Chapter 4: Formation of Compounds - Glencoe...spoon, the crystals shatter. This shattering shows...

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4CHAPTER

118

Formation of CompoundsFormation ofCompounds4

CHAPTER

4.1 The Variety of CompoundsMiniLab 4.1 Evidence of a Chemical

Reaction: Iron VersusRust

4.2 How Elements Form CompoundsMiniLab 4.2 The Formation of Ionic

CompoundsChemLab The Formation and

Decomposition of Zinc Iodide

Chapter PreviewSections Compounds—Submicroscopic

to Macroscopic

Every compound has its own unique set ofphysical and chemical properties. Eachmacroscopic property you observe has a

submicroscopic explanation. The beautifulmosaic of color in this volcanic mineral, knownas eucrite, is the result of submicroscopic combi-nations of matter.

Compounds—Submicroscopicto Macroscopic

119

Wine to Water When substances go through a chemical reaction to formnew compounds, the new substances often have differentphysical properties from the original reactants. How can achemical reaction change a red liquid to a clear liquid?

Materials

• 500-mL beaker • red food coloring• 25-mL graduated cylinder • laundry bleach• water • stirring rod

Safety Precautions

• Perform this experiment in a well-ventilated area orunder a hood.

• Bleach can damage skin or clothing.• Immediately notify your teacher of any spills.

Procedure

1. Pour 300 mL of water into a 500-mL beaker.2. Add three drops of red food coloring and stir the

mixture until the water turns uniformly red.3. Measure 15 mL of bleach and pour the bleach into

the beaker.4. Stir the mixture and observe any changes in the

solution.

Analysis

Research the compound in bleach that could account forthe color change of the mixture. What new compoundsmight be formed during this reaction to account for thecolor change?

Before scanning the chapter, look atthe definition of the word “com-pound.” Then, scan the chapter,noticing the different illustrations ofcompounds. Pick one and try to drawit in a different way, using othershapes, spacing, or perspective. Be imaginative with your drawing.

Reading Chemistry

Review the following conceptsbefore studying this chapter.Chapter 1: how elements and com-pounds differChapter 2: the structure of atoms; thearrangement of electrons in an atom

What I Already Know

Preview this chapter’s content andactivities at chemistryca.com

Start-up ActivitiesStart-up Activities

http://chemistryca.com

4.1

Chemistry is often like detec-tive work. To figure outthe submicroscopic

structure of a com-pound, you first haveto examine some ofits macroscopicproperties. Just asno two humanfingerprints are thesame, no two sub-stances have exactlythe same combina-tion of physical andchemical properties. The structure of a substance at the atomic level deter-mines its macroscopic properties. Therefore, looking at properties of com-pounds provides important clues about their submicroscopic structure andhow they form from atoms. That is powerful detective work. You can beginthis detective work by examining the properties of three familiar compounds:table salt, carbon dioxide, and water.

Salt: A Familiar CompoundWhat is the most popular food additive? In most kitchens, the answer is

salt. It is used in cooking and at the table to enhance the flavor of food.Chemists refer to it as sodium chloride. The chemical name tells you whatelements make up the compound; sodium chloride contains the elementssodium and chlorine.

Even though it’s possible to make sodium chloride from its elements,salt is so abundant on Earth that it is used to manufacture the elementssodium and chlorine. Sodium chloride occurs naturally in large, solid,underground deposits throughout the world and is dissolved in theworld’s oceans. Salt can be obtained by mining these solid deposits,Figure 4.1, and by the evaporation of seawater.

Besides enhancing food’s flavor, sodium chloride is an essential nutrientthat plays crucial roles in living things. If you live in an area that getssnow and ice in the winter, salt is sometimes used to melt ice on roads, asshown in Figure 4.1.

SECTION

Objectives✓ Distinguish theproperties of com-pounds from those of the elements ofwhich they are composed.

✓ Compare and con-trast the properties of sodium chloride,water, and carbondioxide.

✓ Analyze evidenceto conclude that dif-ferences exist in theways compoundsform.

Review VocabularyTransition element:any of the elements inGroups 3 through 12of the periodic table,all of which are metals.

SECTION PREVIEW

120 Chapter 4 Formation of Compounds

The Variety of Compounds

4.1 The Variety of Compounds 121

Figure 4.2Conductivity of SaltIn order to light the bulb,electric current must flowbetween the two electrodes.As you can see, no currentflows through the dry saltcrystals (left). When the salt isdissolved in water, the solu-tion conducts electricity andthe bulb lights (right). Purewater alone does not conductelectricity.

In some regions, trucksspread salt crystals onicy roads to melt theice. The salt lowers themelting point of ice toabout 15°F. If the airtemperature is below15°F, the salt won’t do much good. �

Figure 4.1Getting and Using SaltUnderground salt mining accounts for about 90 percentof the world’s salt production. These salt deposits formedwhen ancient seas evaporated millions of years ago. �

� Salt is the most commonfood seasoning. Someestimates say that 60percent of the averageAmerican’s sodiumintake comes from salt-ing food at the table.

Physical Properties of SaltYou already know some of the physical properties of table salt. It is a white

solid at room temperature. If you look at table salt under a magnifier, you’llnotice that the grains of table salt are little crystals shaped like cubes. Thesecrystals are hard, but when you press down on them with the back of aspoon, the crystals shatter. This shattering shows that the crystals are brittle.If sodium chloride is heated to a temperature of about 800°C, it melts andforms liquid salt. Solid sodium chloride does not conduct electricity, butmelted sodium chloride does. Salt also dissolves easily in water. The resultingsolution is an excellent conductor of electricity, as shown in Figure 4.2.

Evidence of a Chemical Reaction: Iron Versus RustHere’s a chemical reaction whose result everyone has seen—the rust-

ing of iron. Iron metal combines with oxygen in the air to form rust,Fe2O3, iron(III) oxide. This reaction is an example of a familiar com-pound forming from its component elements. One property of ironthat everyone is familiar with is that it is attracted to a magnet. Attrac-tion to a magnet is a simple property to measure—all you need is amagnet.

Procedure1. Obtain a small

wad of fresh steelwool in onesmall paper cupand anothersmall wad ofrusty steel woolin another smallpaper cup.

2. Obtain a 3" � 5"index card and amagnet wrappedin a plastic bag.

3. Test the fresh steel wool with

the magnet. Record your observations.

4. Hold the rusty steel wool overthe white card, and lightly rub

the rusty steel wool betweenyour thumb and forefinger.Some fine rust powder shouldfall onto the white card.

5. Next, hold the card up and slow-ly move the magnet under thecard. Record your observations.

Analysis1. What effect did the magnet

have on the fresh steel wool?

2. What did you observe whenthe magnet was moved underthe card with the rust powder?

3. Was your pile of rust powder apure substance? How does yourexperimental evidence supportyour answer?

4. What evidence do you havethat the rust is a different sub-stance from iron itself?

1


Chemical Properties of SaltThe chemical properties of sodium chloride are not so useful in our

detective work to determine its submicroscopic structure. Salt does notreact readily with other substances. It could sit in a salt shaker for hun-dreds or even thousands of years and still remain salt. It does not have tobe handled in any special way or be stored in a special container. Com-pounds with these chemical properties are referred to as stable or unreac-tive. You can get more clues about salt’s submicroscopic structure byanswering the question: How do the properties of salt compare with theproperties of its component elements, sodium and chlorine?

Properties of SodiumSodium is a shiny, silvery-white, soft, solid element as you can see in

Figure 4.3. From its location on the left side of the periodic table, youknow that it is a metallic element. Sodium melts to form a liquid when it


Figure 4.3A Comparison of Sodium andChlorineSodium is a metal, but it is softenough to be cut with a knife.Where it has been freshly cut,you can see that sodium has asilvery luster that is typical ofmany metals. �

Chlorine, in the flask on the right,is a greenish, poisonous gas atroom temperature. If you’ve everused liquid chlorine bleach, you’veprobably smelled chlorine. �

is heated above 98°C. Sodium must be stored under oil because it reactswith oxygen and water vapor in the air. In fact, it is one of the most reac-tive of the common elements. When a piece of sodium is dropped intowater, it reacts so violently that it catches fire and sometimes explodes.Because of its high reactivity, the free element sodium is never found inthe environment. Instead, sodium is always found combined with otherelements.

Properties of ChlorineThe element chlorine, shown in Figure 4.3, is a pale green, poisonous

gas with a choking odor. Because chlorine kills living cells and is slightlysoluble in water, it is an excellent disinfectant for water supplies andswimming pools. You can tell that chlorine is a nonmetal by its positionin the upper-right portion of the periodic table. Chlorine gas must becooled to �34°C before it turns to a liquid. Like sodium, it is among themost reactive of the elements and must be handled with extreme care.Chlorine is needed for many industrial processes, such as the manufactureof bleaches and plastics. Because of its industrial importance, large quan-tities of chlorine must be transported in railroad tank cars, tanker trucks,and river barges. If a train that is carrying chlorine derails, entire commu-nities are evacuated until the danger of a chlorine leak passes.

Look again at the photo on the opening page of this chapter. Whensodium and chlorine react to form sodium chloride, two dangerous ele-ments combine to form a stable, safe substance that we consume everyday. What could be happening in such a change? You will find the detailsin Section 4.2, but first look at two other common compounds whoseproperties are different from those of sodium chloride.


Carbon Dioxide: A Gas to ExhaleCarbon dioxide is a colorless gas. Take a deep breath and hold it for a few

seconds. What you have inhaled is air, a colorless mixture of nitrogen andoxygen gases with small amounts of argon, water vapor, and carbon diox-ide. Now, exhale. The mixture of gases that you exhale contains more than100 times the amount of carbon dioxide that was in the air that youinhaled, as you can see in Table 4.1. In contrast, the quantity of oxygen isreduced by five percent. While the air was in your lungs, chemical and phys-ical processes exchanged some of the oxygen for carbon dioxide. Carbondioxide is an important chemical link between the plant and animal world.Green plants and other plantlike organisms take in carbon dioxide and giveoff oxygen during photosynthesis. Both plants and animals, includinghumans, use oxygen and give off carbon dioxide during respiration.

Substance Percent in Inhaled Air Percent in Exhaled AirNitrogen 78 75Oxygen 21 16Argon 0.9 0.9Carbon dioxide 0.03 4Water vapor variable (0 to 4) increased

Table 4.1 Composition of Inhaled and Exhaled Air

Physical Properties of Carbon DioxideCarbon dioxide, like sodium chloride, is also a compound, but its prop-

erties differ from those of sodium chloride. For example, salt is a solid atroom temperature, but carbon dioxide is a colorless, odorless, and taste-less gas. When carbon dioxide is cooled below �80°C, the gas changesdirectly to white, solid carbon dioxide without first becoming a liquid.Because the solid form of carbon dioxide does not melt to a liquid, it iscalled dry ice, as shown in Figure 4.4. Carbon dioxide is soluble in water,as anyone who has ever opened a carbonated beverage knows. A watersolution of carbon dioxide is a weak conductor of electricity. You canmake carbon dioxide from its elements by burning carbon in air. Coal and charcoal are mostly carbon.

Chemical Properties of Carbon DioxideLike sodium chloride, carbon dioxide is relatively stable. Carbon diox-

ide is used in some types of fire extinguishers because it does not supportburning, Figure 4.4. Photosynthesis is probably the most significantchemical reaction of carbon dioxide. In photosynthesis, plants use energyfrom the sun to combine carbon dioxide and water chemically to makesimple sugars. Plants use these sugars as raw materials to make many

Some people think thatwhen we inhale, thelungs absorb all theoxygen from theinhaled air and replaceit completely with CO2.You can see from Table4.1 that this notionisn’t true.


other kinds of compounds, from cellulose in wood and cotton to oils suchas corn oil and olive oil. Photosynthesis is only one part of a natural cycleof chemical reactions known as the carbon cycle.

The Properties of CarbonAs with sodium chloride, the properties of carbon dioxide differ from

the properties of its elements. Carbon is a nonmetal and is fairly unreac-tive at room temperature. However, at higher temperatures, it reacts withmany other elements. Charcoal is approximately 90 percent carbon. Asanyone with a charcoal grill knows, carbon burns fairly easily and is anexcellent source of heat, as shown in Figure 4.5. Carbon forms a hugevariety of compounds. In fact, the majority of compounds that make upliving things contain carbon. Carbon compounds are so significant thatan entire branch of chemistry, called organic chemistry, is dedicated totheir study.

Figure 4.4Properties of CarbonDioxide

� Solid carbon dioxide iscalled dry ice. Under ordi-nary conditions, it doesnot melt into a liquid;instead, it changes directlyinto a gas. The dry ice inthis cylinder is immersed inwater and is producingbubbles of carbon dioxidegas. The white vapors yousee here are not carbondioxide, but condensedwater vapor carried alongwith the cold CO2 gas. Youcan tell that carbon diox-ide is denser than air.

� Carbon dioxide does not supportburning. In fact, it is often used toput out fires. Fire extinguisherslike this one are filled with com-pressed carbon dioxide. BecauseCO2 is denser than air, it displacesair and deprives the fire of a sup-ply of oxygen.

Figure 4.5A Chemical Property of CO2

When you burn charcoal to cook chicken,carbon and oxygen in the air combine toproduce carbon dioxide. When elementscombine to form a substance that is morestable, the reaction often gives off energyin the form of heat.


Water, Water EverywhereWater is the third familiar compound you will compare in your detec-

tive work on the submicroscopic structure of compounds. The formalchemical name of water is dihydrogen monoxide, but nobody calls it that.Water covers approximately 70 percent of Earth’s surface and also makesup about 70 percent of the mass of the average human body.

Physical Properties of WaterThe properties of water are different from those of sodium chloride

and carbon dioxide. Water is the only one of the three compounds thatoccurs in Earth’s environment in all three states of matter, as shown inFigure 4.7. At sea level, liquid water boils into gaseous water (steam) at100°C and freezes to solid water (ice) at 0°C. Pure water does not conductelectricity in any of its states. Water is also excellent at dissolving othersubstances. It is often called the universal solvent in recognition of thisvaluable property. Water plays a vital role in the transport of dissolvedmaterials, whether the aqueous solution is flowing down a river; up thexylem in a tree; or through the veins, capillaries, and arteries of your cir-culatory system.

The Properties of OxygenLike carbon, oxygen is another nonmetal. It is a colorless, odorless, and

tasteless gas that makes up about 21 percent of the air you breathe. Whenmaterials such as paper burn in air, they react with oxygen, which is whypeople commonly say that oxygen supports burning.

Oxygen gas becomes a liquid when it is cooled to �183°C, and it isslightly soluble in water. The gills of fishes absorb dissolved oxygen fromtheir water environment. Oxygen is more reactive than carbon and com-

bines with many other elements. A prime example of its reactivi-ty is the process of rusting, in which the element iron com-

bines with oxygen from air. Many of the compoundsthat make up Earth’s crust contain oxygen. Oxygen is

the most abundant element in Earth’s crust, asshown in Figure 4.6.

Oxygen46%

Aluminum8%

Magnesium 4%

Calcium 2.4%Potassium 2.3%

Sodium 2.1%Other < 1%

Iron6%

Silicon28%

Figure 4.6Elemental Composition of Earth’s CrustAs you can see from the graph, oxygen makes upabout 46 percent of Earth’s crust. Nearly all of thisoxygen occurs in compounds with other elements.Silicon is the next most abundant element inEarth’s crust. Therefore, you won’t be surprised tolearn that much of the oxygen is tied up in silicondioxide, commonly known as sand.


Chemical Properties of WaterWater is a stable compound; it doesn’t break down under normal con-

ditions and does not react with many other substances. Perhaps the mostinteresting property of water is its ability to act as a medium in whichchemical reactions occur. Nearly all of the chemical reactions in thehuman body and many important reactions on Earth occur in an aqueousenvironment. Without water, these reactions could not occur or wouldoccur extremely slowly. In addition, water and carbon dioxide are thestarting materials for photosynthesis, the process that makes life on Earthpossible. Now, compare the properties of water with those of its compo-nent elements, hydrogen and oxygen. The properties of oxygen weredescribed on page 126.

The Properties of HydrogenHydrogen is the lightest and most abundant element in

the universe. Hydrogen is usually classified as a nonmetal.Like oxygen, the element hydrogen is an odorless, tasteless,and colorless gas. Hydrogen is a reactive element. Becauseof its reactivity, it is seldom found as a free element onEarth. Instead, it occurs in a variety of compounds, partic-ularly water. Hydrogen reacts vigorously with many ele-ments, including oxygen, as shown in Figure 4.8. This reac-tion forms water. The temperature of hydrogen gas must belowered to a frigid �253°C before it turns to liquid. Hydro-gen does not conduct electricity and is only slightly solublein water.

Figure 4.7Three States of WaterThe iceberg shows water in itssolid state, ice. In the ocean,water is in the liquid state.The atmosphere containswater vapor. Clouds consist ofwater vapor that has con-densed back into smalldroplets of liquid water.

Figure 4.8The Reaction of Hydrogen and OxygenThis welder is using a torch that burns hydro-gen in oxygen. The heat produced is sointense that the torch can even be used under-water. The product of the reaction is water.


Elemental Good HealthDo you know that 60 elements are commonly

found in the human body? Slightly fewer thanhalf of them are essential for life, although sci-entists think most of the others play some role inlife processes. The elements currently known tobe essential are listed in the table below.

Although hydrogen, carbon, oxygen, and nitro-gen make up almost 96 percent of the mass of thehuman body, minerals are also essential for lifeprocesses.

Different roles Because calcium compounds makeup the hard parts of bones and teeth, calcium isneeded for their growth and maintenance. Iron isan important element because it is the active partof the blood’s hemoglobin molecule, which carriesoxygen to the cells. Fluorine helps in the formationand maintenance of teeth and may prevent osteo-porosis, which is the disintegration of bone. Youmay not know that magnesium is necessary for thefunctioning of nerves and muscles. So is potassium.Zinc and selenium are necessary for the activity ofenzymes needed for cell division and growth andfor the functioning of the immune system.

Chemistry

Exploring Further

1. Inferring Why might cooking foods in boilingwater reduce their mineral content?

2. Acquiring Information Why might consumingmore than the Recommended DietaryAllowance of minerals be harmful?

Minerals Non-mineralsF, Na, Mg, Si, P, S, Cl, H, C, O, NK, Ca, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, As, Se, Sn, I

Elements Necessary for Life

Selenium

Zinc and calcium

Chromium

Potassium

Different amounts Maintaining the proper level ofeach mineral in your body is important for health.Nutritionists have established amounts of these ele-ments that you should have in your daily diet. Theamounts are described as the Recommended Dietary

Allowance (RDA) and Estimated Safe and Ade-quate Dietary Intakes (ESAI). For example, theRDA for calcium for teenagers is 1200 mil-ligrams, while that of iodine is 150 micrograms(0.000 150 g). You may think an amount of 150micrograms can’t be too important, but it is cru-cial to the function of your thyroid gland, whichhelps control your metabolism and growth. Ifyou use iodized salt, which contains a littlepotassium iodide, you probably are getting theproper RDA. Likewise, eating a well-balanceddiet of the five food groups will maintain theproper levels of all the minerals that are elemen-tal to good health.

To find out more about the role of minerals inhealthy eating, visit the Chemistry Web site atchemistryca.com

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Using Clues to Make a CaseYou’ve seen that elements combine to form compounds whose properties

differ greatly from those of the elements themselves. Figure 4.9 showsanother example. On the submicroscopic level, this clue indicates thatatoms of elements react chemically to form some combinations that aremuch different from the original atoms. Also, if atoms of elements alwayscombined in the same way, it’s likely that all compounds would be similar.However, you’ve just studied three compounds that have greatly differingproperties. On the submicroscopic level, this clue indicates that atoms mustbe able to combine in different ways to form different kinds of products.With the knowledge of the structure of atoms that you learned in Chapters2 and 3, you can now examine the different ways that atoms can combine.

Figure 4.9Iron Reacting with ChlorineHere is another example of two elementsreacting to form a compound whose proper-ties are different from either element. Heat-ed steel (iron) wool reacts with chlorine gasto form iron(III) chloride, the brown cloudyou see in the flask.

SECTION REVIEWFor more practice

with solving problems,see Supplemental Practice Problems, Appendix B.

Understanding Concepts1. Use water as an example to contrast the proper-

ties of a compound with the elements fromwhich it is composed.

2. Classify the following substances as elements orcompounds.

a) table salt d) chlorine gasb) water e) carbon dioxide gasc) sulfur f) dry ice

3. Give an example of a clue that indicates thatatoms can combine chemically with each otherin more than one way.

Thinking Critically4. Comparing and Contrasting Using sodium

chloride, carbon dioxide, and water as exam-

ples, what can you say about the chemical reac-tivity of compounds compared to the elementsof which they are composed?

Applying Chemistry5. Using the Periodic Table Find the elements

that make up the compounds discussed in thissection. Which elements are metals? Which ele-ments are nonmetals? Compare the propertiesof the three compounds in terms of whethertheir component elements are two nonmetalsor a metal and a nonmetal.

chemistryca.com/self_check_quiz

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4.2

It’s time to return to the submicroscopicworld of atoms in order to explainhow and why atoms combine

to form such a variety of com-pounds. In the 19th century,before electrons were dis-covered, chemists tried tovisualize the way atoms com-bine to form compounds.Some early models picturedatoms with hooks that couldattach to the hooks of otheratoms. Today, chemists knowa great deal more about thestructure of atoms, so theyhave a much clearer model ofchemical combination.

When Atoms CollideWhen elements react, atoms of the elements must collide. It is what

happens during that collision that determines what kind of a compoundforms. How does the reaction of sodium and chlorine atoms to form saltdiffer from the reaction of hydrogen and oxygen atoms to form water?

When atoms collide with each other, what really comes into contact? Asyou learned in Chapter 2, the nucleus is tiny compared to the size of theatom’s electron cloud. Also, the nucleus of an atom is buried deep in thecenter of the electron cloud. Therefore, it is highly unlikely that atomicnuclei would ever collide during a chemical reaction. In fact, reactionsbetween atoms involve only their electron clouds.

When you studied the periodic table in Chapter 3, you learned that theproperties of elements repeat because the pattern of the outermost(valence) electrons repeats in each period. It is this arrangement ofvalence electrons of an atom that is primarily responsible for the atom’schemical properties. You will not be surprised, then, to learn that it is thevalence electrons of colliding atoms that interact. But what kinds of inter-actions between valence electrons are possible? For some additional cluesto what takes place when atoms combine to form compounds, look at agroup of elements with unusual properties—the noble gases.

How Elements Form Compounds

SECTION

Objectives✓ Model two types ofcompound formation:ionic and covalent atthe atomic level.

✓ Demonstrate howand why atomsachieve chemical sta-bility by bonding.

✓ Compare, usingexamples, the effect of covalent and ionicbonding on the physi-cal properties of compounds.

Review VocabularyMetalloid: an elementwith some physicaland chemical proper-ties of metals andother properties ofnonmetals.

New Vocabularyoctet rulenoble gas

configurationionionic compoundionic bondcrystalcovalent bondcovalent compoundmoleculeelectrolyteinterparticle force

SECTION PREVIEW


4.2 How Elements Form Compounds 131

Noble Chemical StabilityOf all of the elements, the elements of

Group 18 are a curious bunch. They arenotorious for their almost completelack of chemical reactivity. In fact,they have some practical uses justbecause they are not reactive, as youcan see in Figure 4.10. Despite the factthat all of these elements occur natural-ly in the environment, not a single com-pound of any of these elements has everbeen found naturally in the environment.

This group of unreactive elements used tobe called the inert gases because chemiststhought these elements could never react toform compounds. However, in the 1960s,chemists were able to react fluorine, under con-ditions of high temperature and pressure, withkrypton and with xenon to form compounds.Since that time, a few additional compounds ofxenon and krypton have been produced. Still, noone has been successful at synthesizing com-pounds of helium, neon, and argon. Now thatchemists know that these elements aren’t com-pletely inert, they are called noble gases.

Figure 4.10Uses for Noble Gases

� Incandescent lightbulbsare filled with noble gases,usually argon and krypton,to protect the filament.The tungsten filament getsso hot that it will reactwith all but the most inertelements.

Noble gases and mixtures of noble gasesare used in eye-catching light displays oftenreferred to as neon lights. These gases giveoff different colors of light when a high-voltage electric current passes throughthem. Neon produces a bright orangecolor, argon produces blue, and heliumgives off yellow-white. Lighting designersobtain different colors by adding mercuryor other substances to the gas mixture.Sometimes, colored glass is used. �

The Octet RuleThe lack of reactivity of noble gases indicates that atoms of these ele-

ments must be stable. Noble gases are unlike any other group of elementson the periodic table because of their extreme stability. As you know, theelements of a vertical group on the periodic table have similar arrange-ments of valence electrons. Each noble gas has eight valence electrons,except for helium, which has two. Because electron arrangement deter-mines chemical properties, the electron arrangements of the noble gasesmust be the cause of their lack of reactivity with other elements.


What does the electron arrangement of noble gases have to do with theway other elements react? Compare the electron arrangements of thenoble gases in Figure 4.11. Today, scientists know that atoms don’t havehooks. Rather, atoms combine because they become more stable by doingso. The modern model of how atoms react to form compounds is basedon the fact that the stability of a noble gas results from the arrangementof its valence electrons. This model of chemical stability is called the octetrule. The octet rule says that atoms can become stable by having eightelectrons in their outer energy level (or two electrons in the case of someof the smallest atoms). In other words, elements become stable by achiev-ing the same configuration of valence electrons as one of the noble gases,a noble gas configuration.

Ways to Achieve a Stable Outer Energy LevelIf atoms collide with enough energy, their outer electrons may

rearrange to achieve a stable octet of valence electrons—a noble gas con-figuration—and the atoms will form a compound. Remember that elec-trons are particles of matter, so the total number of electrons cannotchange during chemical reactions. Next, think about how valence elec-trons might rearrange among colliding atoms so that each atom has a sta-ble octet. There are only two possibilities to consider. The first is a transferof valence electrons between atoms. The second possibility is a sharing ofvalence electrons between atoms. The reactions discussed in Section 4.1provide good examples of both of these possibilities.

8e�

2e�8e�8e�

2e� 2e�8e�

Helium

Neon

Argon

Krypton Xenon

8e� 2e�8e�

8e� 2e�18e�18e� 18e�

Radon, Rn, is the lastmember of the noblegas group. It is aradioactive elementthat is formed by theradioactive decay ofradium. Although it isfound naturally onEarth, its presence isfleeting because itdecays rapidly to otherelements.

Figure 4.11Electron Arrangements of Noble GasesNote that the noble gas atoms have eight electrons inthe outer energy level. This arrangement causes themto be almost completely unreactive. The lone excep-tion to this octet arrangement is helium. The heliumatom has only one energy level, which can containonly two electrons.


Electrons Can Be TransferredYou know that when sodium and chlorine are mixed, a reaction occurs

and a compound, sodium chloride, forms. The photograph at the begin-ning of this chapter showed the macroscopic view of this reaction. Whatcan be happening at the atomic level? Begin by picturing a collisionbetween a sodium atom and a chlorine atom. Locate these atoms on theperiodic table. Sodium is in Group 1, so it has one valence electron. Chlo-rine is in Group 17 and has seven valence electrons.

In previous chapters, you have used Lewis dot structures to representan atom and its valence electrons. You will now be able to use these mod-els to show what happens when atoms combine. The electron dot struc-tures of the atoms are shown below.

How can the valence electrons of atoms rearrange to give each atom astable configuration of valence electrons? If the one valence electron ofsodium is transferred to the chlorine atom, chlorine becomes stable withan octet of electrons. Because the chlorine atom now has an extra elec-tron, it has a negative charge. Also, because sodium lost an electron, itnow has an unbalanced proton in the nucleus and therefore has a positivecharge.

It’s easy to see how chlorine has achieved a stable octet of electrons, buthow does sodium become stable by losing an electron? Look at the posi-tion of sodium on the periodic table. By losing its lone valence electron,sodium will have the outer electron arrangement of neon. Sodium’s stableoctet consists of the eight electrons in the energy level below the level ofthe lost electron. Figure 4.12 summarizes the way that sodium and chlo-rine react.

[ Cl ][Na]� �

� Cl Na � ˇ

ClNa

Figure 4.12The Reaction of Sodium and Chlorine AtomsThe transfer of an electron from a sodium atom to a chlo-rine atom forms sodium and chloride ions. Examine thediagram carefully to see how this transfer gives both ionsa stable octet.

8e� 8e� 2e�8e� 2e�8e�2e�8e� 7e�

Sodium� ion Chloride� ionSodium atom Chlorine atom

� �

2e�1e�

Take a deep breath. Anoble gas, argon (Ar),composes approximate-ly one percent of thatbreath of air you justtook.


Sodium atom + Chlorine atom ˇ Sodium ion + Chloride ionNa + Cl ˇ Na+ + Cl�

Number of protons 11 17 11 17Number of electrons 11 17 10 18Number of outer-level electrons 1 7 8 8

Table 4.2 Reaction of Sodium and Chlorine

Now that each atom has an octet of outer-level electrons, they are nolonger neutral atoms; they are charged particles called ions. An ion is anatom or group of combined atoms that has a charge because of the loss orgain of electrons. Ions always form when valence electrons rearrange byelectron transfer between atoms. A compound that is composed of ions iscalled an ionic compound. The transfer of a single electron changed areactive metal, sodium, and a poisonous gas, chlorine, into the stable andsafe compound, sodium chloride. Note that only the arrangement of elec-trons has changed. Nothing about the atom’s nucleus has changed. Thisresult is clear when you compare the atoms and ions in Table 4.2.

Ions Attract Each OtherRemember that objects with opposite charges attract each other. Once

they have formed, the positive sodium ion and negative chloride ion arestrongly attracted to each other. The strong attractive force between ionsof opposite charge is called an ionic bond. The force of the ionic bondholds ions together in an ionic compound.

Even the smallest visible grain of salt contains several quintillion sodi-um and chloride ions. Every positively charged sodium ion attracts allnearby negatively charged chloride ions and vice versa. Therefore, theseions do not arrange themselves into isolated sodium ion/chloride ionpairs. Instead, the ions organize themselves into a definite cube-shapedarrangement, as shown in Figure 4.13. This well-organized structure is acrystal. Solid substances are composed of crystals. A crystal is a regular,repeating arrangement of atoms, ions, or molecules.

� The macroscopic result ofthis submicroscopicarrangement is a cube-shaped salt crystal.

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Figure 4.13Crystal Structure of NaClPositive sodium ions attractnegative chloride ions to forma cube-shaped arrangementin sodium chloride. In thisarrangement, six chloride ionssurround every sodium ion,and six sodium ions surroundevery chloride ion. The forcesholding each ion in place areionic bonds. �

The Results of Ionic AttractionHow does ionic bonding affect the macroscopic properties of a sub-

stance? The cubic shape of a salt crystal is a result of the cubic arrange-ment of sodium and chloride ions. Given the strong attractive forcebetween sodium and chloride ions and the degree of organization amongthem, it’s not surprising that sodium chloride is a solid at roomtemperature. All particles of matter are in constant motion. Raisingthe temperature of matter causes its particles to move faster. Inorder for a solid to melt, its temperature must be raised until themotion of the particles overcomes the attractive forces and the crys-tal organization breaks down. Breaking the strong crystal structureof sodium chloride requires a lot of energy, which is the reason thatsodium chloride must be heated to more than 800°C before it melts.

When you press salt crystals, they don’t bend or flatten out.When enough force is applied, the salt crystals suddenly shatter.This hardness and brittleness provide macroscopic evidence for thestrength and rigidity of the submicroscopic structure of the saltcrystal. Trying to break an ionic crystal is like trying to break awell-laid brick wall, Figure 4.14.

The Formation of Ionic CompoundsA sodium atom reacts by losing an electron to form a sodium 1�

ion. A chlorine atom gains one electron to form a chloride 1� ion. Inthis MiniLab, you will consider other combinations of atoms.

Procedure 1. Cut three paper disks about

7 cm in diameter for each ofthe elements: Li, S, Mg, O, Ca,N, Al, and I. Use a differentcolor of paper for each element.Write the symbol of each ele-ment on the appropriate disks.

2. Select atoms of lithium andsulfur, and lay the circles side-by-side on a piece of corrugat-ed cardboard.

3. Using thumbtacks of one colorfor lithium and another colorfor sulfur, place one tack foreach valence electron on thedisks, spacing the tacks evenlyaround the perimeters.

4. Transfer tacks from the metallicatoms to the nonmetallic atomsso that both elements achieve

noble gas electron arrange-ments. Add more atoms ifneeded.

5. Once you have created a stablecompound, write the ion sym-bols and charges and the for-mula and name for the resultingcompound on the cardboard.

6. Repeat steps 2 through 5 forthe remaining combinations ofatoms.

Analysis1. Why did you have to use more

than one atom in some cases?Why couldn’t you just takemore electrons from one metalatom or add extra ones to anonmetal atom?

2. Identify the noble gas elementsthat have the same electronstructures as the ions produced.

2

135

Figure 4.14Ionic Crystals Are StrongThe crystal structure of anionic compound is a lot like awell-laid brick wall. It takes agreat deal of force to break abrick wall.


The Formation and Decomposition of Zinc Iodide

Compounds are chemical combinations of ele-ments. Many chemical reactions of elements toform compounds are spectacular but must berun under special laboratory conditions becausethey are dangerous. The reaction of sodium andchlorine to form sodium chloride pictured at thebeginning of Chapter 4 is a good example. If ele-ments react spontaneously to form compounds,that is a good indication that the compound stateis more stable than the free element state. Tobreak a stable compound down into its compo-nent elements, energy must be put into the com-pound. Electricity is often used as this energysource. The process of decomposing a compoundinto its component elements by electricity iscalled electrolysis.

ProblemCan a compound be synthesized from its ele-

ments and then decomposed back into its origi-nal elements?

Objectives

•Compare a compound with its component elements.

•Observe and monitor a chemical reaction.

•Observe the decomposition of the compoundback to its elements.

Materials10 � 150-mm test tubetest-tube racktest-tube holder

100-mL beakerspatulaplastic stirring rodzinciodine crystalsdistilled water9-V battery with terminal clip and leadstwo 20-cm insulated copper wires stripped at

least 1 cm on each end

Safety PrecautionsCAUTION: Iodine crystals are toxic and can stainthe skin. Use care when using solid iodine. The reac-tion of zinc and iodine releases heat. Always use thetest-tube holder to handle the reaction test tube.

1. Obtain a test tube and a small beaker. Placethe test tube upright in a test-tube rack.

2. Carefully add approximately 1 g of zinc dustand about 10 mL of distilled water to the testtube.

3. Carefully add about 1 g of iodine to the testtube. Record your observations in a table likethe one shown.

4. Stir the contents of the test tube thoroughlywith a plastic stirring rod until there is nomore evidence of a reaction. Record anyobservations of physical or chemical changes.

5. Allow the reaction mixture to settle. Using atest-tube holder, carefully pick up the test tubeand pour off the solution phase into the smallbeaker.

6. Add water to the beaker to bring the volumeup to about 25 mL.

7. Obtain a 9-V battery with wire leads and twopieces of copper wire. Attach the copper wires


to the wire leads from the battery. Make surethat the wires are not touching each other.

8. Dip the wires into the solution and observewhat takes place. Record your observations.

9. After two minutes, remove the wires from thesolution and examine the wires. Again, recordyour observations.

1. Observing and Inferring What evidence wasthere that a chemical reaction occurred?

2. Comparing and Contrasting How did youknow the reaction was complete?

3. Making Inferences What term is used todescribe a reaction in which heat is given off?How can you account for the heat given off inthis reaction?

4. Checking Your Hypothesis What evidence doyou have that the compound was decomposedby electrolysis?

5. Drawing Conclusions Why do you think thereaction between zinc and iodine stopped?

1. What role did the water play in this reaction?

2. Do you think zinc iodide is an ionic or cova-lent compound? What evidence do you haveto support your conclusion?

3. The formula of zinc iodide is ZnI2. Use Lewisdot structures to show how it forms from itselements. Hint: Zinc atoms have two valenceelectrons.

Step Observations

3. Addition of iodine to zinc

4. Reaction of iodine and zinc

8. Electrolysis of solution

9. Examinationof wires


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Representing Compounds with FormulasRather than writing out the name sodium chloride every time you refer

to it, you can use a much simpler system. You can write its chemical for-mula, NaCl. The formula of a compound tells what elements make up thecompound and how many atoms of each element are present in one unitof the compound. Water is written as H2O. This formula means that waterconsists of two hydrogen atoms combined with one oxygen atom. Whensodium and chlorine atoms react, they form ions, which then arrangethemselves into a crystal. The ratios of the elements that make up saltcrystals do not change, and so the formula NaCl does not change.

You have seen how the electrons are transferred from sodium to chlo-rine to form a strong crystal arrangement. There are many other ioniccompounds, as you will see later in Chapter 5. Now, look at anotherexample from Section 4.1 to see a different way that atoms can combineto achieve a stable outer level of electrons.

Electrons Can Be SharedIn Section 4.1, you learned that the reaction of hydrogen and oxygen

forms water. What happens when hydrogen and oxygen atoms collide?Hydrogen has only one valence electron. Oxygen, a Group 16 element, hassix valence electrons. Could these atoms achieve the stable electron con-figuration of a noble gas by transfer of electrons? If the oxygen atomcould pick up two more valence electrons, it would have a stable octet—the noble gas configuration of neon.

What about hydrogen? Could a hydrogen lose its single valence elec-tron? Your first inclination might be to treat hydrogen just like sodium,but be careful. If hydrogen loses an electron, it is left with no electrons,and that isn’t the electron structure of a noble gas. Maybe hydrogen couldgain an electron so that its electron arrangement is like that of helium.But, both atoms cannot gain electrons.

Colliding atoms transfer electrons only when one atom has a strongerattraction for valence electrons than has the other atom. In the case of

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sodium and chlorine, chlorine attracts sodium’s valence electron strongly,whereas sodium holds its electron weakly. Therefore, the electron movesfrom sodium to chlorine, forming positive and negative ions in theprocess. You will learn more about this process and the factors that influ-ence it in Chapter 9. In the case of hydrogen and oxygen, neither atomattracts electrons strongly enough to take electrons from the other atom.When atoms collide with enough energy to react, but neither attracts elec-trons strongly enough to take electrons from the other atom, the atomscombine by sharing valence electrons.

To understand how water forms, start by looking at the electron dotstructures of hydrogen and oxygen.

Hydrogen requires one more electron to have the same electronarrangement as helium, while oxygen requires two more electrons to haveneon’s arrangement. Hydrogen and oxygen can share one electron fromeach atom. This sharing is shown by placing two dots representing elec-trons between the atoms.

This arrangement makes hydrogen stable by giving it two valence elec-trons, but it leaves oxygen with only seven valence electrons. Oxygen’soctet can be completed by sharing an electron with another hydrogenatom. (This explains why water has the formula H2O.)

Just as in the case of the formation of sodium chloride by ionic bond-ing, all the parts present before the reaction are still there after the reac-tion. What has changed in the combining of hydrogen and oxygen atoms?The valence electrons no longer reside in the same positions. This is whatalways happens in a chemical reaction; electrons rearrange. Figure 4.15summarizes the reaction between hydrogen and oxygen.

O H H O HH

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Figure 4.15Formation of Water by Electron SharingThe stability of the atoms in a water moleculeresults from a cooperative arrangement inwhich the eight valence electrons (six from oxy-gen and one each from two hydrogens) are dis-tributed among the three atoms. By sharing anelectron pair with the oxygen, each hydrogenclaims two electrons in its outer level. The oxy-gen, by sharing two electrons with two hydro-gens, claims a stable octet in its outer level. Bythis method, each atom achieves a stable noblegas configuration.


Electron Sharing Produces MoleculesThe attraction of two atoms for a shared pair of electrons is called a

covalent bond. Notice that in a covalent bond, atoms share electrons andneither atom has an ionic charge. A compound whose atoms are heldtogether by covalent bonds is a covalent compound. Water is a covalentcompound. Although water is made up of hydrogen and oxygen atoms,these have combined into water molecules, each having two hydrogenatoms bonded to one oxygen atom. A molecule is an uncharged group oftwo or more atoms held together by covalent bonds. Sometimes chemistsrefer to covalent compounds as molecular compounds. The terms meanthe same thing. Figure 4.16 compares a compound with ionic bonds to acompound having covalent bonds.

Figure 4.16Comparing an Ionic and aCovalent CompoundIron(III) chloride is a typicalionic compound. It is crys-talline at room temperature,melts at a high temperature(300°C), and dissolves inwater. �

More Than Two Electrons Can Be SharedWhen charcoal burns, carbon atoms collide with oxygen to form CO2.

Carbon is in Group 14 and has four valence electrons. Oxygen is in Group16 and has six valence electrons.

Carbon and oxygen are like hydrogen and oxygen when it comes to thequestion of sharing or transferring electrons. Neither atom is able toattract electrons away from the other atom. In fact, two nonmetallic ele-ments usually achieve stability by sharing electrons to form a covalentcompound. On the other hand, if the reacting atoms are a metal and anonmetal, they are much more likely to transfer electrons and form anionic compound.

� Ethanol, also known as ethyl alcohol, is a typical covalentcompound. It is a liquid at room temperature but evapo-rates readily into the air. Ethanol boils at 78°C and freezesat �114°C. Unlike many covalent compounds, ethanol dis-solves in water. In fact, ethanol sold in stores as rubbingalcohol contains water.


HISTORY CONNECTION

Connecting to Chemistry

1. Inferring Muchresearch has beendone on developinghydrogen-fueledcars and trucks.Why would such acar be almostentirely nonpollut-ing? What factorsdo you think mightaffect the public’s

acceptance of such avehicle?

2. HypothesizingWhat do you thinkis the reason that somuch energy is pro-duced when hydro-gen reacts with oxy-gen? Recall whathappens when theseatoms bond.

Hydrogen’s Ill-fated LiftsImagine the surprise of the citizens of

Paris on that morning of December 1, 1783.There was Jacques Charles and his assistantgliding above their rooftops in a basket sus-pended from a large balloon. Charles and hisassistant had filled the balloon with hydrogenand became the first humans to ride in sucha lighter-than-air vehicle. By World War I,hydrogen-filled balloons were being used tocarry military personnel aloft to observetroop movements.

In 1936, Germany launched the Hinden-burg, a rigid airship originally designed to usehelium for lift. Helium is slightly less buoyantthan hydrogen, but it is a noble gas. Thus, itdoes not react with anything, whereas hydro-gen burns explosively in air. However, theGermans had to continue the use of hydrogenin the airship because they had no source ofhelium. In the following year, the Hindenburgcarried more than 1300 passengers in trans-atlantic flights. While attempting to dock inLakehurst, New Jersey, on May 1, 1937, theairship exploded and burned when the hydro-gen ignited. Thirty-six people were killed. Theaccident was the final chapter in the use ofhydrogen in lighter-than-air vehicles.

The chemical reaction The reaction thatdestroyed the Hindenburg was the burning ofhydrogen gas.

2H2(g) � O2(g) ˇ 2H2O(g)

Once ignited, thisreaction occursspontaneously. How-ever, the reaction canbe controlled, such asin the main engine ofthe space shuttle.

Space shuttle fuelThe main engine ofthe space shuttle

booster uses liquid hydrogen and oxygen asfuel. These materials are stored in separatesections of a huge external fuel tank attachedbeneath the shuttle. The energy released dur-ing the reaction thrusts the shuttle into orbit.

During the launch of the space shuttleChallenger on January 28, 1986, the reactionbecame uncontrolled. One minute and 13seconds after takeoff, the external tank andshuttle exploded, killing the shuttle’s sevencrew members. The accident was caused bydefects in the design of O-rings that joinedsections of the solid-fuel booster engines,which are attached to the sides of the shuttle.

Challenger explosion

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Consider the reaction of carbon and oxygen to form CO2. Look at elec-tron dot structures for the participating atoms below.

Can you arrange the 16 valence electrons from these three atoms toproduce a molecule in which all three atoms have a stable configuration?You know that at least one bond must exist between the carbon and eachoxygen, so start there. Here’s an approach to the puzzle. Have each oxygenshare an electron with carbon as in the following dot structures.

This arrangement gives carbon six electrons and each oxygen seven, butstill no atom has an octet. What else can be done? There’s no law in chem-istry that says atoms must bond by sharing only one pair of electrons.What happens if they share two pairs? You now have double covalentbonds, as shown in this dot structure.

Now, count all the electrons around each atom, including the ones thatare shared. You’ll see that each atom has a stable octet. Study the result ofthis electron sharing in Figure 4.17. By sharing electrons, the three atomsachieve a more stable arrangement than they had as three separate atoms.A molecule of carbon dioxide is more stable than one carbon atom andtwo oxygen atoms. As with water, the molecule of carbon dioxide is differ-ent from the sum of its parts. The macroscopic properties of carbon diox-ide are a result of the unique properties of carbon dioxide molecules, notthe properties of carbon or oxygen atoms.

O C O

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Figure 4.17Electron Sharing in CO2

When one carbon atom and two oxygen atoms react, thecarbon atom shares two pairs of electrons with each oxy-gen. This arrangement gives all atoms a stable octet.

143

How do ionic and covalent compounds compare?Now you can relate the submicroscopic models of the formation of

NaCl, H2O, and CO2 to their macroscopic properties mentioned in Sec-tion 4.1. When elements combine, they form either ions or molecules. Noother possibilities exist. The particles change dramatically, whether theychange from sodium atoms to sodium ions or hydrogen and oxygenatoms to water molecules. This change explains why compounds have dif-ferent properties from the elements that make them up.

Explaining the Properties of Ionic CompoundsThe physical properties of ionic compounds are directly related to the

fact that ionic compounds are composed of well-organized, tightly boundions. These ions form a strong, three-dimensional crystal structure. Thismodel of the submicroscopic structure explains the general observationthat ionic compounds are crystalline solids at room temperature. Just likeNaCl, these solids are generally hard, rough, and brittle. Ionic compoundsusually have to be heated to high temperatures in order to melt thembecause the attractions between ions of opposite charge are strong. Ittakes a lot of energy to break the well-organized network of bound ions.Compare their appearance with the properties described above. Sometypical ionic compounds are shown in Figure 4.18.

Figure 4.18Comparing Ionic CompoundsHere are some other typical ionic com-pounds. Note that all are solids at roomtemperature and are soluble in water. Notall ionic solids are soluble in water,though.

� Sodium hydrogen carbon-ate (NaHCO3), commonlyknown as baking soda, isalso called sodium bicar-bonate.

Potassium chloride (KCl) isused as a salt substitutebecause it has a taste similarto that of salt. �

� Copper(II) sulfate (CuSO4) issometimes used to treat thegrowth of algae in swimmingpools and in water-treatmentplants.


Figure 4.19Comparing Interpar-ticle Forces in Ionic and Covalent CompoundsIn an ionic compoundsuch as lithium bromide(LiBr), interparticle forcesare strong because of theattraction between ionsof opposite charges. �

� Butane (C4H10) is the fuel commonly found in plastic,disposable lighters. Butane is a covalent compound,but the molecules have no electrical charge, so theattraction between them is weak. In fact, if thebutane were not held under pressure in the lighter, itwould immediately boil away to a gas.

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Another physical property of ionic compounds is their tendency to dis-solve in water. When they dissolve in water, the solution conducts electric-ity, as you saw in Figure 4.2. Ionic compounds also conduct electricity inthe liquid (melted) state. Any compound that conducts electricity whenmelted or dissolved in water is an electrolyte. Therefore, ionic com-pounds are electrolytes. In order to conduct electricity, ions must be freeto move because they must take on or give up electrons. Ionic compoundsin the solid state do not conduct electricity because the ions are lockedinto position. Ionic compounds become good conductors when they melt.This is evidence that the ions are less bound and free to move in the liq-uid state.

Explaining the Properties of Covalent Compounds As with ionic compounds, the submicroscopic model of the formation

of covalent compounds explains many of the properties of these com-pounds. In particular, you can use this model to explain why typical cova-lent compounds such as H2O and CO2 have properties so different fromionic compounds.

In order to explain these differences, consider the submicroscopic orga-nization of covalent compounds. Covalent compounds are composed of

molecules. As you learned in this chapter, the atoms that composemolecules are held together by strong forces—covalent bonds—that make the molecule a stable unit. The molecules themselveshave no ionic charge, so the attractive forces between moleculesare usually weak.

The forces between particles that make up a substance arecalled interparticle forces. These forces are illustrated in Figure4.19. It is the great difference in the strength of interparticle forcesin covalent compounds compared to those of ionic compoundsthat explains many of the differences in their physical properties.


Whereas all ionic compounds are solids at room temperature, manycovalent compounds are liquids or gases at room temperature. Note, how-ever, that many covalent compounds—sugar, for example—will formcrystals if there is enough attractive force between molecules. In Chapter9, you will learn why some molecules attract each other. Many of thecovalent compounds that are solids at room temperature will melt at lowtemperatures. Examples are sugar and the compounds that make up can-dle wax and fat. Compare the properties of some covalent substances inFigure 4.20. Molecular (covalent) compounds do not conduct electricityin the pure state. Many covalent compounds, such as those in gasolineand vegetable oil, do not dissolve in water, although others such as sugarwill dissolve. The solubility of covalent compounds in water varies, but ingeneral, covalent compounds are usually less soluble in water than ioniccompounds. What accounts for these differences?

Figure 4.20Comparing Covalent CompoundsCovalent compounds are composed of molecules in which atoms are bonded by electron sharing. Because of weak interparticle forces between molecules, covalent compounds tend to be gases or liquids at room temperature. They tend to be insoluble in water, although many are extremely soluble.

See page 864 in Appendix F for

Mixing Ionic and Covalent Liquids

Lab

Table sugar (C12H22O11) is called sucrose. It is an example of a covalent compoundthat is a crystalline solid soluble in water.

Gasoline and oil are mixtures of covalentcompounds. Spilled oil does not dissolve inwater, but instead floats on the water inthin layers.

In places where natural gas is not available,many people use propane (C3H8) to heat theirhomes and cook food. It is delivered to busi-nesses and homes in pressurized tank trucks.

Candle wax and butter are mixtures ofcovalent compounds. Because their mole-cules are larger and heavier, they formsolids but melt at low temperatures.

The Rain Forest PharmacyLong ago, Samoan healers dispensed a tea

brewed from the bark of a native rain forest tree,Homalanthus nutans, to help victims of yellowfever, a viral disease. Researchers have since identi-fied the bark’s active ingredient, prostratin. Pros-tratin is now being investigated for possible use asa drug for other viral diseases.

Using plants and products from plants for me-dicinal purposes has a long history. Such commondrugs as aspirin, codeine, and quinine were origi-nally derived from plants. However, only about0.5 percent of all plant species have been studiedintensively for chemical makeup and medicinalbenefit. Because most of the world’s 250 000species of flowering plants live in rain forests,researchers are now combing these areas of densevegetation for new substances to fight diseases.

Screening plants One method of collectingplants for drug research is to select many differentspecies of plants from one locale and test them forpossible medicinal uses. Because of the great diver-sity of plants, researchers hope to find an evengreater diversity of chemical substances. Agenciessuch as the National Cancer Institute routinely usethis screening method. Another method of screen-ing plants for research is the phylogenetic survey,in which researchers study close relatives of plantsalready known to produce beneficial medicinalsubstances. Researchers hypothesize that similari-ties in the evolution of these plants may have pro-duced similar properties in the biochemical sub-stances these plants produce.

Ethnobotanical approach Interest is growing inanother method of screening plants for drugresearch. This method is based on the work ofethnobotanists. Ethnobotanists study the medicinaluses of plants by native peoples. Ethnobotanistshave learned that the sophisticated knowledge con-cerning the use of medicinal plants possessed bythese people can be used to identify plants that may

be important for future research. The discovery ofprostratin is an example of the ethnobotanicalapproach to screening plants for potential use indrug research.

The future Isolating and testing of substancesfrom plants may take years. Within this time, rainforests may be significantly altered. Most rainforests are found in developing countries wherethe citizens wantto work for bet-ter lives andgrow enoughfood to be self-sufficient.Because ofdeforestation toobtain lumber aswell as space foropen-field agri-culture, rain for-est environmentsare disappearing.Some researchersand environ-mental scientistsare setting uppreserves within the rain forests to maintain theforests’ biodiversity. Others are attempting to growrapidly disappearing plant species in rain forestnurseries. With these methods, scientists are tryingto maintain the complex rain forest ecosystems forfuture research.

1. Inferring Why might drug researchers investi-gate a plant that has few pests as a potentiallyuseful plant?

2. Acquiring Information How might researchinto the pharmaceutical value of rain forestsubstances have an impact on the present andfuture uses of rain forests?

Analyzing the Issue


Chemistry and


Interparticle Forces Make the DifferenceInterparticle forces are the key to determining the state of matter of a

substance at room temperature. You already know that ions are held rigid-ly in the solid state by the strong forces between them. Because there aremuch lower interparticle forces in covalent compounds, their moleculesare less tightly held to one another. Therefore, they are more likely to begases or liquids at room temperature.

Because there are no ions in covalent compounds, you do not expectthem to be electrical conductors. Ionic compounds tend to be soluble inwater while molecular compounds do not. This difference is alsoexplained by interparticle forces. Ions are attracted by water molecules,but many covalent molecules are not and, therefore, do not dissolve. Solu-bility in water and the nature of water solutions is a major topic in chem-istry. You will learn more about solutions in Chapter 13.

Connecting IdeasNow that you know that two major types of compounds exist, you may

wonder where else you come into contact with these compounds on aday-to-day basis. In Chapter 5, you’ll look at more examples of ionic andcovalent compounds, including compounds more complex than the sim-ple ones used as examples in this chapter. You’ll also learn the importantpractical skill of naming and writing the formulas of compounds, as wellas how to identify a few special categories of compounds such as acids,bases, and organic compounds.

SECTION REVIEWFor more practice

with solving problems,see Supplemental Practice Problems, Appendix B.

Understanding Concepts1. Explain why sodium chloride is a neutral com-

pound even though it is made up of ions thathave positive and negative charges.

2. Describe two ways in which atoms become sta-ble by combining with each other. Compareand contrast the compounds that result fromtwo kinds of combinations.

3. Why do you think that sodium chloride has tobe heated to 800°C before melting, but candlewax will start to melt at 50°C?

Thinking Critically4. Making Predictions Look at the following dia-

gram and explain how you can determinewhether the compound being formed is ionicor covalent. Do you think it is more likely to bea solid or a gas at room temperature? Explain.

Applying Chemistry5. Using Lewis Structures Potassium metal will

react with sulfur to form an ionic compound.Use the periodic table to determine the numberof valence electrons for each element. Draw aLewis dot structure to show how they wouldcombine to form ions. How would you writethe formula for the resulting compound?

Br [ Br ] Br [Ca]Ca� �

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UNDERSTANDING CONCEPTS1. How is the compound sodium chloride differ-

ent from the elements of which it is composed?

2. If you tried to breathe CO2, you would suffo-cate. Why, then, is CO2 essential to all life onEarth?

3. Why is water essential to life on Earth?

4. Why are Group 18 elements no longer calledinert?

5. Describe the electron arrangement that makesan atom stable. Why is helium stable with adifferent arrangement?

6. Describe two processes by which elements cancombine to form stable compounds. Name thetype of bonding that results from each process.

7. Compare the formation of NaCl from sodiumand chlorine to the formation of CO2 fromcarbon and oxygen. In what ways are they sim-ilar? In what ways do they differ?

4.1 The Variety of Compounds■ When compounds form, they have properties

that differ greatly from the properties of theelements of which they are made.

■ Sodium, a dangerously reactive metal, reactswith chlorine, a poisonous gas, to form thestable compound sodium chloride, table salt.

■ Carbon, usually found as a black solid, reactswith oxygen in the air to form carbon diox-ide, an unreactive gas that produces the fizzin soda pop.

■ Hydrogen, the lightest gaseous element, reactsexplosively with oxygen in the air to formwater, a stable compound on which all lifedepends.

■ The properties of compounds differ widelybecause of differences in what happens totheir constituent atoms when they form.

4.2 How Elements Form Compounds■ Atoms become stable by reacting to achieve

the outer-level electron structure of a noblegas (Group 18).

■ One way to achieve a stable electron structureis to transfer electrons from one atom toanother, thus forming charged ions of oppo-site charge. The ions attract each other toform a crystal. A compound formed in thisway is an ionic compound.

■ Another way that atoms can achieve stabilityis by sharing electrons with other atoms toform molecules. A compound formed byelectron sharing is a covalent compound, alsocalled a molecular compound.

■ Two atoms can share more than one pair ofelectrons.

■ In ionic compounds, the interparticle forcesare attractions between ions of opposite elec-trical charge. These forces are much strongerthan the interparticle forces between mole-cules of covalent compounds.

■ It is the differences in strength of interparticleforces that account for many of the differ-ences in physical properties between ionicand covalent compounds.

VocabularyFor each of the following terms, write a sentence that showsyour understanding of its meaning.

covalent bond ionic bondcovalent compound ionic compoundcrystal moleculeelectrolyte noble gas configurationinterparticle force octet ruleion

REVIEWING MAIN IDEAS

CHAPTER 4 ASSESSMENT

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Chapter 4 Assessment 149

8. Sodium reacts with fluorine to form sodiumfluoride, NaF, a common additive in toothpastethat prevents tooth decay. The electronexchange in this reaction is the same as that forsodium and chlorine. Use the format shownin Table 4.2 to analyze this reaction.

9. What are three general properties of ioniccompounds? How are these properties relatedto the submicroscopic structure of the com-pounds?

10. Why would you never expect to find Na2Cl as astable compound?

11. How is a sodium ion different from a sodiumatom? From a neon atom?

12. The term isoelectronic is used to describe atomsand ions that have the same number of elec-trons. Which of the following are isoelectronic?Na+, Ca2+, Ne, K, O2�, P3�

13. What are three general properties of covalentcompounds? How are these properties related tothe submicroscopic structure of the compounds?

14. How do the basic particles of ionic and cova-lent compounds differ?

APPLYING CONCEPTS15. A bag of pretzels lists the sodium content of the

pretzels. Considering that sodium reacts vio-lently with water, why don’t you explode whenyou eat these sodium-containing pretzels? Whatdoes the label really mean to a chemist?

16. Hydrazine is a compound with the formulaN2H4. What kind of compound is hydrazine?

Describe the formation of hydrazine fromnitrogen and hydrogen atoms.

17. What does it mean to say that a chemical reac-tion is a rearrangement of matter?

18. When two atoms collide, what determineswhether they will react by transferring elec-trons or by sharing electrons?

Everyday Chemistry19. Carbon monoxide gas (CO) binds strongly to

the iron atom of hemoglobin in blood. Howdoes this action cause harm to the body?

History Connection20. Do you think a rocket powered by hydrogen

and oxygen causes a great deal of atmosphericpollution? Explain.

Chemistry and Society21. List two economic factors that contribute to

the loss of rain forests.

THINKING CRITICALLYApplying Concepts22. Several times in this chapter, it was stated that

elements combine to form compounds. Con-sidering what you learned in Section 4.2, whatdoes the word combine mean in chemistry?

Interpreting Chemical Structures23. MiniLab 2 Which of these compounds could

the following “tack model” represent: magne-sium chloride, potassium sulfide, calciumoxide, or aluminum bromide?

Ne atomNa� ion

2e�8e�2e�8e�


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Observing and Inferring24. MiniLab 1 Brass is an alloy of copper and

zinc. Neither metal is magnetic. People whobuy brass antiques at auctions, shows, or shopsoften carry a small magnet with them. Whatdo you think they learn with this magnet?

Making Predictions25. ChemLab In the electrolysis of zinc iodide,

what limits your ability to recover all of theoriginal zinc and iodine?

Observing and Inferring26. ChemLab Why does the reaction between zinc

and iodine speed up when the iodine is dis-solved in water?

Interpreting Data27. ChemLab In the synthesis of zinc iodide, the

zinc was added in excess. What does in excessmean? What two experimental observationsdid you make to confirm this?

CUMULATIVE REVIEW28. When carbon in charcoal burns in air to form

CO2, is the process endothermic or exother-mic? How do you know? (Chapter 1)

29. Describe the particle structure of an atom ofpotassium. Assume an atomic mass of 39.When a potassium atom forms an ion by react-ing with chlorine, how will its structurechange? (Chapter 2)

30. How does the electron arrangement of Group18 elements relate to their chemical properties?(Chapter 3)

SKILL REVIEW31. Interpreting a Graph Look at the solubility

graph and answer the following questions.

a) Which of the compounds on the graph are ionic compounds? How do you know?

b) How many grams of KBr will dissolve in 100 g of water at 60°C?

c) People think that salt is very soluble in water.Is that thinking accurate? Defend your answer.

WRITING IN CHEMISTRY32. Write an article about salt mining in the Unit-

ed States. Find out where the major saltdeposits are located and how they came to bethere. Describe what happens to the salt after itis removed from the mine. Tell how table salt ismade, and find out what else it containsbesides sodium chloride.

PROBLEM SOLVING33. When hydrogen and chlorine combine chemi-

cally, they form hydrogen chloride, HCl.Hydrogen chloride is a gas at room tempera-ture; it becomes a liquid if it is cooled to�85°C. On the basis of this evidence, do youthink that hydrogen chloride is ionic or cova-lent? Explain.

240

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120

100

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60

40

20

0

10 20 30 40 50 60 70 80 90 100

Temperature (�C)

Solu

bili

ty (

g/1

00 g

H O

)

NaClOKNO

NaCl

KBr

2

3

3



Standardized Test Practice

1. How do the properties of the compound mag-nesium oxide (MgO) compare with the prop-erties of the elements magnesium and oxygen?

a) The compound has completely differentproperties from the two elements.

b) The compound has the same properties asthe two elements.

c) The compound has similar properties as thetwo elements.

d) The properties of compounds and elementscannot be compared.

2. Why does oxygen make up such a large per-centage of Earth’s crust?

a) Large quantities of oxygen are trapped insoil and rock deep beneath Earth’s surface.

b) Large quantities of water, which is made ofoxygen, are trapped beneath Earth’s surface.

c) Low temperatures far beneath Earth’s surfacefreeze atmospheric oxygen into solid form.

d) Oxygen is highly reactive and bonds withother substances forming solid compounds.

3. According to the chart above, how does thedensity differences of water in different statesensure the survival of aquatic life in Canada?

a) In the winter, ice floats on lakes and ponds,insulating aquatic creatures from cold temperatures.

b) In the winter, ice sinks to the bottom of lakesand ponds, causing organisms to hibernate.

c) In the summer, water vapor sinks to thebottom of ponds and lakes, carrying needed oxygen.

d) In the winter, water vapor sinks to the bottom of ponds and lakes, carrying neededoxygen.

4. Why is the gas hydrogen rarely found in ele-mental form on earth?

a) Hydrogen is a rare element.b) Hydrogen is highly reactive.c) Hydrogen easily forms compounds.d) Hydrogen is an unreactive atmospheric gas.

5. What would be the result if nitrogen gasreplaced noble gases inside incandescent light bulbs?

a) The tungsten filament in the bulb wouldburn with a dimmer light.

b) The tungsten filament in the bulb wouldcombust.

c) Chemical reactions would make the lightbulb last longer.

d) Chemical reactions would reduce the lifes-pan of the bulb.

6. The octet rule says

a) atoms become less stable with eight elec-trons in their outer energy level.

b) atoms become more stable with eight electrons in their outer energy level.

c) atoms will change their configurations andbecome a noble gas.

d) atoms will become more reactive and chemically bond with noble gases.

7. A compound composed of electrically chargedatoms is called a(n)

a) octet compound.b) chemically stable compound.c) ionic compound.d) covalent compound.

Standardized Test Practice 151

Practice Under Test-Like ConditionsAsk your teacher to set a time limit. Then do all ofthe questions in the time provided without referring-to your book. Did you complete the test? Could youhave made better use of your time? What topics doyou need to review? Show your test to your teacherfor an objective assessment of your performance.

Test Taking Tip

chemistryca.com/standardized_test

Water Density

State of Water Density (g/mL)

Solid 0.917

Liquid 1.000

Gaseous (25�C) 0.0008

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Chapter 4: Formation of Compounds - Glencoe...spoon, the crystals shatter. This shattering shows...

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