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

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We can explain the absence of noble gases by looking at their chemical

reactivity—or lack of it. Life is a very complex process of chemical change involv-

ing regulated chemical reactions. The noble gases generally do not participate in

chemical change because they are so unreactive, so we would not expect to ﬁnd

them playing any part in the chemistry of life.

The fact that life is made primarily of the smaller elements, from a small sec-

tion of the periodic table, may be linked to the fact that these elements are gener-

ally more abundant on Earth, and in the universe at large, than elements with

very large atoms. As we’ll see in Chapter 21, larger atoms can be built from

smaller ones. Although there are exceptions, the larger atoms are generally less

abundant than smaller ones in nature. In any case, nearly all of the mass of a

human is made up of atoms of hydrogen (atomic number = 1), carbon (atomic

number =6), nitrogen (atomic number =7), and oxygen (atomic number = 8).

We are largely made from some of the smallest atoms of the environment.

Phosphorus (atomic number = 15) and sulfur (atomic number = 16) are also

important as part of proteins, DNA, and RNA.

Another feature of the elemental composition of human life that you may

have identiﬁed is that a relatively high number of the elements found within us

consists of transition metals, although we contain these in very small amounts.As

we mentioned earlier, the chemical versatility of transition metal ions, such as

their tendency to readily form ions of differing charges, makes them useful in

enzymes—biological catalysts of chemical reactions.

It is also notable that humans contain elements from every block of the peri-

odic table except the f-block, which comprises the lanthanides and actinides. The

elements in this block are composed of very large atoms. Many of the actinides

are unstable and therefore radioactive (see Chapter 21).

HERE’S WHAT WE KNOW SO FAR

■

The elements within a group have a common number of electrons in their

highest energy level.

■

The elements in the periodic table are arranged into three main sections: the

metals, nonmetals, and metalloids (or semimetals).

■

The structure of the periodic table includes “blocks”deﬁned in terms of which

type of orbital is being ﬁlled. This gives us the s-block, p-block, d-block, and

f-block elements.

■

The elements that predominate in living things are carbon, hydrogen, nitro-

gen, and oxygen, with small concentrations of phosphorus, sulfur, and transi-

tion metals.

278 Chapter 7 Periodic Properties of the Elements

A seven-base-pair segment of double-

stranded DNA. The phosphorus atoms

are yellow.

Elements of the

Human Body

Abundance

Element (percent, by mass)

Oxygen 61.0

Carbon 23.8

Hydrogen 10.0

Nitrogen 2.6

Calcium 1.4

Phosphorus 1.1

Sulfur 0.2

Silicon 0.02

Potassium 0.02

Sodium 0.014

Chlorine 0.012

Fluorine 0.004

Magnesium 0.003

Iron 0.0004

Strontium 0.00005

Bromine 0.00003

Lead 0.00002

Trace amounts of Mn,

I, Cr, Mo, Se, and V

TABLE 7.13TABLE 7.13

7.4 The Concept of Periodicity 279

7.4 The Concept of Periodicity

Why is the structure of the periodic table useful to know? Once we understand the

trends as we traverse the groups and periods in the periodic table, we can make

reasonable predictions about the chemical and physical behavior of any element in

the periodic table. We will add to our understanding in subsequent chapters, even

learning how to assess the likely nuclear behavior of an element. Once we are

armed with this understanding, the periodic table serves as a most wonderful

guide to the formation and interaction of substances.

The basis of the periodic table is

periodicity, which will be revealed in many of

the properties we explore in the remainder of this chapter. When chemists talk of

“periodic properties” among the elements, they mean the way in which char-

acteristic properties recur in a periodic manner as we move through the periodic

table. For example, element number 3 in the table, lithium, is a very reactive

metal that forms an alkali on reaction with water. Element 4 (beryllium), is less

reactive, and as we move through elements 5 (boron), 6 (carbon), 7 (nitrogen), 8

(oxygen), 9 (ﬂuorine), and 10 (neon), we ﬁnd elements that become steadily less

like lithium in chemical reactivity and physical properties. Then suddenly, with

element 11 (sodium), we ﬁnd another very reactive metal that forms an alkali on

reaction with water. The similarities between the reactivities of lithium and

sodium are so striking that it is clear that we are observing some signiﬁcant re-

peating feature of reactivity as we move through the periodic table. If we then

move on to elements 12 through 19, we ﬁnd the same thing happening again.

Elements 12 through 18 have properties less and less like those of sodium and

lithium, and then suddenly, with element 19 (potassium), the property of being

a very reactive metal that forms an alkali on reaction with water recurs. These

chemical characteristics, which we call the characteristics of the alkali metals,

recur in a periodic manner as we move through the periodic table.

That is the basic concept of chemical periodicity, and we could have chosen

various other properties to make the same point. For example, the property of

being a very unreactive gaseous element that exists as free individual atoms

occurs in elements with atomic numbers 2, 10, 18, 36, 54, and 86. As with the

alkali metals, we have a chemical property—the lack of reactivity, in this case—

that recurs in a systematic, or periodic, manner as we move through the periodic

table. The number of elements we have to pass by before a characteristic property

recurs is not constant. It is 8, 8, 18, 18, 32 in the series above (see Exercise 7.4), but

the crucial fact is that each characteristic property does recur periodically as we

move through the periodic table.

EXERCISE 7.4 Explaining the Periods Between Periodicities

The periodic property of being an “inert gas” recurs after we move on to 8, then 8,

then 18, then 18, then 32 intervening elements in the periodic table. Can you suggest

the underlying physical reason why the periodicity follows this particular pattern?

Solution

The pattern is a consequence of the way in which electrons ﬁll up orbitals before a

particular electron arrangement associated with a particular property recurs. The

fundamental characteristic associated with being an inert gas is to have a stable octet

of eight electrons in the atom’s highest energy level, or two electrons in a completely

full energy level in the case of helium. The Aufbau principle indicates that after

helium, the stable octet will recur after the 2s and 2p orbitals have been ﬁlled, which

means eight electrons must be added before we arrive at neon, then another eight

before we arrive at argon. As we then move through Period 4, ten 3d electrons must

be added before the 4p orbitals become ﬁlled, so this time we must move through 18

FIGURE 7.7

This three-dimensional periodic table,

known as the ElemenTree, was devel-

oped by Canadian Fernando Dufour to

emphasize the periodic relationships of

the elements.

Visualization: Periodic Table

Trends

The chelating agent known as

18-crown-6 binds to sodium cations.

elements before the stable octet recurs. Similar reasoning applies to the gap between

krypton and xenon. Then the gap jumps up to 32, between xenon and radon, as a

consequence of the ﬁlling of f orbitals that must occur before the 6p orbitals of

radon can be ﬁlled. This explanation is most clearly visualized by looking at the long

version of the periodic table in Figure 7.5 or the pyramidal version in Figure 7.7.

PRACTICE 7.4

Judging on the basis of the periodicity of the periodic table, indicate what would be

the hypothetical atomic number of a metal that would be more reactive than Fr.

How many electrons would it take to ﬁll a noble gas in the hypothetical Period 8 of

the periodic table?

See Problems 29–34.

7.5 Atomic Size

Heavy metal poisoning is unfortunately a common occurrence, particularly

among children living in homes painted with lead-based paints. In the hospital,

once the poisoning is recognized, the patient is treated with a chelating agent.

Chelating agents (see Chapter 18) speciﬁcally grab ions of a particular size and

help to remove them from the body. The size of the ion is extremely important in

how well it binds to the chelating agent. If it is too small or too large, it will not

associate with the chelating agent and be excreted from the body in that manner.

Figure 7.8 reveals the key trends and periodicities found in the size of atoms.

What do you note about the size of the atoms in the periodic table? Atomic size de-

creases as we move from left to right along any period, and it increases as we move

down any group. Figure 7.8 uses atomic radii as a measure of atomic size. The

atomic radius is deﬁned as half the distance between the nuclei in a molecule con-

sisting of identical atoms. It is also known as the

covalent radius of an atom,

because it indicates the size of the atom when that atom is involved in covalent

bonding. Not all elements form such molecules. Some atomic radii values must

be estimated via indirect methods, such as comparing the distances between

atomic nuclei when the atoms are bound within chemical compounds instead of

in a molecule made of identical atoms. Values for the radii of metal atoms can

also be obtained by analyzing the distance between the nuclei of the atoms within

the solid structure of the metal concerned. These values are called

metallic radii.

The uncertainties involved in measuring atomic radii, and the various meth-

ods that can be used, mean that the values we obtain should be regarded as ap-

proximate, and you may ﬁnd slightly different values quoted in different sources.

The basic trends, however, are absolutely clear: Atomic size decreases along periods

and increases down groups. How can we explain these data? As we move down the

group, we encounter atoms with increasing numbers of occupied electron energy

levels. Each new occupied energy level, corresponding to a higher principal quan-

tum number, includes electron orbitals of greater average radius than those of the

previous level, so the atoms grow larger as we move down a group.

It might seem more difﬁcult to explain why we ﬁnd atoms of de-

creasing size as we move along any period, because the atoms actually

contain more matter as we travel from left to right, thanks to the

steady addition of protons, electrons, and neutrons. However, the

additional electrons are being added to energy levels already present

in previous elements of the period, and they are held around nuclei

whose positive charge is steadily increasing across the period. This in-

creased positive charge draws the electrons of the occupied energy

levels closer to the nucleus as we move along a period, so the atomic

radius steadily decreases.

280 Chapter 7 Periodic Properties of the Elements

Application

Atomic radii can be measured by mea-

suring the distance between the nuclei

of atoms in a metal. Measuring the dis-

tance between nuclei in a molecule

gives the covalent radius

in molecules.

Video Lesson: Periods and

Atomic Size

7.5 Atomic Size 281

152

113

186

160

143

117

110

104

227

197

122

121

117

114

110

247

215

163

140

141

143

133

130

265

217

170

175

155

167

140

145

VIIIAVIIAVIAVAIVAIIIAIIAIA

Atomic radius increases

Atomic radius decreases

FIGURE 7.8

Atomic radii for selected atoms.

IIA

IIIB IVB VB VIB VIIB VIIIB IB IIB

IIIA IVA VA VIA VIIA

VIIIA

Period

Trend in atomic radius.

Visualization: Determining the

Atomic Radius of a Nonmetal

(Carbon)

Visualization: Determining the

Atomic Radius of a Nonmetal

(Chlorine)

Visualization: Determining the

Atomic Radius of a Metal

(Molybdenum)

IA IIA IIIA IVA VA VIA VIIA VIIIA

800

1086

1402

1005

869

941

1314

780

715

703

1176

1356

1527

2088

1681

1255

1060

947

1042

834

813

(926)

1009

1143

708

761

580

579

558

589

1311

899

2377

735

590

549

503

520

495

419

409

382

FIGURE 7.9

First ionization energies for selected

elements.

7.6 Ionization Energies

Most of the elements in the periodic table are metals. When they react with other

elements to form compounds, they generally do so by losing one or more elec-

trons to form positive ions. An input of energy is required to remove one or more

electrons from an atom, and the energy involved is known as

ionization energy.

The smaller the ionization energy, the more likely it is that ionization will occur.

All elements, not just the metals, have ionization energy values. By examining the

ionization energies of the elements, we can see why metals are particularly prone

to form positive ions, and we can see how the differing reactivities of metals

matches their differing ionization energies.

The

ﬁrst ionization energy (I

) of an element is the energy required to remove

the highest energy electron from an atom of the element in the gaseous ground

state. The ﬁrst ionization energy is typically quoted per mole of atoms being ion-

ized and can be symbolized in general terms, using X to represent any element, as

X(g) → X

(g) + e

−

For example, the ﬁrst ionization energies for sodium and chlorine are

Na(g) → Na

(g) + e

−

= 495 kJ/mol

Cl(g) → Cl

(g) + e

−

= 1255 kJ/mol

Why is the ﬁrst ionization energy of sodium so much smaller than that of chlorine?

Note the resulting ion’s electron conﬁguration. The sodium ion has a noble gas

electron conﬁguration. The chlorine ion does not. The ﬁrst ionization energies of

selected main-group elements are listed in Figure 7.9.

The

second ionization energy (I

), of an element is the energy required to

remove one electron from each singly charged +1 ion of an element in the

gaseous state:

(g) → X

(g) + e

−

Third, fourth, and all subsequent ionization energies can be similarly deﬁned, as

successive electrons are removed.

We can see the link between ionization energies and reactivities by examining

some of the alkali metals. Lithium, sodium, and potassium are all reactive metals

that form alkalis on reaction with water. These metals, when they react, form ions

of Li

,Na

, and K

, respectively. We can get an indication of how readily the for-

mation of these ions occurs by considering these ionization energies:

Li(g) → Li

(g) + e

−

= 520 kJ/mol

Na(g) → Na

(g) + e

−

= 495 kJ/mol

K(g) → K

(g) + e

−

= 419 kJ/mol

282 Chapter 7 Periodic Properties of the Elements

Video Lesson: Ionization Energy

EXERCISE 7.5 Implications of Ionization Energy Trends

The reactions of lithium, sodium, and potassium metals with water to form an

alkaline solution are as follows:

2Li(s) + 2H

O(l) → 2Li

(aq) + 2OH

−

(aq) + H

(g)

2Na(s) + 2H

O(l) → 2Na

(aq) + 2OH

−

(aq) + H

(g)

2K(s) + 2H

O(l) → 2K

(aq) + 2OH

−

(aq) + H

(g)

Given the values listed in the text for the ﬁrst ionization energies of these metals,

which metal will react most vigorously if added to water? If we were to put a chunk

of cesium metal in water, how might it react?

First Thoughts

Ionization energy is a measure of the energy required to remove the highest energy

electron from an atom. It is an endothermic process, so when comparing elements

in an otherwise equivalent reaction, we ﬁnd that the lower the ionization energy, the

less energy will be required, and the more readily the reaction will occur.

Solution

The potassium metal has the lowest ionization energy of the three, so it will react

most readily with water; sodium will react less readily and lithium the least impres-

sively. Figure 7.10 shows the reactions of these metals with water. Note the ﬂame

that potassium makes upon reaction with water. What about cesium? Its ionization

energy of 376 kJ/mol means that it will react even more vigorously than potassium.

In fact, the reaction with water is explosive.

7.6 Ionization Energies 283

The metals lithium (left), sodium

(center), and potassium (right).

FIGURE 7.10

The relative reactivities of metals can

sometimes be quite clearly visualized.

Sodium

Potassium

Calcium

Gold

Further Insights

Francium would be expected to react even more vigorously than cesium. However,

francium is exceedingly rare and has no stable isotopes. One place that prepares

samples for scientiﬁc study is the Nuclear Structure Laboratory at the State Univer-

sity of New York at Stony Brook. An isotope known as francium-210 (along with

ﬁve neutrons) is prepared by slamming oxygen-18 nuclei atoms into gold-197

nuclei using a nuclear accelerator:

197

Au +

O →

210

Fr + 5 n

This, and other isotopes of francium, can be produced at the rate of about 1 million

atoms per second. Keep in mind, though, that 1 million atoms are not very many

compared to the 6.02 × 10

atoms that make up a mole. The francium that is

collected can be studied to determine its physical and chemical properties, but the

nuclei decay so quickly that it is impractical to collect francium-210 for other uses.

PRACTICE 7.5

Which of the metals in Group IIA would you expect to react most vigorously with

water? Explain your answer in terms of the ionization energies of the elements.

See Problems 45, 46, 73, and 75.

We have just identiﬁed one of the general trends linking ionization energy

values with an element’s position in the periodic table. Ionization energies gener-

ally decrease moving down any group. We can explain this trend by noting that as

we move down a group, the valence electrons being removed are coming from

energy levels of successively higher principal quantum numbers. Therefore, as we

move down a group, the electrons being removed are screened from the nucleus

by additional inner energy levels occupied with electrons.

The other main trend in ionization energies is that they generally increase from

left to right along any period. We can rationalize this as due to the extra energy

needed to release an electron from the pull of the increasing nuclear charge found

as we move along a period.

These trends in ionization energy can be clearly seen in the laboratory if

we choose metals with signiﬁcantly different reactivities (see Figure 7.10). The

trends can also be plotted graphically, as shown in Figure 7.11, in a way that

makes the periodicity in this property apparent. As we move through Figure 7.11,

we see the ionization energies rising and falling in a periodic manner.

284 Chapter 7 Periodic Properties of the Elements

18 36 54 86

500

1000

1500

2000

2500

Period

Ionization energy (kJ/mol)

Atomic number

FIGURE 7.11

The values of ﬁrst ionization energies

for the elements in the ﬁrst six periods.

7.6 Ionization Energies 285

The trend in ﬁrst ionization energy. The size of the sphere

for each atom indicates the relative magnitude of the ﬁrst

ionization energy.

EXERCISE 7.6 Differences in Ionization Energies

Here are the ﬁrst three ionization energies of sodium and magnesium, in kilojoules

per mole:

Sodium 495 4560 6920

Magnesium 735 1445 7730

There is a very large difference between the ﬁrst and second ionization energies of

sodium. With magnesium, however, the very large difference is seen between the

second and third ionization energies. Why do the large differences in ionization

energies occur where they do?

First Thoughts

What does the ionization energy tell us? It is the energy required to remove an

electron from the atom. The key question, then, is “What is it about the electron

conﬁguration of each element that results in such a sudden increase in ionization

energy?”

Solution

Sodium atoms have one outer-shell electron ([Ne]3s

), whereas magnesium atoms

have two ([Ne]3s

). Once sodium has lost one electron, the next electron must come

from a deeper electron shell, closer to the nucleus, so a great deal more energy is re-

quired to remove that second electron. In the case of magnesium, the transition to

removing electrons from an inner shell does not occur until the two outer electrons

have been removed, so the big jump in ionization energy is seen between the second

and third ionization energies, rather than between the ﬁrst and second.

Further Insights

Iron is an example of a transition element that has several possible oxidation states,

including Fe

and Fe

. A most wonderful and trustworthy website to retrieve ion-

ization energy data (and so much more!) is the WebElements site at http://www

.webelements.com. Look up the ionization energy for the ions of iron at that site.

Where are the sudden increases in ionization energy, and how do they compare to

those found in a nonmetal such as phosphorus? Can you explain the differences be-

tween these two elements?

PRACTICE 7.6

Predict, and draw a graph of, the ionization energy versus electron ionized for

sulfur. Compare your graph to the measured values at the WebElements site.

See Problems 47, 48, 53, 54, 96, and 99.

7.7 Electron Afﬁnity

Our look at ionization energies concerned the energy required to generate

positive ions (cations) from neutral atoms. The other side of the “ionization coin”

is the generation of negative ions (anions) from neutral atoms. A negatively

charged ion forms when an atom gains electrons, and the energy changes associ-

ated with electron gain are called electron afﬁnities. The

electron afﬁnity is the

energy change associated with the addition of an electron to an atom in the

gaseous state, and as usual, we quote the values in kilojoules per mole of atoms.

For example,

F(g) + e

−

→ F

−

(g) −328 kJ/mol

Cl(g) + e

−

→ Cl

−

(g) −349 kJ/mol

Br(g) + e

−

→ Br

−

(g) −325 kJ/mol

I(g) + e

−

→ I

−

(g) −295 kJ/mol

At(g) + e

−

→ At

−

(g) −270 kJ/mol

The electron afﬁnity values for Group VIIA reveal one of the general trends in

electron afﬁnity and also one of the complications. The values indicate to us that

electron afﬁnities generally become less negative (or more positive), correspond-

ing to less energy being released, as we move down a group.

Does this trend make

sense?

Because added electrons are farther from the nuclei of larger atoms, they

interact less with the positive charge of the nucleus. The trend makes sense. There

are many exceptions, however, such as more energy released when a chloride ion

forms, in the above list, than upon formation of the ﬂuoride anion.Why does ﬂu-

orine not ﬁt the trend? In searching for a possible explanation, we should note

that ﬂuorine is the smallest of this group of atoms, so the electrons in

the outer p orbitals, to which the extra electron is being added, will be

much closer together than in the other atoms of this group. The larger

electron–electron repulsions between these closely spaced electrons,

compared to those of other halogen atoms, will be associated with

increased potential energy.

Figure 7.12 shows many more electron afﬁnity values. It allows us

to conﬁrm the general trend just identiﬁed within groups, in addition

to revealing exceptions. Figure 7.12 also conﬁrms that elements that

form stable negative ions have large negative electron afﬁnity values,

whereas those that form stable positive ions have much lower negative

electron afﬁnity values, or even positive values.

Note, as well, that the values in Figure 7.12 indicate that electron

afﬁnity values generally get more negative as we move from the left to

the right along a period in the periodic table.

Does this trend make

sense?

We know from our understanding of the trend associated with

atomic radii that atoms get smaller as we move to the right along the

period table. For the same reasons we discovered as we move down a

group, the electron afﬁnity generally tends to get more negative as we

move to the right along a period.

286 Chapter 7 Periodic Properties of the Elements

VIIA

IIIA

IIA

IVA

VIA

VIIIA

–60

–27

–328

–53

–48

–47

–46

–349

–325

–295

–270

–154

–141

–200

–195

–190

–183

FIGURE 7.12

Electron afﬁnity values, in kilojoules per mole

(kJ/mol), for selected elements. Electron afﬁnity

values generally decrease as we go up and to the

right in the periodic table. More negative values

indicate a greater afﬁnity for electrons.

Video Lesson: Electron Afﬁnity

7.8 Electronegativity 287

EXERCISE 7.7 Trends

Many of the periodic properties identiﬁed in the periodic table are related to each

other. Is there a relationship between electron afﬁnity and the number of valence

electrons?

First Thoughts

We should ﬁrst construct a crude periodic table with arrows pointing to the most

negative electron afﬁnity. Superimposing the number of valence electrons on this

drawing will reveal any relationship.

Solution

As the number of valence electrons gets larger, the atomic radius gets smaller across

a period. As the radius gets smaller, the electron afﬁnity becomes more negative.

Therefore, as the number of valence electrons increases, the electron afﬁnity be-

comes more negative. This trend, however, is valid only within a given period in the

periodic table, because the electron afﬁnity changes and the number of valence elec-

trons remain constant as we move down a group.

Further Insights

Periodic trends are abundant in the periodic table. Size, electron afﬁnity, reactivity,

and ionization potential are all related to position on the periodic table. This table

has trends in many other properties, such as the formation of acids or bases upon

reaction of elemental oxides with water, the tendency to form chlorides, the ten-

dency to behave like metals, and many more.

PRACTICE 7.7

Is there a relationship between the trends observed in ionization energy and elec-

tron afﬁnity? Explain.

See Problems 57–62, 74, 76, 79, and 80.

7.8 Electronegativity

In our discussion about electron afﬁnities, we are actually considering a rather

contrived situation, because most elements rarely exist in the form of free

gaseous atoms. A more meaningful characteristic of elements, one that is related

to their interaction with additional electrons and therefore is more ﬁrmly rooted

in real chemical behavior, is their electronegativity. The

electronegativity of an

atom is a measure of the ability of the atom in a molecule to attract shared elec-

trons to itself. Atoms in molecules share electrons to create covalent bonds. How-

ever, if the atoms sharing electrons are different, they will contain nuclei with

different charges and different electron arrangements. The result is that the elec-

trons are not shared equally. Electronegativity values have been calculated for

every atom in the periodic table. However, we generally do not quote electroneg-

ativity values for the noble gases (Group VIIIA) because, as a consequence of

their exceptional stability, they do not readily form bonds.

Electronegativity values for the elements in the periodic table are one of the most

powerful quantities chemists have that explain and predict the behavior of molecules

and ions. Their importance cannot be overstated. In 1932, Linus Pauling deﬁned

the concept in terms of bond energies. The Pauling scale of electronegativities

and the variations of it are the bases of much of our discussion about bonding

in subsequent chapters. To illustrate the utility of electronegativity values, let’s

consider two essential vitamins that we must regularly consume: vitamin A and

vitamin C. Deﬁciencies in these vitamins are responsible for serious maladies.

Application

Video Lesson: An Introduction

to Electronegativity