Saence Crnñent~J sta~idardsS Students know the structure of theatom and know it is composed of protons,neutrons and electrons.

What You’ll Learnp Describe the structure of

the atom and whereprotons, neutrons, andelectrons are located.

p compare the mass, size,and charge of the threebasic particles of an atom.

p Describe two observationsthat Dalton’s atomic theorysupported.

Why It’s ImportantAn understanding of thenature of the atom is thefirst step toward Learningwhat the world is made of.

Vocab~sbrymatteratomnucleusprotonneutronelectron

Rev~ew ve@abidarymass: a measure of theamount of matter in anobject (p. 11)

Matter is made of tiny particles called atoms.

Real-World Rndh~g CornieCtiO~ How can you figure outwhat’s inside a wrapped box without opening it? Exploring theatom is like exploring that box. Atoms can’t be observed directlywith your eyes, so how have scientists learned about what’sinside them?

What is the current atomic model?Would it surprise you to learn that the chair you are sitting

on and the air you breathe are made up of the same thing? Theworld you live in is made of matter. Matter is anything thathas mass and takes up space. Things you can see, such as yourchair, and things you can’t see, such as air, are matter. Matter isdifferent from light, heat, and sound. These are forms of energy.Matter is made up of atoms. An atom is a very small particlethat makes up all matter. Only recently have scientists been ableto see the surface of an atom.

Inside the AtomIn the early 1980s, a powerful new instrument called the

atomic-force microscope was invented. The atomic-force microscope can magnify an object up to one million times. This magnification is great enough for the surfaces of individual atoms tobe seen, as shown in Figure 1. If further magnification were possible, you might be able to see inside an atom. You probablywould be surprised to find that most of the atom is empty space.In this space, particles are moving. No one has ever seen insidean atom, so how do scientists know what atoms are made of?

Atoms—Basic Unitsof Matter

lii

174 chapter 4

parts of Atoms—Protons, Neutrons, and ElectronsMany experiments performed by scientists during the last 200

years have established what is inside an atom. An atom is mostlyempty space surrounding a tiny nucleus. The nucleus is a regionthat is located at the center of an atom and contains most of theatom’s mass. Figure 2 shows that the nucleus contains positivelycharged particles and neutral particles. A positively chargedparticle located in the nucleus is a proton. A neutral particle,which has no charge, located in the nucleus is a neutron. Atomsalso contain particles called electrons. An electron is a negativelycharged particle that moves in the space surrounding the nucleus.

The Size of AtomsAs tiny as atoms are, electrons, protons, and ‘neutrons are even

smaller. The data in Table 1 show that protons and neutrons haveabout the same mass. Electrons have only about 1/2,000 the massof a proton or a neutron. If you held a textbook and placed a paperclip on it, you wouldn’t notice the added mass because the mass ofa paper clip is small compared to the mass of the book. In a similar way, the masses of an atom’s electrons are negligible comparedto an atom’s mass. An atom’s protons and neutrons are packedtightly into a tiny nucleus. Visualize the nucleus as the size of anant. How large would the atom be? Amazingly, the atom would bethe size of a football stadium.

Election~C

Nucleus

Figure 2 An atom of lithium hasthree electrons, three protons, andfour neutrons.Describe the locations of the protons, theneutrons, and the elections.

or

Lithium atom

WoRD ORIGINnucleusfrom Latin nucula; meanslittle nut

:ro‘agis top05-

Particle Charge Mass (9)

:mhl. PiOpérities ~ • atom__rticle.

Mass (amu)

Proton +1 1.6727 x 10—24 1.007316

Neutron 0 1.6750 x 10—24 1.008701

Electron —1 9.110 x 1O~~— ~

0.000549

Lesson 1 • Atoms—Basic Units of Matter 175

ACADEWIIC VOCABULARYaccuraze(adjective) free from erroror mistakeThe scale at the doctor’s office isaccurate.

Is there historical evidence of atoms?The idea that matter is made of tiny indivisible particles was

I proposed as early as 400 B.C. But experimental evidence to supportthe idea of atoms was not available until the seventeenth andeigthteenth centuries. Actually, the current understanding of,atomic structure has developed over the last several hundred/years.Each time new evidence becomes available, the model of at91micstructure becomes clearer and more accurate.

Democrftus and the AtomGreek philosopher Democritus (c. 460—370 B.c.) was the firs~

person to use the word atom. Atom comes from the Greek word\atoma, which means ccindivisible.” Indivisible describes somethingthat cannot be divided into smaller pieces. Democritus provided amuch more detailed idea of the atom than any that ever had.been

I proposed. He thought that atoms were very small, solid sphereswith no holes and no empty space inside.

Democritus argued that atoms were indivisible. He imaginedcutting a piece of matter into smaller and smaller pieces. Hehypothesized that eventually he would come to a point at whichhe could not cut any more pieces. He would have come to a piececonsisting of one atom that could not be divided.

The student in Figure 3 is illustrating Democritus’s experiment.She is cutting a piece of aluminum in half, and again in half, overand over again. The pieces become smaller and smaller, but eachis still aluminum. Suppose she could continue to cut beyond thepoint where the pieces are too small to see. She would eventuallyreach a point where the final piece is just one indivisible aluminum atom. An atom is the smallest piece that still is aluminum.

What was Democritus’s idea of the atom?

Figure 3 Democritus’s ideaswere based on reasoning ratherthan experiments. This pictureis recreating Democritus’s concept of the indivisible atom.

iTh Chapter 4 • understanding the Atom

The Law of Conservation of MassWhat happens to the atoms in substances during a chemical

reaction? A chemical reaction is a process in which the atoms inthe starting materials rearrange to form products with differentproperties. French scientist Antoine Lavoisier (AN twan luhvWAH see ay) (1743—1797) conducted experiments that helpedanswer this question. Lavoisier placed a carefully measured massof solid mercury(II) oxide into a sealed container. When he heatedthe container, he saw something different. The red powder of

• rnercury(IT) oxide had changed into a silvery liquid and a gas. Thesilvery liquid ~tas mercury. Lavoisier established that the gas produced was a component of air. This component is oxygen. In hisexperiments Lavoisier recorded the masses of the starting materials and of the products. He found that the total mass of the start-

• ing materials was always the same as the total mass of theproducts. Experiments such as this led to the recognition of thelaw of conservation of mass. This law states that the mass of theproducts always is the same as the mass of the starting materials.

What data did Lavoisier record in his experiments’

The Law of Definite ProportionsBy 1799, J. L. Proust had completed a different series of experi

ments. Proust analyzed a variety of pure compounds to determinetheir compositions. He found that any pure compound alwayscontains the same elements in the same proportion by mass. Thisprinciple is called the law of definite proportions. The law appliesto any compound no matter where the sample comes from or howlarge or small it might be. Figure 4 illustrates that water’s composition is the same whether the sample comes from your kitchen sinkor from a glacier on Mars. Water always contains two hydrogenatoms and one bxygen atom. The law of definite proportions provided evidence to support the work of John Dalton as he developed

I

ACADEMiC VOCABULARYproportion(noun) the relation of onepart to another or to the wholeA large proportion of thepeople present were students.

his atomic model.

Figure 4 The law of definite proportions could beillustrated in a similar wayfor every pure substance.

I

Sulphur Magnesia Lime Soda Potash

0000©Strontian Barytes Iron Zinc Copper

Dalton’s Atomic Model

H N Co PHydrogen Nitrogen Carbon Oxygen Phosphorus

5 Mg Ca Na KSulfur Magnesium Calcium Sodium Potassium

Sr Ba I Z CuStrontium Barium Iron Zinc Copper

English schoolteacher and scientist John Dalton (1766—1844)was interested in the physical properties of gases. Like Lavoisierand Proust, Dalton made careful measurements of starting materials and products in a number of chemical reactions. To record hisresults accurately, he invented symbols for the known elements.As Figure 5 shows, these are more complex than modern symbols,but they helped scientists communicate better.

Dalton gathered information from his own observations andfrom the findings of other scientists. He put these results together.Dalton then proposed a new atomic theory. His atomic theoryconsists of five principles. Notice that the second principle isanother way of stating the law of conservation of mass.

1. All matter is made up of atoms.2. Atoms are neither created nor destroyed in chemical

reactions.3. Atoms of different elements combine in whole-number ratios.4. Each element is made of a different kind of atom.5. The atoms of different elements have different masses and

properties.

Which principle states the law of conservation

‘~S1~~ of mass?Dalton brought all that was known about the atom into a reason

able theory. Other scientists then could continue his work. Theycould improve Dalton’s theory or prove that it was wrong. Overtime, Dalton’s theory was modified as new evidence became available. Scientists now know that nuclear reactions can convert atomsof one element into atoms of a different element. We also knowthat atoms are made of smaller particles.

Hydrogen Azote

Dalton’s Atomic Symbols

~00Carbon Oxygen Phosphorus

Modern Atomic Symbols

Figure 5 Dalton created pictures for each ofthe elements. These werehelpful for writing downhis results, just as ourmodern symbols are.

178 chapter 4 • understanding the Atom

Looking Back at the LessonThe ancient Greeks taught that matter consists of tiny indivisi

ble particles called atoms. However, the Greeks couldn’t prove theexistence of atoms. It wasn’t until the seventeenth century that scientists began to look for evidence of the atom. Their experimentsdemonstrated the law of conservation of mass and the law of definite proportions. With these important ideas, Dalton describedhis atomic model. Dalton’s model started the development of themodern model of the atom. That model consists of even tinier particles called protons, neutrons, and electrons. You’ll read moreabout these particles in Lesson 2.

1. Explain the difference betweena neutron and a nucleus. 0

2. An atom contains equal numbersof ______ and ______

0Understanding Main Ideas

3. Which has no charge? 0A. electronsB. protonsC. neutronsD. nucleus

4. Name the particles that makeup an atom and tell wherethey are located. 0

5. Explain in your own wordswhat is meant by the law ofdefinite proportions. 0

6. Describe how Lavoisier wasable to demonstrate the law ofconservation of mass. 0

7. Show that the ratio of thenumber of atoms of hydrogen to the number of atomsof oxygen in the compoundwateris2tol.

8. Compare Copy and fill in thegraphic organizer below tocompare the mass and thevolume of a proton with themass and the volume ofan electron. 0

Mass VolumeProtonNeutron

Applying Science

9. Design an experiment thatconfirms the law of conservation of mass. 0

10. Assess the reasons whyDalton, not Democritus, iscredited with being the“Father of the Atom.” 0

I

Using Vocabulary

Standards CheckSummarizeCreate your own lesson summary as you write a script fora television news report.

1. Review the text after thered main headings andwrite one sentence abouteach. These are the headlines of your broadcast.

2. Review the text and write2—3 sentences about eachblue subheading. Thesesentences should tell who,what, when, where, andwhy information abouteach red heading.

3. Include descriptive detailsin your report, such asnames of reporters andlocal places and events.

4. Present your news reportto other classmates aloneor with a team.

sciencóriiJneFor more practice, visit StandardsCheck at ca8.msscience.com.

Lesson 1 • Atoms—Basic Units of Matter 179

Vacab aspectral lineenergy levelelectron cloud

Review Vocabularyelectromagneticspectrum: the entire rangeof electromagnetic wavesof different wavelengths(p. 428)

the Atom_________ Scientists have put together a detailed model of

atoms and their parts.

Real-World Reading Connection Imagine you are adetective. You go to a crime scene. You can only make observations and analyze clues because there are no witnesses to thecrime. Similarly, scientists make observations and gather cluesthat help them build a model of the atom even though theycannot see inside one.

How were electrons discovered?Since the time of the ancient Greeks, around 400 B.c.,

scientists thought atoms were the smallest units of matter. Butmore than 2,000 years later, in the late 1800s, a series of experiments led scientists to a better understanding of atoms. Theylearned that atoms are made of even smaller particles. Many ofthese experiments used a cathode-ray tube similar to theone inFigure 6. Cathode rays are given off at the cathode, which is anegatively charged disk. A cathode ray is a stream of particlesthat can be seen when an electric current is passed through avacuum tube. The cathode rays travel to the positively chargeddisk at the other end of the tube.

Figure 6 What is the positively charged disk called?

Sde~ce Couftert~“j Standards~Students know the structure of theatom and know it is composed of protons,neutrons, and electrons.

I

Discovering Parts of

What You’ll Learnp. Describe the arrangement

of electrons, protons, andneutrons within an atom.

p. Explain how Rutherforddeveloped his model ofthe atom.

p. List the evidence thatshowed the existence ofelectrons, protons, andneutrons.

p. Compare Thomson’s,Rutherford’s, and Bohr’smodels of the atom.

Why It’s ImportantThe structure of the atom isthe key to understandingchemistry.

Figure 6 The electron was discovered using a cathode-raytube similar to the one in the photo.

Low pressure gas.......~Cathode ray

/A

Opening conne~edto a vacuum pump

Battery

182 Chapter 4 • understanding the Atom

Thomson’s ExperimentsIn 1897, English scientist J. J. Thomson wanted to find out how

electric currents affect cathode rays. He changed the cathode-raytube by putting charged metal plates above and below the tube, asshown in Figure 7. One plate was positively charged. The otherplate was negatively charged. Thomson found that the cathode raysdid not follow a straight path down the tube. Instead, they bent inthe direction of the positive plate. Recall that opposite chargesattract one another and like charges repel one another. Thomsonconcluded that the particles in a cathode ray must have a negativecharge. He named the newly discovered particles electrons.

Thomson also was able to use the cathode-ray tube to measurethe mass of the charged particles. To his surprise, he found thatthe mass of an electron is much smaller than the mass of an atom.He concluded that atoms are not indivisible, as Dalton had proposed. Thomson also realized that atoms must contain positivecharges to balance the negative charges of the electrons. His findings must have been true because atoms are neutral.

What did Thomson learn from his experiment about%~ the mass of electrons?

Thomson’s Atomic ModelWith this new information, Thomson proposed a new model

for the atom. Instead of a solid, neutral sphere that had the samematter all the way through, Thomson’s model of the atom contained both positive and negative charges. He proposed that anatom was a positively charged sphere. The electrons were mixedevenly through the sphere, similar to how raisins are mixed incookie dough. Figure 8 shows a cutaway view of an atom in whichthe small spheres represent the electrons.

SCIENCE USE V. COMMON USEchargeScience Use a definitequantity of electricityThe electron has a negativecharge.Common Use an expense,cost, or feeWhat is the charge foradmission?

Thomson’s Model

Positivelycharged plate

+

Figure 7 Using this experimental setup, J. J.Thomson found that cathode rays were attractedto the positively charged plate above the tube.Infer What must be the charge on the cathode rays?

Negativelycharged plate

Figure 8 Thomsonsuggested that electrons mixed evenly intothe positively chargedspherical atom.

Positivelycharged sphere

Negativelycharged electrons

Lesson 2 • Discovering Parts of the Atom 183

I Rutherford—Discovering the NucleusThe discovery of electrons stunned scientists and made them

want to find out more about the atom. Ernest Rutherford was aresearch student of J. J. Thomson at the Cavendish Laboratory inEngland. Rutherford was interested in understanding the structureof Thomson’s model of the atom. By 1911, Rutherford had a laboratory and students of his own. Rutherford expected his studentsto find that electrons and positive charges were mixed together inan atom. But as you will read in the next section, what they foundwas another surprise.

The Good Foil ExperimentTwo of Rutherford’s students set up a series of experiments to

see if Thomson’s model was correct. Particles with a positivecharge, called alpha particles, were shot through a sheet of thingold foil. The apparatus is shown in Figure 9. A detector beyondthe gold foil glowed with a spot of light wherever the particles hit.Rutherford thought the positive charge of the gold atom wasspread evenly throughout the atom. At no place would the speeding alpha particles come upon a charge large enough to stronglyrepel them. Figure 10 shows a close-up view of what Rutherfordmight have expected. The alpha particles would speed through thefoil with only slight changes in their paths. This was the result predicted by the Thomson model.

Why did Rutherford think the alpha particles wouldmove straight through the gold foil?

~i/©~fl~;r~

(noun) the collecting of information about a particularsubjectShe did research on atoms at thelibrary.

Electron

0!

:0 0 0

0

Figure 9 Predicted Outcome The path of an alphaparticle is shown by a burst of light where the particle hits.

Alpha particlesource Expected path of

0 0 IA0

Evenlydistributedpositive charge

— Path of alpha particles

Figure 10 Rutherfordexpected most particlesto crash through the goldfoil with little change indirection.

184 Chapter 4 • Understanding the Atom

An Unexpected ResultWhat happened was another surprise. Notice in Figure 11 that

most of the alpha particles did pass directly through the foil withno bending of their paths. But sometimes, particles were stronglybounced off to the side. Astoundingly, one particle in about 8,000bounced straight backward. Rutherford later described his amazement by saying, “It was quite the most incredible event that hasever happened to me in my life. It was almost as incredible as ifyou had fired a fifteen-inch shell at a piece of tissue paper and itcame back and hit you.” Thomson’s model of the atom did notwork. How did Rutherford know this?

Interpreting the EvidenceRutherford realized that if positive charges were spread evenly

in atoms, all the alpha particles would have passed through the foilwith only a small change in direction. He also recognized that apositively charged particle could be bounced directly backward.This would happen only if the alpha particle bumped into something with much greater mass and positive charge than the alphaparticle itself. Think about this similar situation. Imagine that youare running very fast. If you bump into a dangling leaf, you won’teven notice. You just keep running along a straight path. But if youcrash into a tree branch, you will very likely be knocked off yourcourse. A head-on collision with a tree trunk might even bounceyou straight backward. Figure 12 shows an artist’s view of howRutherford must have visualized charged particles bouncing offthe nucleus of a gold atom.

Alpha particlesource cJncePts In

To see animation of Rutherford’s experimentvisit ca8.msscience.corn.

Electron

-z&~----

/Emptyspace

I.

I

FI

Nucleus(dense positivecharge)

— Path of alpha particles

Figure 12 Some alphaparticles must have hit amassive particle in thegold atom.Explain how Rutherfordknew that Thomson’s modelof the atom was not correct.

Lesson 2 • Discovering Parts of the Atom 185

Evidence

Most of the alpha particles passed rightthrough the gold foil.

The charged particles that bounced backcould not have been knocked off courseunless they had hit a mass much largerthan their own.

A few of the alpha particles bounced The positive charge is concentra1directly back. in a small space within an atom]

Rutherford’s Atomic ModelUsing the observations of his students, Rutherford drew some

conclusions, which are summarized in Table 2. Most of the alphaparticles passed directly through the gold atoms. For this to happen, the atoms must have contained mostly empty space. Becausesome alpha particles were strongly deflected from their paths,those particles must have come near a large positive charge. Veryfew alpha particles were bounced completely backward. Those particles that did bounce back must have collided with a mass havinga large positive charge.

Drawing on these conclusions, Rutherford revised Thomson’smodel of the atom. Figure 13 shows Rutherford’s new atomicmodel. Notice that most of the volume of an atom is empty space.At the center is the nucleus. An atom’s electrons move very fast inthe empty space surrounding the nucleus.

Thinking about Rutherford’s results, American poet RobertFrost wrote a very short poem, The Secret Sits.

“We dance round in a ring and suppose,But the Secret sits in the middle and knows:’

What do you think sits in the middle? What dances roundin a ring?

~~1

How doelectrons move?

Conclusion

ILprocedure~flI~i. complete a lab

form.safety

Most of the mass of an atom isconcentrated in a small spacewithin the atom

1. Draw a straight Linedown the center of a10-cm x 10-cm blockof foam with a ruler.

3. Break 20 toothpicks inhalf. Poke the halvesinto the foam so theyare Like the nucleus ofan atom.

4. Use round, dried peasas electrons. Aim andflick the peas downthe center tine on theblock.

5. Make a diagram toshow where the electrons came out Use aprotractor to measurethe angle the electronsmade compared to thecenter tine, which isthe path they wouldhave followed if theydid not hit any atoms.

Analysis1. Describe how your

arrangement of toothpicks was like the nucleiof atoms in a block ofmetal. Why did thetoothpicks representjust the nuclei insteadof the whoLe atoms?

i. Describe problemsyou had with thisexperiment.

[igure 13 Rutherford’s atomincluded a positively chargednucleus. Electrons moved in the 1:space around the nucleus.

Nucleus

186 chapter 4 • understanding the Atom

Rutherford’s Model

completing Rutherford’s ModelRutherford used cathode-ray tubes for other experiments. He

wanted to find out about the positive charge in an atom’s nucleus.The result of these experiments was the discovery of anotherparticle~ called the proton. A proton is an atomic particle witha +1 charge. Rutherford and his students knew the approximatemass of a proton. They could determine how many protons werein atoms. However, they couldn’t account for all of the mass of anatom. Rutherford predicted that an atom contains another undiscovered particle. But, it wasn’t until 1932 that the existence of theneutron was proved by English physicist James Chadwick. A neutron is a neutral atomic particle with a mass similar to a protonbut has no charge. An atom’s neutrons occupy the nucleus alongwith its protons. Neutrons were difficult to find because they haveno charge, unlike protons and electrons. Both protons and electrons are deflected by a magnetic field.

Compare and contrast protons and neutrons.

WoRD OaoG~r~iprotonfrom Greek protos; meansfirst

Weakness in the Rutherford ModelRutherford’s model explained much of the experimental evi

dence, but it also brought up new questions. How are electronsarranged in atoms? How can differences in the chemical behaviorof different elements be explained? For example, why does oxygenreact easily with metals? Why is argon not very reactive? One cluecame from the observation that elements give off colored lightwhen heated in a flame. Figure 14 shows the bright colors of theelements barium, sodium, strontium, and potassium when theyare placed in a flame. Each element creates its own flame color.Some elements are used in fireworks to produce the brilliant colorsof a display. Rutherford’s model could not explain where this lightcomes from.

Figure 14 Scientistswanted to know whatcauses the colored lightwhen elements areheated.Identify the color producedwhen barium is placed in aflame.

Lesson 2 • Discovering Parts of the Atom 187

LA

—

Bohr and theHydrogen Atom

In 1918, Danish scientist Niels Bohr began toanswer some of the questions about Rutherford’smodel. Rutherford had proposed that electronscould move around the nucleus at any distancefrom the nucleus. He thought electrons mightmove like the ball on a string, shown in the topillustration of Figure 15. In the figure, a boy hastied a soft sponge ball to a long string and isslowly twirling it above his head. The ball doesn’thave much energy and moves in a small circle.Suppose the boy releases more string and twirlsmore energetically. The bottom illustration ofFigure 15 shows that the ball moves in a largercircle farther from his head. Depending on theenergy the boy provides and the length of thestring he releases, the ball could circle his head atany distance up to the length of the string. Bohrshowed that Rutherford’s idea that electronscould circle the nucleus at any distance wasincorrect. His experiments convinced him thatelectrons did not behave like a twirling ball thatcould travel in circles of any diameter. Electronscould only move in circles with certain diameters, like the planets that circle the Sun. Like theplanets~ an electron’s path around the nucleushad a definite radius.

What did Bohr compare the path ofan electron to?

Bohr came to this conclusion by studying thehydrogen atom. He chose hydrogen because it isthe simplest element, with only one electron.Bohr was interested in the light given off byhydrogen gas when it is excited. Atoms becomeexcited when they absorb energy by being heatedin a flame or by electricity. Figure 16 shows theelement neon in an advertising sign. The red lightis produced when neon is excited by electricity.

Bohr wanted to know what was happeninginside an atom to cause it to release energy in theform of colored light. Was there a connectionbetween the light and the structure of theatom?

Figure 15 By gradually letting out morestring and twirling faster, the.ball will travel inincreasingly large circles.

‘1

Short String and Low Energy

4

0

Longer String and Greater Energy

iss Chapter 4 • understanding the Atom

The Spectrum of HydrogenTo understand the light given off by excited atoms, think about

the rainbow of colors you see when ordinary light moves througha prism. The colors red, orange, yellow, green, blue, and violetblend into each other in a continuous spectrum of colors. Recallthat colors at the red end of the spectrum have longer wavelengthsand lower energies. Colors at the violet end have shorter wavelengths and higher energies. Visible light is just a small section ofall the possible wavelengths in the electromagnetic spectrum.Ultraviolet rays have shorter wavelengths and higher energies thandoes visible light. Infrared rays have longer wavelengths and lowerenergies than does visible light. You cannot see ultraviolet rays orinfrared rays. The electromagnetic spectrum is the whole range ofelectromagnetic waves with different energies and wavelengths.

t~~t~fl Arrange visible light, infrared rays, and ultraviolet rays~ in order of their energies,from lowest to highest.

How is the energy of electrons related to the electromagneticspectrum? The light given off by excited hydrogen atoms doesn’thave a continuous spectrum of colors. Instead, hydrogen gives offlight of specific colors, as shown in Figure 17. The narrow bands ofred, green, blue, and violet light given off by an excited hydrogenatom are called its spectral lines.

Visible spectrum

ACAD~MlC VocA~uLARYvisible(adjective) capable of beingseen with the eyeOn a clear night, the stars arevisible in the night sky

Lesson 2 • Discovering Parts of the Atom 189

Figure 16 Neon gas is excitedby electricity and glows red.

IHydrogen

Neon

Figure 17 The Tight given off byhydrogen and neon is not continuous like the rainbow of color produced by white light. Each elementhas its own specific spectral lineswith specific energies.

IlL

Spectrall Lines and Energy LeveilsA spectral line is a single wavelength of light that can be seen

when the light from an excited element is passed through a prism.If you compare the spectrum of hydrogen to the spectrum of lightin Figure 17, you’ll notice that hydrogen has a red line and then agreen line. Between those lines, all the colors you see in the spectrum of sunlight are missing. The same is true for the colorsbetween hydrogen’s green line and its blue line. Each color is a di~Iferent wavelength and energy. Bohr knew that if the electrons inan excited atom could have every possible energy, they would giveoff light just like the spectrum of sunlight. But hydrogen gives offonly specific wavelengths of light. That means that an excitedhydrogen atom releases only certain amounts of energy. Becauseelectrons only can have certain amounts of energy, they can movearound the nucleus only at distances that correspond to thoseamounts of energy. These regions of space in which electrons canmove about the nucleus of an atom are called energy levels.

Energy levels can be compared to the ladder shown in Figure 18.You can stand on the ladder only at the level of each step, notbetween levels, similarly, electrons can be only at certain energylevels, not between levels. If an electron absorbs energy from aflame or from an electric current, it can jump from a lower energylevel to a higher energy level. When the electron falls back downfrom a higher energy level to a lower one, it releases energy. InFigure 19, energy levels are compared to a staircase in which thesteps are not evenly spaced.

Figure 19Electronsclimb anenergy staircase as theymove to upperenergy levels.They give offenergy in theform of lightwhen they fallback down.

k tween the spectrum of~ What is the difference Lie , h ~~ hydrogen and the spectrum of sun ig t.

Figure 18 A person canmove on a ladder only bystanding on the steps. Anelectron can move in anatom only by jumpingfrom energy level toenergy level.

Is

Energyabsorbed

i~o Chapter 4 • understanding the Atom-J

Bohr’s Atomic Mod&Bohr proposed that what he had learned from studying the

hydrogen atom applied to all atoms. Like Rutherford’s model,Bohr’s atomic model contains a nucleus. Electrons move in circlesaround the nucleus. But, as shown in Figure 20, the electrons canmove only in circles with certain diameters. Each of these circles,called energy levels, has its own energy. The energy levels are at setdistances from the nucleus and have specific energies.

Electrons in the Bohr AtomIn Bohr’s model of the atom, each energy level can hold a given

number of electrons. The way that electrons are placed in energylevels is similar to the way students might fill the rows of seats inan auditorium. Students fill the front row closest to the stage first.Then they fill the second row. When the second row is filled, theycontinue to the third row and beyond until all students are seated.Maybe the last occupied row is full of students. Or, maybe it isonly partly filled.

Similarly, electrons fill the lowest energy level first. The lowestenergy level is closest to the nucleus and can hold two electrons.When this first energy level is full, electrons begin to fill the second level. The second energy level can hold eight electrons. Whenthe second energy level is filled, electrons go to the next higherlevel. The last occupied energy level may or may not be completelyfilled. Figure 21 shows how electrons are placed in the elementswith atomic numbers 1—10.

(7~. Figure 21 Which two atoms have filled energy levels?~Which atom has four electons in its outer energy level?

Helium2

He

C

lithium Beryllium3 4Li Be

:., /

,, S. S

S 0 3

Boron Cathon Nitrogen Oxygen Fluorine Neon5 6 7 8 9 10B C N 0 F Ne

S S S -~ C. C. SS. S •• S Se’ S cc S ~i, £ I ~e S

0 0 0 0 0 -00 0 0 5 0 5 S

. • •

I Figure 20 In Bohr’satom, electrons orbit thenucleus at set distances.

4V

Bohr’s Model

Figure 21 As the number of electrons increases fromone to ten, two electrons fill the lowest energy level.Then, eight electrons fill the second energy level.

Hydrogen

H

0

Lesson 2 • Discovering Parts of the Atom 191

U

3

e3

3 S

0 • 3 4 4 4

C

e .3

Helium Neon Lithium Sodium

[~igure22 Heliumandneon, with filled outerenergy levels, are unreactive. Lithium and sodium,with one electron each intheir outer energy levels,are very reactive.

Bohr’s Model and Chemical PropertiesWhy do elements have different chemical properties? Bohr’s

model provided an answer. The clue to the chemical properties ofan element is in the number of electrons in its outer energy level.Elements with the exact number of electrons to fill their outermostenergy level are unreactive. Figure 22 shows that helium and neonhave filled outer energy levels. This means these elements do notcombine with other atoms to form compounds~ or new substances.As you might guess, elements with partially filled outer energy levels are likely to form compounds. Figure 22 shows that lithiumand sodium have one electron in their outermost energy levels.Both are very reactive metals.

Limitations of Bohr’s Atomic ModelBohr’s model explained much about chemical behavior. He pro

posed that energy levels were like circular orbits. That idea seemedto work for the simple hydrogen atom, but it did not work formore complex elements. If electrons don’t, travel in circular orbits,how do they move in the space around the nucleus?

The Electron CloudToday, scientists think of an electron in an atom as being in an

electron cloud. An electron cloud is a region surrounding anatomic nucleus where an electron is most likely to be found. Electrons move rapidly from one place to another. They can be anywhere. But they are more likely to be closer to the nucleus thanfarther aw~y because of the attraction of the negatively chargedelectrons for the positively charged nucleus. Figure 23 shows a diagram of an electron cloud. The electron cloud is much larger thanthe diameter of the nucleus. If the nucleus were the size of aperiod~ the atom would have a diameter of about 5 m. Figure 24summarizes how knowledge about the atom has increased throughexperiments.

—===-~=-=-.-

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Figure 23 An electroncan be anywhere, evenoutside the edges of thisillustration. The electroncloud shows only wherethe electrons are mostlikely to be found.

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192 chapter 4 • ~~derstanding the Atom

Rutherford

Ernest Rutherford’s experiments showed thatmost of an atom’s mass is squeezed into a tinynucleus. In the remaining space, electronsmove in orbits of all possible diameters.

Bohr

Neils Bohr used information from line spectrato define the orbits of electrons as having certain definite diameters.

Electron Cloud

In the current model, electrons occupy aspace around the nucleus, but it is impossibleto tell where an electron is at any particulartime.

[i~gure 24 Development of Atomic Models The history of the development of thecurrent model of the atom is an example for how science works. Models are proposedand tested. As more is learned, models are revised to fit the new observations.

John Dalton’s picture of the atom was a simple, neutral sphere of indivisible matter thatwas the same throughout. Atoms of differentelements, however, were different from oneanother.

J. J. Thomson’s amazing discovery of the electron showed that atoms were ndt indivisible.They contained negative electrons and positivecharges to make them neutral.

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Lesson 2 • Discovering Parts of the Atom 193

Bow has the atomic model changed?For Dalton, atoms were simple spheres of matter. Now the

model of the atom is an electron cloud. How did this change takeplace? J. J. Thomson showed that the atom contains even smallerparticles~ called electrons. Rutherford proved that the atom has anucleus packed with protons. Chadwick found out that neutronsalso share space in the nucleus. Neils Bohr hypothesized that electrons move in energy levels. Today, scientists know that themotions of electrons can’t be known. The electron cloud modelshows only where electrons are most likely to be. Nevertheless, thismodel has been useful to chemists.

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Create your own lesson Using vocabulary 7. compare Copy and fill in the5ummary as you organize ~ graphic organizer below. Corn-an outline. [ 1.~ pareThomson’s; Rutherford’s;own words. 0 and Bohr’s atomic models to

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Atomic Model ~ pOPI 194 chapter4 • Understanding the Atom

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Science Content‘%J StandardsW Students know how to use the periodictable to identify elements in simplecompounds.Wstudents know each element has aspecific number of protons in the nucleus(the atomic number) and each isotope of theelement has a different but specific numberof neutrons in the nucleus.W Construct appropriate graphs fromdata and develop quantitative statementsabout the relationship between variables.

Explain how elementsdiffer.

~ Identify elements andatomic masses on theperiodic table.

).Explain how two isotopesdiffer.

$. Explain how two ionsdiffer.

Why It’s ImportantTo understand their chemicalbehavior, it’s important toknow how the atoms of anelement can differ.

Vocabtdaryelementatomic numbermass numberisotopeaverage atomic massion

~‘aview VocabWaryperiodic table: table of theelements arranged accordingto repeated changes inproperties (Grade 5)

Elements, Isotopes, andIons—How Atoms Differ

Atoms of a particular element always have thesame number of protons.

Real-World Reading Connection You touch a doorknoband get a shock. Electrons are moving between the doorknoband your hand. Electrons can move from one atom to anotheratom. Why does this happen?

Different Elements—DifferentNumbers of Protons

Early Greek philosophers thought of matter as combinationsof four basic elements. These elements were earth, water, fire,and air. Today, an element is defined differently. An element isa pure substance made from atoms that all have the same number of protons. All atoms of the same element have the samenumber of protons. For example, all aluminum atoms have13 protons. That means that all atoms that have 13 protons arealuminum atoms. The number of protons in the atom of anelement is the element’s atomic number. Figure 25 shows someelements with their atomic numbers.

Lesson 3 • Elements, Isotopes, and Ions—How Atoms Differ 195

What You’ll Learn

Figure 25 Atoms of each of these elements havetheir own identifying atomic numbers.

Atomic Number and thePeriodic Table

How can you find out how many protons an element has if youdon’t know its atomic number? You can use the periodic table ofelements, shown in Figure 26. Elements in the periodic table arearranged horizontally in order of increasing atomic numbers. Theelements are also arranged vertically in groups with similar chemical properties. In almost all periodic tables, the block for each element gives the element’s atomic number, name, and symbol.Often, the atomic mass also is included.

fl Figure 26 How many elements are in the first row of~44ê~ the periodic table? How many are in the second and

third rows?

In the periodic table shown in Figure 26, the blocks for most ofthe elements are colored light blue. These elements are metals.Notice that Enost of the elements are classified as metals. Theblocks for nonmetals are yellow. These blocks are located-at theright side of the table. Between the metals and nonmetals arethe semimetals. These elements are represented by the greenblocks. Semimetals are elements that have properties similar tothose of both metals and nonmetals.

2

____ Figure 26 PeriodicTableof Elements 10Hydrogen The periodic table is a way of organizing the Helium

2 elements and understanding the relationships 13 14 15 16 17

~ among their chemical properties.Lithium Beryllium J Boron Carbon Nitrogen Ouygeri Finorine Neon

3 4 I-I 5 6 7 ~ 9 10Li Be B C N 0 F Me

Sodium M~~iUf j Aleminom Silicon Phoophnreu Sulfur Chlorine Argon11 12 j~3 4 5 6 7 ~ 9 10 11 12 13 14 15 16 17 18Na Mg Al Si P 5 Cl Ar

Peounolom Calcium Scandium Titanium Vanadium Chinmium Manganese Iron Cobalt Nickel Cnpper Zinc Gallinm Germanium Annenic Selenium Bromine Krypton19 20 21 22 23 24 25 26 27 28 29 30 31 32- 33 3~ 35 35IC Ca Sc Ti V Cr Mn Fe -Co Ni Cu Zn Ga Ge As Se Br Kr

Rubidiem Strontium Yttrium Zircoriom Niobium Mai~num Technetium Ruthenium ~edi,m Palludiom Silver Cadmium Indlem 6n Antimony Tellurium iodine Xenon37 36 39 40 41 42 43 ~ 45 46 47 40 49 5° 51 52 53 ~4Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd in Sn Sb To B Xe

Cesium Barium Lanthanum hafnIum tantalum tuegnien Rhenium Onmlum Iridium PlatInum Gold Mercury Thallium Lead Bismuth Po!onium Antatine Radon

6 ~5 56 57 72 73 74 7S 76 77 70 75 0° 01 02 03 04 05 06Cs Ba Is Hf Ta W Re Os Ir Pt Au Hg Ti Pb Si Po At Rn

Francium Radium Actinium ruareofurdium Oabnium Seaborgium Bohritm Ilusslum Meiluerium nurmnoaduuim nncntgenhm Uounbi,m unanquadoaoe

7 87 00 ~9 104 105 106 107 108 109 ~110 111 112 114Fr Rn Ac Rf Db Sq Bin 11$ Mt Os Rg Uub Uuq

1- — r~-[ y~-~Curium ~.,o,nd~lam~um rt,mothinm Samarium ~oplum e,~i~m Terbium Gynponuium Holmium tnbium Thulium Ynrrblnm Laieuium

~ 59 60 61 62 63 64 65 66 67 68 69 70 71Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Thorium Fnoartiniom Umnlum Neptunium Plutonium Ametltitm Curium torkeliom Califotnium Einsteinium Feinliom Muodrimlum Nuhelium Lawencltm99 91 92 93 94 95 B6 97 ~ g~ 105 101 102 103Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Ir

i~o Chapter 4 • Understanding the Atom

Isotopes—Different Numbersof Neutrons

In Lesson 2, you read that an atom’s nucleus contains neutronsin addition to protons. Unlike protons, neutrons have no charge.Atoms of the same element always have the same number of protons. However, atoms of the same element sometimes have different numbers of neutrons.

Protons, Neutrons, and Mass Number- The atomic number of an element tells you the number of protons in its atoms. But, how can you know the number of neutronsan atom has? To find out, you need to know the atom’s mass number. An atom’s mass number is the sum of the number of protonsand neutrons the atom has. Subtract the number of protons (theatomic number) from the mass number to calculate the number ofneutrons.

Number of neutrons = mass number — number of protons

Figure 27 shows the nuclei of two different atoms of neon.The atomic number of both atoms is 10, so each atom shouldhave 10 protons (+). But one atom has a mass number of 20. Theother’s mass number is 22. A third type of neon atom also occursnaturally, but in very small amounts.

~Figure 27 Count the number of neutrons in each~model. Check your results using the equation above.

You may have read previously that an element is a substancecomposed of one kind of atom. Now you are learning that atomsof an element can have different numbers of neutrons in theirnuclei. But, as you will read in Chapters 7 and 8, the chemicalbehavior of an element doesn’t depend on the contents of itsnucleus. All atoms of the same element act the same chemicallybecause they have the same number of electrons.

Neon-22 nucleus

-ALesson 3 • Elements, Isotopes, and Ions—How Atoms Differ 197

Neon-20 nucleus

Figure 27 Count the protons andneutrons in each nucleus. Showthat the mass number equals thenumber of protons plus the number of neutrons.

w~.: ~ .

Recall that all the atoms of a particular element have the sameatomic number. Having the same atomic number means that theseatoms contain the same number of protons. In contrast, you haveread that not all atoms of an element have the same mass number.This means that atoms of the same element can have differentnumbers of neutrons. Neon was the example shown in Figure 17.Atoms of the same element that contain different numbers of neutrons are called isotopes.

Because most elements have more than one isotope, each element has an average atomic mass. The average atomic mass of anelement is the weighted average mass of the mixture of an element’s isotopes. The most common of carbon’s isotopes has sixprotons and six neutrons. If you add the number of protons andneutrons (6 + 6), you find that this isotope has a mass number of12. Another of carbon’s isotopes has seven neutrons. Add to thenumber its 6 protons and you have its mass number of 13. Can yousee why carbon’s third isotope has a mass number of 14? The symbols for these three isotopes are C-12, C-13, and C-14. What otherproperty, shown in Table 3, is different for one of the three carbonisotopes besides the number of neutrons?

What are the atomic number and mass number of~M~4$-~ the most common isotope of carbon?

Using IsotopesCarbon-14 is radioactive. Radioactive isotopes have unstable

nuclei that break down and release particles, radiation, and energy.This property makes an isotope useful for a variety of purposes.Carbon-14 is useful for dating bones, wood, and charcoal up to75,000 years old. Geologists use uranium-238 to determine the ageof rocks. In hospitals and clinics, radioactive isotopes help diagnose and treat many medical conditions. In Figure 28, you canfind out what a tracer element is and how tracers are used in avariety of ways.

Isotope symbolTable 3 ComparisOn of Three Carbon Isotopes ___

AtomicN umber

N umberof

Neutrons

MassNumber

Radioactive?

Isotopes

ACADEWIIC VOCABULARY

(verb) to show differenceswhen compared The studentwrote a poem that contrastedwinter and summer.

WORD OrnGal3

from Greek isos (means equal)and topos (means place)

198 chapter 4 • ~~derstanding the Atom

conceP~ In/ Interactive Table Organize information about

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Deuterium j Tritium

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I IsotopeS of HydrogenThe atomic number of hydrogen is 1. This element is in the first

block of the periodic table. All hydrogen atoms have one proton.The most common isotope of hydrogen sometimes is called protium. protium (PROH tee um) has no neutrons. Its mass numberis the same as its atomic number-_.one. Two other hydrogen isotopes are called deuterium (doo TEER ee um) and tritium (TRIHtee urn). These isotopes do have neutrons. Deuterium has one proton and one neutron. Its mass number is 2. Tritium has one protonand two neutrons. Its mass number is 3. Table 4 illustrates howthe three isotopes differ. How many electrons are in the atoms ofeach of the hydrogen isotopes?

Protium, deuterium, and tritium are the only isotopes of anyelement that have special names. They have the same chemicalproperties. However, they have different physical properties. Of thethree isotopes, tritium is the only one that is radioactive. Tritium isuseful in scientific research because its radioactivity makes it easyto detect. Scientists also use deuteriurn to study chemical reactions.

f7~!J~1~ Table 4 what’s the name of the isotope of hydrogen~ that has two neutrons?

~0~s~Gaining or Losing ElectronsBecause the number of protons and the number of electrons are

equal, an atom is neutral. The positive and negative charges of thetwo types of particles balance. However, atoms can lose or gainelectrons. An atom that has lost or gained electrons doesn’t havethe same number of electrons as protons. This means the atomis no longer neutral. It has become an ion. An ion is an atom thatis no longer neutral because it has gained or lost electrons. Ionsform substances called ionic compounds.

Table 4 lsotoPeS~0f_Hydrogen

Name Protium

zoo Chapter 4 • ~~derstanding the Atom

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+ Figure 29 The sodiumo atom has eleven protons and

S c eleven electrons. The sodium

—4 + ~ ion has eleven protons andonly ten electrons.

C C Determine what neutral atom

has ten electrons.

Sodium atom Sodium ion One electron

Positive Ions—Losing ElectronsWhen an atom loses an electron, it has more protons than elec

trons. As a result, it has a positive charge. An atom with a positivecharge is called a positive ion. A positive ion is represented by theelement’s symbol with a superscript plus sign (+). A positivelycharged hydrogen ion is written H~. If an atom loses two electrons,the symbol for the ion has the superscript 2+• For example, calciumloses two electrons and forms the positive ion Ca2~. Figure 29

shows a diagram of sodium becoming an ion. When you look atthe figure, remember that electrons do not move in circular orbits.Diagrams are drawn in this way for ease of use.

As you read, refer to the periodic table in Figure 26. Note thepositions of the elements being discussed. Elements on the left sideof the periodic table are most likely to lose electrons to form positive ions. For example, elements in Group 1, such as lithium andpotassium, easily lose one electron to form ions. These ions have+1 charge, Lit, and Kt Elements in Group 2, such as magnesiumand calcium, easily lose two electrons. These elements form ionswith +2 charges, Mg2~ and Ca2~. Some members of Group 13 canlose three electrons and form ions with +3 charges. Aluminum,for example, forms the ion Al3~.

Negative Ions—Gaining ElectronsWhen an atom gains an electron, it forms an ion with a negative

charge. A negative ion has more electrons than protons. Elementson the right side of the periodic table are most likely to form negative ions. Elements in Group 17 easily gain one electron. For example, fluorine and chlorine can form ions with a —1 charge. A singlenegative charge is shown as a superscript minus sign. The ions offluorine and chlorine are represented as F- and Cl-. Oxygen andsulfur are in Group 16. These atoms can gain two electrons toform ions with —2 charges, O2 and 52-S Positive and negative ionsattract each other because of their opposite charges. In this way,compounds are formed.

Lesson 3 • Elements, Isotopes, and Ions—How Atoms Differ 201

j

Reviewing Elements,Isotopes, and Ions

You have read that all the atoms of an element have the samenumber of protons. The number of protons in an element is itsatomic number. Elements are arranged in the periodic tableaccording to their atomic numbers. Some atoms of the same element may have different numbers of neutrons in their nuclei.These different types of atoms are called isotopes. The total number of protons and neutrons in an atom is its mass number. Someatoms can lose electrons to become positive ions. Other atoms cangain electrons to become negative ions. In the next chapter, youwill see how ions can combine to form ionic compounds.

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