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7 Nuclear Chemistry: The Change of One Element into Another

7 Nuclear Chemistry: The Change of One Element into Another

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2.8 RADIOACTIVITY



• A nuclear reaction involves a change in an atom’s nucleus, usually producing a

different element. A chemical reaction, by contrast, involves only a change in the

way that different atoms are combined. A chemical reaction never changes the

nuclei themselves or produces a different element.

• Different isotopes of an element have essentially the same behavior in chemical

reactions but often have completely different behavior in nuclear reactions.

• The energy change accompanying a nuclear reaction is far greater than that

accompanying a chemical reaction. The nuclear transformation of 1.0 g of

uranium-235 (235

92U) releases more than one million times as much energy as the

chemical combustion of 1.0 g of methane.



2.8 RADIOACTIVITY

Scientists have known since 1896 that many nuclei are radioactive—they undergo a

spontaneous decay and emit some form of radiation. Early studies of radioactive isotopes, or radioisotopes, by Ernest Rutherford in 1897 showed that there are three

common types of radiation with markedly different properties: alpha (a), beta ( b ),

and gamma (g) radiation, named after the first three letters of the Greek alphabet.



Alpha (a) Radiation

Using the simple experiment shown in Figure 2.7, Rutherford found that a radiation

consists of a stream of particles that are repelled by a positively charged electrode,

attracted by a negatively charged electrode, and have a mass-to-charge ratio identifying them as helium nuclei, 42He2+. Alpha particles thus consist of two protons and

two neutrons.

Figure 2.7

Beta radiation is strongly deflected

toward the positive electrode.

β rays

+

Gamma radiation

is undeflected.

γ rays



The radioactive source in the shielded

box emits radiation, which passes

between two electrodes.



α rays



Alpha radiation is deflected

toward the negative electrode.



Because the emission of an ␣ particle from a nucleus results in a loss of two protons and two neutrons, it reduces the mass number of the nucleus by 4 and reduces

the atomic number by 2. Alpha emission is particularly common for heavy radioactive isotopes. Uranium-238, for example, spontaneously emits an ␣ particle and

forms thorium-234.

2 protons

2 neutrons

4 nucleons



90 protons

144 neutrons

234 nucleons



92 protons

146 neutrons

238 nucleons



238

92 U



4

234

2 He + 90 Th



The effect of an electric field on a, b,

and g radiation.



49



50



Chapter 2 ATOMS, MOLECULES, AND IONS



Note how the nuclear equation for the radioactive decay of uranium-238 is written. The equation is said to be balanced because the total number of neutrons and

protons, collectively called nucleons, or nuclear particles, is the same on both sides

of the equation and the number of charges on the nuclei and on any elementary

particles (protons and electrons) is the same on both sides. In the decay of 238

92U to

give 42He and 234

Th,

for

instance,

there

are

238

nucleons

and

92

nuclear

charges

on

90

both sides of the equation.



Beta (b) Radiation

Further work by Rutherford in the late 1800s showed that b radiation consists of a

stream of particles that are attracted to a positive electrode (Figure 2.7), repelled by a

negative electrode, and have a mass-to-charge ratio identifying them as electrons, -10e

or b-. Beta emission occurs when a neutron in the nucleus spontaneously decays into

a proton plus an electron, which is then ejected. The product nucleus has the same

mass number as the starting nucleus because a neutron has turned into a proton, but

it has a higher atomic number because it has the newly created proton. The reaction

of 131I to give 131Xe is an example:

54 protons

77 neutrons

131 nucleons



0 nucleons

but −1 charge



53 protons

78 neutrons

131 nucleons



131

53 I



131

0

54 Xe + −1 e



Writing the emitted b particle as -10e in the nuclear equation makes clear the

charge balance of the nuclear reaction. The subscript in the 131

53I nucleus on the left

(53) is balanced by the sum of the two subscripts on the right (54 - 1 = 53).



Gamma (g) Radiation

Gamma (g) radiation is unaffected by either electric or magnetic fields (Figure 2.7)

and has no mass. Like visible light, ultraviolet rays, and X rays, g radiation is simply

electromagnetic radiation of very high energy, which we’ll discuss in more detail in

Section 5.1. Gamma radiation almost always accompanies a and b emission as a

mechanism for the release of energy, but it is often not shown when writing nuclear

equations because it changes neither the mass number nor the atomic number of the

product nucleus.



Positron Emission and Electron Capture

In addition to a, b, and g radiation, two other types of radioactive decay processes

also occur commonly: positron emission and electron capture. Positron emission occurs

when a proton in the nucleus changes into a neutron plus an ejected positron (+10e or

b+), a particle that can be thought of as a positive electron. A positron has the

same mass as an electron but an opposite charge.

The result of positron emission is a decrease in the atomic number of the product

nucleus but no change in the mass number. Potassium-40, for example, undergoes

positron emission to yield argon-40, a nuclear reaction important in geology for dating

rocks. Note once again that the sum of the two subscripts on the right of the nuclear

equation (18 + 1 = 19) is equal to the subscript in the 40

19K nucleus on the left.

18 protons

22 neutrons

40 nucleons



0 nucleons

but +1 charge



19 protons

21 neutrons

40 nucleons



40

19 K



40

0

18 Ar + 1 e



2.8 RADIOACTIVITY



51



You might already know that the acronym PET used in medical imaging stands

for positron emission tomography. A chemical compound containing a positronemitting isotope, usually 18F, is injected into the body and accumulates at a certain

site, such as in a tumor. When decay occurs, the emitted positron reacts with a nearby

electron and is instantly annihilated, releasing gamma rays whose position in the

body can be detected.

Electron capture is a process in which the nucleus captures one of the surrounding

electrons in an atom, thereby converting a proton into a neutron. The mass number of

the product nucleus is unchanged, but the atomic number decreases by 1, just as in

positron emission. The conversion of mercury-197 into gold-197 is an example:

79 protons

118 neutrons

197 nucleons



Inner-shell

electron

80 protons

117 neutrons

197 nucleons



197

80 Hg



ϩ



0

–1 e



᭡ A PET scan of a 62 year old man with a

brain tumor, as indicated by the yellow

and orange area in the lower left portion

of the brain.



197

79 Au



Characteristics of the different kinds of radioactive decay processes are summarized in Table 2.2.



TABLE 2.2



A Summary of Radioactive Decay Processes



Process

Alpha emission



Symbol

4

2He



or a

-



0

-1e



or b



Gamma emission



0

0g



or g



Positron emission



0

1e



Electron capture



E. C.



Beta emission



+



or b



Change in

Atomic Number



Change in

Mass Number



Change in

Neutron Number



-2



-4



-2



+1



0



-1



0



0



0



-1



0



+1



-1



0



+1



WORKED EXAMPLE 2.7



BALANCING NUCLEAR EQUATIONS

Write a balanced nuclear equation for each of the following processes:

4

(a) Alpha emission from curium-242: 242

96Cm : 2He + ?

0

(b) Beta emission from magnesium-28: 28

12Mg : -1e + ?

118

0

(c) Positron emission from xenon-118: 54Xe : 1e + ?



STRATEGY



The key to writing nuclear equations is to make sure that the number of nucleons is the

same on both sides of the equation and that the number of nuclear charges plus electron or positron charges is the same.

SOLUTION



(a) In ␣ emission, the mass number decreases by 4 and the atomic number decreases by

4

238

2, giving plutonium-238: 242

96Cm : 2He + 94Pu

(b) In ␤ emission, the mass number is unchanged and the atomic number increases by

0

28

1, giving aluminum-28: 28

12Mg : -1e + 13Al

(c) In positron emission, the mass number is unchanged and the atomic number

0

118

decreases by 1, giving iodine-118: 118

54Xe : 1e + 53I



52



Chapter 2 ATOMS, MOLECULES, AND IONS



Ī PROBLEM 2.11



Write a balanced nuclear equation for each of the following processes:

0

(a) Beta emission from ruthenium-106: 106

44Ru : -1e + ?

4

(b) Alpha emission from bismuth-189: 189

83Bi : 2He + ?

204

(c) Electron capture by polonium-204: 84Po + -10e : ?



Ī PROBLEM 2.12



What particle is produced by decay of thorium-214 to radium-210?

214

90Th



:



210

88Ra



+ ?



CONCEPTUAL PROBLEM 2.13 Identify the isotopes involved, and tell what type of

decay process is occurring in the following nuclear reaction:



Neutrons



81



80



79

67



68



69



Atomic number



2.9 NUCLEAR STABILITY

Why do some nuclei undergo spontaneous radioactive decay while others do not?

Why, for instance, does a carbon-14 nucleus, with 6 protons and 8 neutrons, spontaneously emit a b particle, whereas a carbon-13 nucleus, with 6 protons and 7

neutrons, is nonradioactive?

The answer has to do with the neutron/proton ratio in the nucleus and with the

forces holding the nucleus together. To see the effect of the neutron/proton ratio on

nuclear stability, look at the grid in Figure 2.8. Numbers on the side of the grid give

the number of neutrons in different nuclei, and numbers along the bottom give the



200

180



Combinations outside the

band are not stable enough

to be detected.



160

Number of neutrons (N)



The “island of

stability” near

114 protons and

184 neutrons

corresponds to

a group of

superheavy

nuclei that are

predicted to be

radioactive but

stable enough to

be detected. The

first member of

this group was

reported in 1999.



140



Band of stability



120

100

80

60

1:1 neutron/proton ratio



Figure 2.8



The band of nuclear stability. The band

indicates various neutron/proton

combinations that give rise to nuclei that

are either nonradioactive or that are

radioactive but decay slowly enough to

exist for a measurable time.



40

20

0



0



10



20



30



40



50



60



70



80



Number of protons (Z)



90



100 110 120



2.9 NUCLEAR STABILITY



number of protons. The first 92 elements are naturally occurring, while the remainder are the artificially produced transuranium elements. (Actually, only 90 of the

first 92 elements occur naturally. Technetium and promethium do not occur naturally

because all their isotopes are radioactive and have very short lifetimes. Francium and

astatine occur on Earth only in very tiny amounts.)

When the more than 3600 known isotopes are plotted on the neutron/proton grid

in Figure 2.8, they fall in a curved band sometimes called the band of nuclear stability.

Even within the band, only 264 of the isotopes are nonradioactive. The others decay

spontaneously, although their rates of decay vary enormously. On either side of the

band is a so-called sea of instability representing the large number of unstable

neutron–proton combinations that have never been detected. Particularly interesting

is the island of stability predicted to exist for a few superheavy isotopes near

114 protons and 184 neutrons. The first members of this group—287114, 288114, and

289

114—were prepared in 1999 and do indeed seem to be stable enough to live for

several seconds before they decay.

Several generalizations can be made about nuclear stability:

• Every element in the periodic table has at least one radioactive isotope.

• Hydrogen is the only element whose most abundant isotope (11H) contains more

protons (1) than neutrons (0).

• The ratio of neutrons to protons gradually increases, giving a curved appearance

to the band of stability.

• All isotopes heavier than bismuth-209 are radioactive, even though they may

decay slowly and be stable enough to occur naturally.

A close-up look at a segment of the band of nuclear stability (Figure 2.9) shows the

interesting trend that radioactive nuclei with higher neutron/proton ratios (top side



130



125

Nuclei with higher

neutron/proton ratios

tend to undergo β

emission.



120



Number of neutrons (N)



115

Beta emission

110

Nonradioactive

105

Positron emission or

electron capture

100

Alpha emission

95



Nuclei with lower

neutron/proton ratios

tend to undergo

positron emission,

electron capture,

or α emission.



90



Figure 2.9



85



80



65



70



75



Number of protons (Z)



80



A close-up look at the band of nuclear

stability. This look at the region from

Z = 66 (dysprosium) through Z = 79

(gold) shows the types of radioactive

processes that various radioisotopes

undergo.



53



54



Chapter 2 ATOMS, MOLECULES, AND IONS



of the band) tend to emit b particles while nuclei with lower neutron/proton ratios

(bottom side of the band) tend to undergo nuclear decay by positron emission, electron capture, or ␣ emission.

The trend shown in Figure 2.9 makes sense if you think about it: The nuclei on the

top side of the band are neutron-rich and therefore undergo a process—b emission—

that decreases the neutron/proton ratio by converting a neutron into a proton. The

nuclei on the bottom side of the band, by contrast, are neutron-poor and therefore

undergo processes that increase the neutron/proton ratio. Take a minute to convince

yourself that ␣ emission does, in fact, increase the neutron/proton ratio for heavy

nuclei in which n>p.

This process decreases

the neutron>proton ratio:



e Beta emission:



Neutron : Proton + b -



These processes increase

the neutron>proton ratio:



Positron emission:

Electron capture:

L Alpha emission:



Proton : Neutron + b +

Proton + Electron : Neutron

A

A-4

4

ZX : Z - 2 Y + 2He



Ī PROBLEM 2.14



(a) Of the two isotopes 173Au and 199Au, one decays by b emission and one decays by

␣ emission. Which does which?

(b) Of the two isotopes 196Pb and 206Pb, one is nonradioactive and one decays by

positron emission. Which is which?



2.10 MIXTURES AND CHEMICAL COMPOUNDS;

MOLECULES AND COVALENT BONDS

Although only 90 elements occur naturally, there are far more than 90 different kinds

of matter on Earth. Just look around, and you’ll surely find a few hundred. All the

many kinds of matter you see can be classified as either mixtures or pure substances

(Figure 2.10). Pure substances, in turn, can be either elements or chemical compounds.



Matter



Pure substances



Mixtures



Elements



Chemical compounds



Figure 2.10



A scheme for the classification of matter.



A mixture is simply a blend of two or more substances added together in some

arbitrary proportion without chemically changing the individual substances themselves. Thus, the constituent units in the mixture are not all the same, and the

proportion of the units is variable. Hydrogen gas and oxygen gas, for instance, can be



2.10 MIXTURES AND CHEMICAL COMPOUNDS; MOLECULES AND COVALENT BONDS



mixed in any ratio without changing them (as long as there is no flame nearby to initiate reaction), just as a spoonful of sugar and a spoonful of salt can be mixed.

A chemical compound, in contrast to a mixture, is a pure substance that is formed

when atoms of different elements combine in a specific way to create a new material

with properties completely unlike those of its constituent elements. A chemical compound has a constant composition throughout, and its constituent units are all

identical. For example, when atoms of sodium (a soft, silvery metal) combine with

atoms of chlorine (a toxic, yellow-green gas), the familiar white solid called sodium

chloride (table salt) is formed. Similarly, when two atoms of hydrogen combine with

one atom of oxygen, water is formed.

To see how a chemical compound is formed, imagine what must happen when

two atoms approach each other at the beginning of a chemical reaction. Because the

electrons of an atom occupy a much greater volume than the nucleus, it’s the electrons that actually make the contact when atoms collide. Thus, it’s the electrons that

form the connections, or chemical bonds, that join atoms together in compounds.

Chemical bonds between atoms are usually classified as either covalent or ionic. As a

general rule, covalent bonds occur primarily between nonmetal atoms, while ionic

bonds occur primarily between metal and nonmetal atoms. Let’s look briefly at both

kinds, beginning with covalent bonds.

A covalent bond, the most common kind of chemical bond, results when two

atoms share several (usually two) electrons. A simple way to think about a covalent

bond is to imagine it as a tug-of-war. If two people pull on the same rope, they are

effectively joined together. Neither person can escape from the other as long as both

hold on. Similarly with atoms: when two atoms both hold on to some shared electrons, the atoms are bonded together (Figure 2.11).



᭡ The crystalline quartz sand on this

beach is a pure compound (SiO2), but the

seawater is a liquid mixture of many

compounds dissolved in water.



+



The two teams are joined together because both are tugging on the same rope.



Figure 2.11



A covalent bond between atoms is analogous to a tug-of-war.



The unit of matter that results when two or more atoms are joined by covalent

bonds is called a molecule. A hydrogen chloride molecule (HCl) results when a

hydrogen atom and a chlorine atom share two electrons. A water molecule (H2O)

results when each of two hydrogen atoms shares two electrons with a single oxygen

atom. An ammonia molecule (NH3) results when each of three hydrogen atoms

shares two electrons with a nitrogen atom, and so on. To visualize these and other

molecules, it helps to imagine the individual atoms as spheres joined together to

form molecules with specific three-dimensional shapes, as shown in Figure 2.12. Balland-stick models specifically indicate the covalent bonds between atoms, while

space-filling models accurately portray overall molecular shape but don’t explicitly

show covalent bonds.



55



+



Similarly, two atoms are

joined together when both

nuclei (+) tug on the same

electrons (dots).



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