolecules during phasechanges.

3.3.4 Describe each of the six phasechanges.

3.3.5 Identify phase changes asendothermic or exothermic.

Build VocabularyWord-Part Analysis List on the boardthe following word parts and meanings:-ion, “the act of” or “the result of anaction”; -ic, “related to or characterizedby”; endo-, “inside”; exo-, “outside”;therm, “heat”; -ize, “to become.” Havestudents identify these word parts in thevocabulary terms. Discuss the terms’meanings with students.

Reading Strategya. Liquid b. Liquid c. Gas d. Liquide. Gas f. Gas

INSTRUCT

Characteristics ofPhase ChangesUse VisualsFigure 15 Note that water is one ofthe few substances that exist naturallyas a solid, liquid, and gas under ordinaryconditions. Have students look at thefigure and read the caption. Ask, Whichtwo phases of water are visible in thephotograph? (Solid and liquid) Wherewould the third phase most likely befound? (In the air)Visual, Logical

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Reading Focus

1

Section 3.3

Print• Laboratory Manual, Investigation 3A• Reading and Study Workbook With

Math Support, Section 3.3Transparencies, Section 3.3

Technology• Probeware Lab Manual, Lab 1• Interactive Textbook, Section 3.3• Presentation Pro CD-ROM, Section 3.3• Go Online, NSTA SciLinks, Phases of matter;

PHSchool.com, Data sharing

Section Resources

States of Matter 85

In Figure 16, a state of matter is listed at each corner of the triangle.Each arrow in the diagram represents a different phase change. Eachpair of arrows represents a set of reversible changes. For example, thearrow starting at the solid phase and ending at the liquid phase repre-sents melting. The arrow starting at the liquid phase and ending at thesolid phase represents freezing.

Melting, freezing, vaporization, condensation, sublimation,and deposition are six common phase changes. All phase changesshare certain characteristics related to energy and temperature.

Temperature and Phase Changes One way torecognize a phase change is by measuring the temperature ofa substance as it is heated or cooled. The temperature ofa substance does not change during a phase change.

Naphthalene (NAF thuh leen) is a compound that is sometimes usedin mothballs. Figure 17 is a graph of the data collected when a solidpiece of naphthalene is slowly heated. Temperature readings are takenat regular intervals. At first the temperature rises as the solid naphtha-lene warms up. But at 80°C, the temperature of the naphthalene stopsrising. The temperature remains at 80°C, which is the melting point ofnaphthalene, until melting is complete.

If liquid naphthalene is placed in an ice-water bath, the temperatureof the liquid will drop until it reaches 80°C. It will remain at 80°C untilall the liquid freezes. The temperature at which a substance freezes—itsfreezing point—is identical to the temperature at which it melts. Thefreezing and melting points of naphthalene are both 80°C.

If naphthalene is heated after it has completely melted, its tempera-ture begins to rise again. The temperature keeps rising until it reaches218°C, which is the boiling point of naphthalene. Until boiling is com-plete, the temperature remains at 218°C.

Time (minutes)

Tem

per

atu

re (

�C)

100

90

80

70

60

50

40

30

200 1 2 3 4 5 6 7 8

Heating Curve for Naphthalene

Figure 17 This graph shows whathappens to the temperature of asolid sample of naphthalene asthe sample is slowly heated. Using Graphs What happenedto the temperature in the intervalbetween four and seven minutes?

Figure 16 This diagram lists sixphysical changes that can occuramong the solid, liquid, andgaseous phases of a substance.Interpreting Diagrams Explainwhy the changes are grouped intothe pairs shown on the diagram.

Melting

Freezing

Condensation

VaporizationSubl

imat

ion

Dep

ositi

on

Solid Liquid

Gas Build Reading LiteracyRelate Text and Visuals Refer topage 190D in Chapter 7, whichprovides the guidelines for relating text and visuals.

As students read the section, have themrefer to Figure 17. Ask, In what state isthe naphthalene at 20°C? In what stateis the naphthalene at 90°C? (Solid, liquid)What are the melting and freezingpoints of naphthalene? (Both are 80°C.)What would the heating curve fornaphthalene look like if the graph were extended beyond a temperatureof 100°C? (The heating curve would riseuntil it reached the boiling point ofnaphthalene, which is 218°C.)Visual, Logical

FYIThe plateau on the graph in Figure 17 isnot completely flat as it should be whilea substance melts. This graph representsactual data collected in a lab. It is likelythat there were some impurities in thenaphthalene sample, which caused someminor fluctuation in the melting point.

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States of Matter 85

Customize for English Language Learners

Build a Science GlossaryDirect students to Figure 16. Before they read,have them use the figure to make a list of thesix terms in the arrows for a glossary. Students

should add definitions for the terms to theglossary as they read the section. Encouragestudents to include examples of the six phasechanges along with the definitions.

Answer to . . .

Figure 16 The groupings representpairs of reversible changes.

Figure 17 The temperature remainedfairly constant.

Energy absorbed

86 Chapter 3

Energy and Phase Changes During a phase change, energyis transferred between a substance and its surroundings. The directionof the transfer depends on the type of phase change. Energy iseither absorbed or released during a phase change.

The ice sculpture in Figure 18 isn’t going to last forever. When thetemperature of the air rises above 0°C or when sunlight shines directlyon the ice, an ice sculpture begins to melt. Melting is an example of anendothermic change. During an endothermic change, the systemabsorbs energy from its surroundings.

The amount of energy absorbed depends on the substance. Forexample, one gram of ice absorbs 334 joules (J) of energy as it melts.This amount of energy is the heat of fusion for water. Fusion is anotherterm for melting. The heat of fusion varies from substance to substance.

One gram of water releases 334 joules of energy to its surroundingsas it freezes, the same amount of energy that is absorbed when one gramof ice melts. Farmers use this release of energy to protect their crops.When farmers expect temperatures to drop slightly below 0°C, theyspray the crops with water as shown in Figure 19. As water freezes, itreleases heat. The flow of heat slows the drop in temperature and helpsprotect the crops from damage. Freezing is an example of an exother-mic change. During an exothermic change, the system releases energyto its surroundings.

An understanding of phase changes can be useful in many situa-tions. The How It Works box explains how the design of an ice rink inUtah allows the manager of the rink to control the hardness of the ice.

How much energy does one gram of ice absorb asit melts?

Figure 18 This ice sculpture of adog sled was carved at a winterfair in Fairbanks, Alaska. The icesculpture will start to melt if thetemperature rises above 0ºC orsunlight shines directly on the ice.

Figure 19 Energy released as iceforms on these strawberry plantskeeps the plants from freezing attemperatures slightly below 0°C.Applying Concepts Explainwhy a farmer would need tokeep spraying the plants withwater while the temperatureremains below freezing.

Energy released

86 Chapter 3

Integrate BiologyOne of the main causes of frost damagein crops is ice-nucleating bacteria. Thesebacteria provide a nucleus around whichice can form. When ice forms on plants, itpierces their cell walls and causes the cellsto desiccate, or dry out because waterescapes from the cells. Have studentssearch the Internet for information onice-nucleating bacteria. Have studentspresent what they learn in a pamphletthat offers suggestions on how to preventfrost damage from bacteria. Verbal

Energy Transfer

Purpose Students observe that thephase change from ice to water isendothermic.

Materials tray of ice cubes

Procedure Place a tray of ice cubes on a counter at the beginning of class. As theice begins to melt, ask students where theenergy to melt the ice comes from.

Expected Outcome Students willobserve that the ice melts. They will inferthat the energy to melt the ice comesfrom the air and the counter becausethese materials are warmer than the ice.Visual

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Section 3.3 (continued)

Preventing Frost Damage For a frostcontrol system to work, a farmer needs to startthe sprinklers just before the temperature atground level reaches 0°C. Placing a shallowpan of water in the lowest part of the field is aneasy way to monitor the temperature. As soonas ice begins to form around the edge of thepan, the sprinklers should be turned on. Thesprinklers should stay on until the temperaturerises enough to start melting the ice.

This method of frost control is effective for strawberry plants down to about �6.6°C.When strawberries are in bloom, sprinklers are best started at 1.1°C. Blueberries withopen blooms are safe from damage at tem-peratures as low as 0°C. Because of the cost of installation and operation, frost controlsystems are generally used for crops that candemand high unit prices, such as oranges,strawberries, Blueberries, and asparagus.

Facts and Figures

States of Matter 87

Custom-Tailored Ice At the Utah Olympic Oval the hardness of the ice can becontrolled. The ice must be cold and hard for long speed-skating races, where the length of a skater’s glide isimportant. For shorter races, the skaters need moretraction, so the ice is made a little warmer and softer.Interpreting Diagrams What is the purpose of thesand layer?

Speed SkaterThis skater is racing at the UtahOlympic Oval, one of theworld’s most technicallyadvanced ice rinks.

Sand layer To preventthe gravel layer from

freezing, which coulddamage the rink structure,the sand layer is kept at atemperature above 0°C.

Gravel layer This layerprovides the foundation

for the rink.

Chilled concrete slab Thetemperature is controlled by cold

salt water, which is pumped throughpipes embedded in the slab. The saltwater remains liquid because its freezingpoint is lower than that of pure water.

Paint layer withmarkings

Paint layer with logos

Paint layer withbackground color

Ice layersThe skating surface is formed with asmany as 24 layers of ice and paint. Warm water is used for the top layers because it contains less dissolved air, and canproduce a denser, harder, frozen surface.

Skating surface The surface is built up fromthin layers of ice to a depth of less than 2 cm.

The ice can be as cold as –8°C. The temperature andthe hardness of the ice are controlled by tiny changesto the temperature of the underlying concrete layer.

Lubricant

Insulating layers These layersprevent heat from rising up

into the concrete slab from thewarmer sand layer underneath.

Custom-Tailored IceSkaters traveling at speeds up to80 km/h will be slowed by a surface thatis not smooth, level, and uniformly hard.The conditions of the ice at the UtahOlympic Oval, which was constructedfor the 2000 Winter Olympics, can bechanged overnight. For sprinters, the iceis softer so that their traction, or abilityto grip the ice, is greater. A hardersurface allows long-distance skaters to glide more easily across the ice.

Controlling the indoor climate is also important in maintaining the icesurface. The Utah Olympic Oval has twoair-filtering systems and a dehumidifierthat can reduce humidity to 3%. The airtemperature varies only 3°C from the iceto the ceiling. Controlling the conditionsabove and below the ice allows forprecise control of the ice surface.

Interpreting Diagrams The sandlayer prevents the gravel layer fromfreezing and damaging the rinkstructure.Visual, Logical

For EnrichmentHave students research playing surfacesused in other sports, such as football,tennis, and golf. Explain that differentsurfaces affect an athlete’s performance.Students could compare the advantagesand disadvantages of artificial turf andgrass football fields; or cement, clay, andgrass tennis courts. They could also findout how the length of grass affects shotsat different locations on a golf course.Logical

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States of Matter 87

Answer to . . .

Figure 19 Until the temperature risesabove freezing, the farmer mustcontinue to provide energy to keep theplants from freezing.

334 joules

88 Chapter 3

Melting and FreezingIn water, hydrogen and oxygen atoms are combined in small unitscalled molecules. Each water molecule contains two hydrogen atomsand one oxygen atom. The arrangement of molecules in waterbecomes less orderly as water melts and more orderly as water freezes.

Melting In ice, attractions between water molecules keep the mol-ecules in fixed positions. When ice cubes are removed from a freezerand placed in an empty glass, heat flows from the air to the ice. As theice gains energy, the molecules vibrate more quickly. At the meltingpoint of water, 0°C, some molecules gain enough energy to overcomethe attractions and move from their fixed positions. When all the mol-ecules have enough energy to move, melting is complete. Any energygained by the water after the phase change increases the average kineticenergy of the molecules, and the temperature rises.

Freezing When liquid water is placed in a freezer, energy flowsfrom the water to the air in the freezer, and the water cools down. Asthe average kinetic energy of its molecules decreases, they move moreslowly. At the freezing point of water, some molecules move slowlyenough for the attractions between molecules to have an effect. Whenall the molecules have been drawn into an orderly arrangement, freez-ing is complete. Any energy removed from the ice after the phasechange decreases the average kinetic energy of the molecules, and thetemperature of the ice drops.

Often, people think of cold temperatures when theyhear the term freezing. But substances that are solids atroom temperature can freeze at temperatures that are quitehigh. For example, silicon freezes at 1412°C (2574°F). As acomparison, you can bake cookies at 177°C (350°F).

Vaporization and CondensationFigure 20 shows how food cools and stays cold in a refrig-erator. The process depends on a substance that changesfrom a liquid to a gas to a liquid over and over again.During these phase changes, energy flows from the insideof the refrigerator to the outside.

The phase change in which a substance changes froma liquid into a gas is vaporization. Vaporization is anendothermic process. That is, a substance must absorbenergy in order to change from a liquid to a gas. Onegram of water gains 2261 joules of energy when it vapor-izes. This amount of energy is the heat of vaporizationfor water. The heat of vaporization varies from substanceto substance.

Figure 20 In a refrigerator, a pairof phase changes keep the foodcold. Energy from inside the foodcompartment is used to change aliquid to a gas in the evaporator.This energy is released when thecompressed gas changes back to aliquid in the condenser.

Energy removedfrom food

compartment

Energyreleased to

surroundings

Compressor

Condenser

Evaporator

For: Links on phases of matter

Visit: www.SciLinks.org

Web Code: ccn-1033

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Melting and FreezingUse VisualsFigure 20 Keeping food cool requiresthe transfer of energy. Ask, What changeoccurs in the evaporator? Is thischange exothermic or endothermic?(Liquid to gas; the change is endothermicbecause energy is absorbed from the foodcompartment) What change occurs in the condenser? Is this changeexothermic or endothermic? (Gas isliquid; the change is exothermic becauseenergy is released to the surroundings)Visual

FYIRefrigeration is also discussed in Chapter 16, where the focus is on theneed for work to be done for heat toflow from an area of lower temperatureto an area at higher temperature. In thischapter, the focus is on the endothermicand exothermic phase changes.

Vaporization andCondensationUse CommunityResourcesInvite an appliance repairperson to cometo the class and speak about compoundsused in air conditioner and refrigeratorcooling coils. Suggest that the speakerdescribe the types of compounds used,how the compounds have changed overtime, and whether there are any problemswith the replacement compounds. Interpersonal

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Chlorofluorocarbons Under the Clean AirAct, which became effective in November of1995, it became illegal to release into the aircompounds that harm Earth’s ozone layer.Chlorofluorocarbons (CFCs) are among thebanned substances. CFCs were used as refrig-erants in cooling devices, such as refrigeratorsand air conditioners.

Because CFCs do not dissolve easily in waterand are generally not very reactive, they tendto stay in the atmosphere once they are

released. In the lower atmosphere, theirpresence causes few problems. However, thesecompounds move into the stratosphere wherethey are exposed to high energy radiation. Thisradiation causes the CFCs to react and releasechlorine atoms. These chlorine atoms reactwith oxygen to form chlorine monoxide, whichcauses the ozone to decompose into oxygen.Hydrogen fluorocarbons (HFCs) have replacedCFCs in many cooling devices.

Facts and Figures

Download a worksheet on phasesof matter for students to complete,and find additional teacher supportfrom NSTA SciLinks.

Scientists distinguish two vaporization processes—boiling and evaporation. Evaporation takes placeat the surface of a liquid and occurs at temperaturesbelow the boiling point.

Evaporation If you go outside after a rain showeron a sunny, warm day, you may notice puddles ofwater. If you return to the same location after a fewhours, the puddles may be gone. This disappearanceof the puddles is due to evaporation. Evaporation isthe process that changes a substance from a liquid to agas at temperatures below the substance’s boiling point.

Figure 21 shows what is happening as water evaporates from asmall, shallow container. Some molecules near the surface are movingfast enough to escape the liquid and become water vapor. (A vapor isthe gaseous phase of a substance that is normally a solid or liquid atroom temperature.) The greater the surface area of the container, thefaster the water evaporates.

What happens if the water is in a closed container? As the waterevaporates, water vapor collects above the liquid. The pressure causedby the collisions of this vapor and the walls of the container is calledvapor pressure. The vapor pressure of water increases as the tempera-ture increases. At higher temperatures, more water molecules haveenough kinetic energy to overcome the attractions of other moleculesin the liquid.

Boiling As you heat a pot of water, both the temperature and thevapor pressure of the water increase. When the vapor pressure becomesequal to atmospheric pressure, the water boils. The temperature atwhich this happens is the boiling point of water.

The kinetic theory explains what happens when water boils. As thetemperature increases, water molecules move faster and faster. Whenthe temperature reaches 100°C, some molecules below the surface ofthe liquid have enough kinetic energy to overcome the attraction ofneighboring molecules. Figure 22 shows that bubbles of water vaporform within the liquid. Because water vapor is less dense than liquidwater, the bubbles quickly rise to the surface. When they reach thesurface, the bubbles burst and release water vapor into the air.

How does the surface area of a liquid affect therate of evaporation?

89

Figure 22 Boiling takes place throughouta liquid. Applying Concepts Explain whythe temperature of water does not riseduring boiling.

Figure 21 Evaporation takesplace at the surface of a liquid.

Build Science SkillsDesigning Experiments The rate ofevaporation is affected by surface area.Ask students to design an experimentusing water in a beaker with gradationsto test this statement. Have studentsidentify the variables that will need to be controlled (type of liquid, volume of liquid, temperature of liquid, time).Ask, What will the manipulatedvariable be? (Surface area) What will the responding variable be? (Rate ofevaporation) How will you measure therate of evaporation? (One acceptableanswer is measuring the liquid level after a specified interval of time.) Allow studentsto conduct approved experiments andsummarize their results in a report. Verbal, Logical

Students often think that atoms andmolecules expand as the temperaturerises. Use water’s expansion when itfreezes to challenge this misconception.Explain that the space between watermolecules is greater in ice than in liquidwater because of the arrangement ofmolecules in ice. When ice melts, thespace between molecules decreases andthe volume decreases. Logical

FYIEventually, no more vapor can collectabove the liquid in a closed container.The water continues to evaporate, but some of the vapor condenses inreturn. This is an example of a dynamicequilibrium. Chemical (and physical)equilibrium is introduced in Chapter 7.

Water, especially tap water, containsdissolved gases such as oxygen andnitrogen. When water is heated, thesolubility of these gases in waterdecreases, which causes small bub-bles to form in heated water before it begins to boil.

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Answer to . . .

Figure 22 Energy absorbed by the system during boiling is used toovercome attractions among watermolecules.

The greater the surfacearea, the faster the

water evaporates.

Observing Phase Changes

Materials250-mL Erlenmeyer flask, graduated cylinder,thermometer, dry ice

Procedure1. Pour 150 milliliters of water into a 250-mL

Erlenmeyer flask. Place a thermometer in theflask. CAUTION Wipe up any spilled water rightaway to avoid slips and falls.

2. Observe what happens after your teacher addsa small piece of dry ice to the flask. (Dry ice issolid carbon dioxide.) CAUTION Dry ice candamage skin on contact. Do not touch the dry ice.

3. Record the temperature of the water just afterthe dry ice is added and again after it is nolonger visible.

Analyze and Conclude1. Observing What happened when the dry ice

was added to the water?

2. Analyzing Data Did adding the dry icecause the water to boil? Explain your answer.

3. Inferring What was the source of thebubbles in the water?

4. Formulating Hypotheses What caused acloud to form above the flask?

5. Applying Concepts What phase changesoccurred in the flask?

Figure 23 Water vapor from theair condensed into drops of liquidwater on these blades of grass.

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The boiling point of a substance depends on the atmospheric pres-sure. The normal boiling point of water at sea level is 100°C. At higherelevations, the atmospheric pressure is lower. Do you know thatDenver, Colorado, is called the mile-high city? This nickname is basedon Denver’s location at one mile above sea level. In Denver, the vaporpressure of water will equal atmospheric pressure at temperaturesbelow 100°C. The boiling point of water in Denver can be as low as95°C. Food does not cook as quickly at 95°C as it does at 100°C. Pastatakes longer to cook in Denver than in New Orleans, Louisiana, a citythat is located near sea level.

Condensation Have you ever come out of a shower to find yourbathroom mirror clouded over? The “cloud” on the mirror is caused bywater vapor that cooled as it came in contact with the mirror. Thewater vapor transferred heat to the mirror and condensed into liquidwater. Condensation is the phase change in which a substance changesfrom a gas or vapor to a liquid. This process is also responsible for themorning dew on the blades of grass in Figure 23. Condensation is anexothermic process.

90 Chapter 3

Observing PhaseChanges

ObjectiveAfter completing this lab, students willbe able to • identify examples of condensation and

sublimation.

Skills Focus Observing

Prep Time 15 minutes

Advance Prep Obtain dry ice from a local supplier or from a scientificsupply house, ice cream wholesaler, or compressed gas dealer. Do not storedry ice in an airtight container. Leave awindow open in your car if you have dryice in the car. Use a hammer to carefullybreak the dry ice into pea-sized pieces.

Class Time 20 minutes

Safety Do not handle dry ice with barehands. It damages skin tissue on contact.Wear safety goggles, a lab apron, andleather gloves when handling dry ice. Use forceps to dispense a pea-sized pieceof dry ice into each flask. Caution studentsnot to touch dry ice and to use care whenhandling glassware to avoid breakage.Provide only nonmercury thermometers.

Teaching Tips• Do this lab after students have studied

sublimation.• If students are having trouble with

Question 2, ask at what temperaturewater boils. If they are having troublewith Question 4, remind them thatwater vapor exists above the liquid.

Expected Outcome When dry ice is added to water, vigorous bubblingbegins and a fog appears above the liquid.

For EnrichmentChallenge students to propose a way tohave solid, liquid, and gaseous watertogether in the same test tube. One wayis to place ice in liquid water in a closedcontainer. Water vapor will collect abovethe surface of the water. Visual, Logical

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Section 3.3 (continued)

Analyze and Conclude1. Bubbles formed in the water, and a fogformed above the water. The temperature of thewater decreased.2. No, because the water was not at 100°C whenbubbles appeared.3. Carbon dioxide gas from the dry ice

4. Cold carbon dioxide gas chilled the air abovethe water, causing water vapor to condense intotiny liquid droplets.5. Sublimation of carbon dioxide and condensa-tion of water occurred. Students may also mentionvaporization of water. Logical, Kinesthetic

Section 3.3 Assessment

Reviewing Concepts1. Name six common phase changes.

2. What happens to the temperature of asubstance during a phase change?

3. How does the energy of a system changeduring a phase change?

4. What happens to the arrangement ofwater molecules as water melts and freezes?

5. What is the difference betweenevaporation and boiling?

6. Explain why sublimation and deposition areclassified as physical changes.

Critical Thinking7. Applying Concepts How can the mass of a

pile of snow decrease on a sunny day whenthe air temperature does not rise above 0°C?

8. Drawing Conclusions At room temperature,table salt is a solid and acetone is a liquid.Acetone is the main ingredient in nail polishremover. What conclusion can you draw aboutthe melting points of these materials?

Sublimation and DepositionDirectors of concerts and plays sometimes use dry ice to create afog-like special effect. Dry ice is the common name for the solidform of carbon dioxide. At room temperature, dry ice can directlychange from a solid to a colorless gas. Sublimation is the phasechange in which a substance changes from a solid to a gas or vaporwithout changing to a liquid first. Sublimation is an endothermicchange. As dry ice sublimes, the cold carbon dioxide vapor causeswater vapor in the air to condense and form clouds.

Where does the name dry ice come from? Solid carbon dioxidedoes not form a liquid as its temperature rises. Suppose 100 steaksare shipped from Omaha, Nebraska, to a supermarket in Miami,Florida. The steaks will spoil unless they are kept cold during thetrip. If regular ice is used, water collects in the shipping containeras the ice melts. If the steaks are shipped in dry ice, the containerand the steaks stay dry during the journey. Figure 24 shows anotheruse of dry ice.

When a gas or vapor changes directly into a solid without firstchanging to a liquid, the phase change is called deposition. Thisexothermic phase change is the reverse of sublimation. Depositioncauses frost to form on windows. When water vapor in the air comesin contact with cold window glass, the water vapor loses enoughkinetic energy to change directly from a gas to a solid.

States of Matter 91

Steps in a Process Write a paragraphdescribing three steps that must occur for awater molecule to start on the surface of hotbath water and end up on the surface of abathroom mirror. Note whether the phasechanges that take place during the processare endothermic or exothermic. (Hint: Usewords such as first, next, and finally to showthe order of events.)

Figure 24 A technician at TinkerAir Force Base in Oklahoma hangsa mosquito trap. The trap is baitedwith dry ice because mosquitoesare attracted to carbon dioxide.

Sublimation andDepositionUse VisualsFigure 24 Have students examineFigure 24. Explain that the trap washung as part of a study of West Nile virus,a mosquito-borne disease. Mosquitoesacquire the virus from infected birds andthen transmit the virus to other animalsand humans. Mosquitoes are collectedand sent to a laboratory for identificationby species and gender. Samples of similarfemale mosquitoes are blended and theirRNA is analyzed for West Nile virus. Ask,Why is dry ice used to bait the mos-quito trap rather than gaseous carbondioxide? (The large amount of carbondioxide stored in dry ice is released slowlyover time rather than all at once.)Logical

ASSESSMENTEvaluateUnderstandingHave students write the six terms thatdescribe phase changes. Next to eachterm, students should describe the typeof change. For example, for melting,students would indicate that the phasechange is from a solid to a liquid. Tellstudents to include arrows pointing upto identify exothermic changes andarrows pointing down to identifyexothermic changes.

ReteachUse Figure 16 to review the six phasechanges described in this section. Ask,What do the phase changes with redarrows have in common? (They are allexothermic processes.) What do thephase changes with blue arrows havein common? (They are all endothermicprocesses.)

First, the water molecule on the surfaceof the bath water evaporates—anendothermic change. Next, randommotion of the molecule carries it tothe surface of the mirror. Finally, themolecule condenses on the mirror—an exothermic change.

If your class subscribes tothe Interactive Textbook, use it toreview key concepts in Section 3.3.

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States of Matter 91

5. Evaporation takes place at the surface of aliquid and occurs at temperatures below theboiling point.6. A substance’s identity does not changeduring sublimation or deposition.7. When snow absorbs energy from sunlight,it either melts or sublimes.8. Table salt must have a melting point aboveroom temperature, and acetone must have amelting point below room temperature.

Section 3.3 Assessment

1. Melting, freezing, vaporization,condensation, sublimation, and deposition2. The temperature of a substance does notchange during a phase change.3. Energy is either released or absorbed duringa phase change.4. The arrangement of molecules becomes lessorderly as water melts and more orderly aswater freezes.