1 Atomic Structure—Protons, Electrons, and Neutrons

2.1  Atomic Structure—Protons, Electrons, and Neutrons







chapter outline



chapter goals



2.1



Atomic Structure–Protons,

Electrons, and Neutrons



See Chapter Goals Revisited (page 96) for

Study Questions keyed to these goals.



2.2



Atomic Number and Atomic Mass



•



2.3



Isotopes



Describe atomic structure and define

atomic number and mass number.



2.4



Atomic Weight



•



2.5



The Periodic Table 



2.6



Molecules, Compounds, and

Formulas



Understand the nature of isotopes and

calculate atomic masses from isotopic

masses and abundances.



•



2.7



Ionic Compounds: Formulas,

Names, and Properties 



Know the terminology of the periodic

table.



•



2.8



Molecular Compounds: Formulas

and Names



Interpret, predict, and write formulas

for ionic and molecular compounds.

Name ionic and molecular compounds.



2.9



Atoms, Molecules, and the

Mole 



•

•

•



Explain the concept of the mole and

use molar mass in calculations.



•



Derive compound formulas from

experimental data.



2.10 Describing Compound Formulas

2.11



Hydrated Compounds



51



Understand some properties of ionic

compounds.



T



he chemical elements are forged in stars and, from these elements, molecules such as water and ammonia are made in outer space. These simple

molecules and much more complex ones such as DNA and hemoglobin are

found on earth. To comprehend the burgeoning fields of molecular biology as well

as all modern chemistry, we have to understand the nature of the chemical elements

and the properties and structures of molecules. This chapter begins our exploration

of the chemistry of the elements, the building blocks of chemistry, and the compounds they form.



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Nucleus (protons and neutrons)



2.1 Atomic Structure—Protons, Electrons,

and Neutrons

Around 1900 a series of experiments done by scientists in England such as Sir

Joseph John Thomson (1856–1940) and Ernest Rutherford (1871–1937) established a model of the atom that is still the basis of modern atomic theory. Atoms

themselves are made of subatomic particles, three of which are important in

chemistry: electrically positive protons, electrically negative electrons, and, in all

except one type of hydrogen atom, electrically neutral neutrons. The model

places the more massive protons and neutrons in a very small nucleus (Figure

2.1), which contains all the positive charge and almost all the mass of an atom.

Electrons, with a much smaller mass than protons or neutrons, surround the nucleus and occupy most of the volume. In a neutral atom, the number of electrons

equals the number of protons.

The chemical properties of elements and molecules depend largely on the electrons in atoms. We shall look more carefully at their arrangement and how they influence atomic properties in Chapters 6 and 7. In this chapter, however, we first

want to describe how the composition of the atom relates to its mass and then to the

mass of compounds. This is crucial information when we consider the quantitative

aspects of chemical reactions in later chapters.



kotz_48288_02_0050-0109.indd 51



Electron cloud



FIGURE 2.1   The structure of

the atom. All atoms contain a nucleus

with one or more protons (positive

electric charge) and, except for one

type of H atom, neutrons (no charge).

Electrons (negative electric charge) are

found in space as a “cloud” around the

nucleus. In an electrically neutral atom,

the number of electrons equals the

number of protons. Note that this

figure is not drawn to scale. If the

nucleus were really the size depicted

here, the electron cloud would extend

over 200 m. The atom is mostly

empty space!

51



11/22/10 9:14 AM



52



c h a p t er 2   Atoms, Molecules, and Ions



2.2 Atomic Number and Atomic Mass

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Atomic Number

All atoms of a given element have the same number of protons in the nucleus. Hydrogen is

the simplest element, with one nuclear proton. All helium atoms have two protons,

all lithium atoms have three protons, and all beryllium atoms have four protons.

The number of protons in the nucleus of an element is given by its atomic number,

which is generally indicated by the symbol Z.

Currently known elements are listed in the periodic table inside the front cover

of this book and on the list inside the back cover. The integer number at the top of

the box for each element in the periodic table is its atomic number. A sodium atom

(Na), for example, has an atomic number of 11, so its nucleus contains 11 protons.

A uranium atom (U) has 92 nuclear protons and Z = 92.



Relative Atomic Mass and the Atomic Mass Unit

Copper

29



Cu



Atomic number

Symbol



• Historical Perspective on the

Development of Our Understanding

of Atomic Structure  A brief history

of important experiments and the

scientists involved in developing

the modern view of the atom is on

pages 334–343.



With the quantitative work of the great French chemist Antoine Laurent Lavoisier

(1743–1794), chemistry began to change from medieval alchemy to a modern field

of study (page 112). As 18th- and 19th-century chemists tried to understand how the

elements combined, they carried out increasingly quantitative studies aimed at

learning, for example, how much of one element would combine with another.

Based on this work, they learned that the substances they produced had a constant

composition, so they could define the relative masses of elements that would combine to produce a new substance. At the beginning of the 19th century, John Dalton

(1766–1844) suggested that the combinations of elements involve atoms and proposed a relative scale of atom masses. Apparently for simplicity, Dalton chose a mass

of 1 for hydrogen on which to base his scale.

The atomic mass scale has changed since 1800, but like the 19th-century

chemists, we still use relative masses with the standard today being carbon. A carbon atom having six protons and six neutrons in the nucleus is assigned a mass

value of exactly 12. From chemical experiments and physical measurements, we

know an oxygen atom having eight protons and eight neutrons has 1.33291 times

the mass of carbon, so it has a relative mass of 15.9949. Masses of atoms of other

elements are assigned in a similar manner.

Masses of fundamental atomic particles are often expressed in atomic mass

units (u). One atomic mass unit, 1 u, is one-twelfth of the mass of an atom of carbon with six

protons and six neutrons. Thus, such a carbon atom has a mass of exactly 12 u. The

atomic mass unit can be related to other units of mass using the conversion factor 1

atomic mass unit (u)=1.661ì1024 g.



Mass Number



How Small Is an Atom? The radius

of the typical atom is between 30 and

300 pm (3 × 10–11 m to 3 × 10–10 m). To

get a feeling for the incredible smallness of an atom, consider that

1 cm3 contains about three times as

many atoms as the Atlantic Ocean

contains teaspoons of water.



Because proton and neutron masses are so close to 1 u, while the mass of an electron is only about 1/2000 of this value (Table 2.1), the approximate mass of an atom

can be estimated if the number of neutrons and protons is known. The sum of the

number of protons and neutrons for an atom is called its mass number and is given

the symbol A.

A = mass number = number of protons + number of neutrons



For example, a sodium atom, which has 11 protons and 12 neutrons in its nucleus,

has a mass number of 23 (A = 11 p + 12 n). The most common atom of uranium

has 92 protons and 146 neutrons, and a mass number of A = 238. Using this information, we often symbolize atoms with the following notation:

Mass number

Atomic number



A

ZX



Element symbol



The subscript Z is optional because the element’s symbol tells us what the

atomic number must be. For example, the atoms described previously have



kotz_48288_02_0050-0109.indd 52



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2.2 Atomic Number and Atomic Mass







53



Table 2.1 Properties of Subatomic Particles*



Mass

Particle



Grams



Electron



9.109383 × 10



0.0005485799



1−



0

−1e



Proton



1.672622 × 10−24



1.007276



1+



1

1p



or p+



Neutron



1.674927 × 10−24



1.008665



0



1

0n



or n



Atomic Mass Units

−28



Charge



Symbol

or e−



*These values and others in the book are taken from the National Institute of Standards and Technology

website at http://physics.nist.gov/cuu/Constants/index.html



the symbols 23

11Na or

“uranium-238.”



238

92U,



EXAMPLE 2.1



Atomic Composition



or just



23



Na or



238



U. In words, we say “sodium-23” or



Problem What is the composition of an atom of phosphorus with 16 neutrons? What is its mass

number? What is the symbol for such an atom? If the atom has an actual mass of 30.9738 u, what

is its mass in grams? Finally, what is the mass of this phosphorus atom relative to the mass of a

carbon atom with a mass number of 12?

What Do You Know? You know the name of the element and the number of neutrons. You also

know the actual mass, so you can determine its mass relative to carbon-12.

Strategy The symbol for phosphorus is P. You can look up the atomic number (which equals the

number of protons) for this element on the periodic table. The mass number is the sum of the

number of protons and neutrons. The mass of the atom in grams can be obtained from the mass

in atomic mass units using the conversion factor 1 u = 1.661 × 10−24 g. The relative mass of an

atom of P compared to 12C can be determined by dividing the mass of the P atom in atomic mass

units by the mass of a 12C atom, 12.0000 u.

Solution A phosphorus atom has 15 protons and, because it is electrically neutral, also has

15 electrons. A phosphorus atom with 16 neutrons has a mass number of 31.

Mass number = number of protons + number of neutrons = 15 + 16 = 31

The atom’s complete symbol is 3115P.

Mass of one 31P atom = (30.9738 u) × (1.661 × 10−24 g/u) =  5.145 × 10−23 g

Mass of 31P relative to the mass of an atom of 12C: 30.9738/12.0000 =  2.58115

Think about Your Answer Because phosphorus has an atomic number greater than carbon’s,

you expect its relative mass to be greater than 12.

Check Your Understanding

1.



What is the mass number of an iron atom with 30 neutrons?



2.



A nickel atom with 32 neutrons has a mass of 59.930788 u. What is its mass in grams?



3.



How many protons, neutrons, and electrons are in a 64Zn atom?



rEvIEW & cHEcK FOr SEctIOn 2.2

1.



The mass of an atom of manganese is 54.9380 u. How many neutrons are contained in one

atom of this element?

(a)



2.



25



(b) 30



(c)



29



(d) 55



An atom contains 12 neutrons and has a mass number of 23. Identify the element.

(a)



kotz_48288_02_0050-0109.indd 53



C



(b) Mg



(c)



Na



(d) Cl



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54



c h a p t er 2   Atoms, Molecules, and Ions



2.3 Isotopes

Solid H2O



Solid D2O



FIGURE 2.2   Ice made from

“heavy water.”  Water containing



ordinary hydrogen (11H, protium) forms a

solid that is less dense (d = 0.917 g/cm3

at 0 °C) than liquid H2O (d = 0.997 g/cm3

at 25 °C), so it floats in the liquid.

(Water is unique in this regard. The

solid phase of virtually all other substances sinks in the liquid phase of

that substance.) Similarly, “heavy ice”

(D2O, deuterium oxide) floats in “heavy

water.” D2O-ice is denser than liquid

H2O, however, so cubes made of D2O

sink in liquid H2O.



© Cengage Learning/Charles D. Winters



Liquid H2O



In only a few instances (for example, aluminum, fluorine, and phosphorus) do all

atoms in a naturally occurring sample of a given element have the same mass. Most

elements consist of atoms having several different mass numbers. For example, there

are two kinds of boron atoms, one with a mass of about 10 (10B) and a second with a

mass of about 11 (11B). Atoms of tin can have any of 10 different masses. Atoms with

the same atomic number but different mass numbers are called isotopes.

All atoms of an element have the same number of protons—five in the case of

boron. To have different masses, isotopes must have different numbers of neutrons.

The nucleus of a 10B atom (Z = 5) contains five protons and five neutrons, whereas

the nucleus of a 11B atom contains five protons and six neutrons.

Scientists often refer to a particular isotope by giving its mass number (for example, uranium-238, 238U), but the isotopes of hydrogen are so important that they

have special names and symbols. All hydrogen atoms have one proton. When that is

the only nuclear particle, the isotope is called protium, or just “hydrogen.” The isotope of hydrogen with one neutron, 21H, is called deuterium, or “heavy hydrogen”

(symbol = D). The nucleus of radioactive hydrogen-3, 31H, or tritium (symbol = T),

contains one proton and two neutrons.

The substitution of one isotope of an element for another isotope of the same

element in a compound sometimes can have an interesting effect (Figure 2.2). This

is especially true when deuterium is substituted for hydrogen because the mass of

deuterium is double that of hydrogen.



Isotope Abundance

A sample of water from a stream or lake will consist almost entirely of H2O where the

H atoms are the 1H isotope. A few molecules, however, will have deuterium (2H) substituted for 1H. We can predict this outcome because we know that 99.985% of all hydrogen atoms on earth are 1H atoms. That is, the abundance of 1H atoms is 99.985%.





Percent abundance ϭ



number of atoms of a given isotope

ϫ 100% (2.1)

total number of atoms of all isotoopes of that element



The remainder of naturally occurring hydrogen is deuterium, whose abundance is

only 0.015% of the total hydrogen atoms. Tritium, the radioactive 3H isotope, occurs naturally in only trace amounts.

Consider again the two isotopes of boron. The boron-10 isotope has an abundance of 19.91%; the abundance of boron-11 is 80.09%. Thus, if you could count

out 10,000 boron atoms from an “average” natural sample, 1991 of them would be

boron-10 atoms and 8009 of them would be boron-11 atoms.



Determining Atomic Mass and Isotope Abundance

• Atomic Masses of Some Isotopes

Atom

Atomic Mass (u)

4

He

4.0092603

13

C

13.003355

16

O

15.994915

58

Ni

57.935346

60

Ni

59.930788

79

Br

78.918336

81

Br

80.916289

197

Au

196.966543

238

U

238.050784

• Isotopic Masses and the Mass

Defect  Actual masses of atoms are always less than the sum of the masses

of the subatomic particles composing

that atom. This is called the mass defect

and the reason for it is discussed in

Chapter 23.



kotz_48288_02_0050-0109.indd 54



The masses of isotopes and their abundances are determined experimentally using a

mass spectrometer (Figure 2.3). A gaseous sample of an element is introduced into

the evacuated chamber of the spectrometer, and the atoms or molecules of the sample are converted to positively charged particles (called ions). The cloud of ions

forms a beam as they are attracted to negatively charged plates within the instrument. As the ions stream toward the negatively charged detector, they fly through a

magnetic field, which causes the paths of the ions to be deflected. The extent of deflection depends on particle mass: The less massive ions are deflected more, and the

more massive ions are deflected less. The ions, now separated by mass, are detected

at the end of the chamber. Chemists using modern instruments can measure isotopic

masses to as many as nine significant figures.

Except for carbon-12, whose mass is defined to be exactly 12 u, isotopic masses

do not have integer values. However, the isotopic masses are always very close to

the mass numbers for the isotope. For example, the mass of an atom of boron-11

(11B, 5 protons and 6 neutrons) is 11.0093 u, and the mass of an atom of iron-58

(58Fe, 26 protons and 32 neutrons) is 57.9333 u.



11/18/10 2:05 PM



55



2.4 Atomic Weight







IONIZATION



ACCELERATIO N



DEFLECTION



Magnet



Electron gun



A mass spectrum is a plot of the

relative abundance of the charged

particles versus the ratio of

mass/charge (m/z).



Heavy ions

are deflected

too little.



e−e−e−

e−e−e−

e−e−e−



∙

Gas inlet



DETECTION



20Ne+



∙



Repeller Electron

trap

plate



To mass

analyzer



22Ne+



Accelerating

plates



21Ne+



Magnet



Light ions

are deflected

too much.

To vacuum pump



1. A sample is introduced as a vapor

into the ionization chamber.

There it is bombarded with highenergy electrons that strip electrons

from the atoms or molecules of the

sample.



2. The resulting positive particles are

accelerated by a series of negatively

charged accelerator plates into an

analyzing chamber.



Detector



3. This chamber is in a magnetic

field, which is perpendicular to the

direction of the beam of charged

particles.

The magnetic field causes the beam

to curve. The radius of curvature

depends on the mass and charge of

the particles (as well as the

accelerating voltage and strength of

the magnetic field).



Relative Abundance



VA P O RIZATION



100

80

60

40

20

0



20



21



22



m/z



4. Here, particles of 21Ne+ are focused

on the detector, whereas beams of ions

of 20Ne+ and 22Ne+ (of lighter or

heavier mass) experience greater and

lesser curvature, respectively, and so

fail to be detected.

By changing the magnetic field, charged

particles of different masses can be

focused on the detector to generate the

observed spectrum.



FIGURE2.3 Mass spectrometer. A mass spectrometer will separate ions of different mass and

charge in a gaseous sample of ions. The instrument allows the researcher to determine the accurate mass

of each ion, whether the ions are composed of individual atoms, molecules, or molecular fragments.



rEvIEW & cHEcK FOr SEctIOn 2.3

Silver has two isotopes, one with 60 neutrons (percent abundance = 51.839%) and the other

with 62 neutrons. What is the symbol of the isotope with 62 neutrons, and what is its percent

abundance?

(a)



107

47Ag,



51.839%



(b)



107

47Ag,



48.161%



(c)



109

47Ag,



51.839%



(d)



109

47Ag,



48.161%



2.4 AtomicWeight

Every sample of boron has some atoms with a mass of 10.0129 u and others with a

mass of 11.0093 u. The atomic weight of the element, the average mass of a representative sample of boron atoms, is somewhere between these values. For boron, for

example, the atomic weight is 10.81. If isotope masses and abundances are known,

the atomic weight of an element can be calculated using Equation 2.2.

 % abundance isotope 1 

Atomic weight ϭ 

 (mass of isotope 1)



100

 % abundance isotope 2 

ϩ

 (mass of isotope 2) ϩ . . .



100



(2.2)



• Atomic Mass, Relative Atomic

Mass, and Atomic Weight The atomic

mass is the mass of an atom at rest.

The relative atomic mass, also known

as the atomic weight or average atomic

weight, is the average of the atomic

masses of all of the element’s isotopes.

The term atomic weight is slowly being

phased out in favor of “relative atomic

mass.”



For boron with two isotopes (10B, 19.91% abundant; 11B, 80.09% abundant), we find

 19.91 

 80.09 

ϫ 10.0129 ϩ 

ϫ 11.0093 ϭ 10.81

Atomic weight ϭ 

 100 

 100 



Equation 2.2 gives an average mass, weighted in terms of the abundance of each

isotope for the element. As illustrated by the data in Table 2.2, the atomic weight of an

element is always closer to the mass of the most abundant isotope or isotopes.



kotz_48288_02_0050-0109.indd 55



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56



c h a p t er 2   Atoms, Molecules, and Ions

Table 2.2  Isotope Abundance and Atomic Weight

Element

Hydrogen



Boron



Neon



Magnesium



Symbol



Atomic

Weight



Mass

Number



Isotopic

Mass



Natural

Abundance (%)



H



1.00794



1



1.0078



99.985



D*



2



2.0141



0.015



T†



3



3.0161



0



10



10.0129



19.91



11



11.0093



80.09



20



19.9924



90.48



21



20.9938



0.27



22



21.9914



9.25



B



10.811



Ne



20.1797



Mg



24.3050



24



23.9850



78.99



25



24.9858



10.00



26



25.9826



11.01



*D = deuterium; †T = tritium, radioactive.



The atomic weight of each stable element is given in the periodic table inside

the front cover of this book. In the periodic table, each element’s box contains the

atomic number, the element symbol, and the atomic weight. For unstable (radioactive) elements, the atomic weight or mass number of the most stable isotope is given

in parentheses.



EXAMPLE 2.2



Calculating Atomic Weight from Isotope Abundance



© Cengage Learning/Charles D. Winters



Problem  Bromine has two naturally occurring isotopes. One has a mass of 78.918338 u and an

abundance of 50.69%. The other isotope has a mass of 80.916291 u and an abundance of 49.31%.

Calculate the atomic weight of bromine.

What Do You Know?  You know the mass and abundance of each of the two isotopes.

Strategy  The atomic weight of any element is the weighted average of the masses of the isotopes in a representative sample. To calculate the atomic weight, multiply the mass of each isotope by its percent abundance divided by 100 (Equation 2.2).

Solution

Atomic weight of bromine = (50.69/100)(78.918338) + (49.31/100)(80.916291) =  79.90 u 



Elemental bromine.  Bromine is a

deep orange-red, volatile liquid at room

temperature. It consists of Br2 molecules in which two bromine atoms are

chemically bonded together. There are

two, stable, naturally occurring isotopes

of bromine atoms: 79Br (50.69% abundance) and 81Br (49.31% abundance).



Think about Your Answer  You can quickly estimate the atomic weight from the data given.

There are two isotopes, mass numbers of 79 and 81, in approximately equal abundance. From this,

we would expect the average mass to be about 80, midway between the two mass numbers. The

calculation bears this out.

Check Your Understanding

Verify that the atomic weight of chlorine is 35.45, given the following information:



Example 2.3



35



Cl mass = 34.96885; percent abundance = 75.77%



37



Cl mass = 36.96590; percent abundance = 24.23%



Calculating Isotopic Abundances



Problem  Antimony, Sb, has two stable isotopes: 121Sb, 120.904 u, and 123Sb, 122.904 u. What are

the relative abundances of these isotopes?



kotz_48288_02_0050-0109.indd 56



11/18/10 2:05 PM



57



2.4 Atomic Weight



What Do You Know? You know the masses of the two isotopes of the element and know their

weighted average, the atomic weight, is 121.760 u (see the periodic table).

Strategy You can predict that the lighter isotope (121Sb) must be the more abundant

because the atomic weight is closer to 121 than to 123. To calculate the abundances recognize

there are two unknown but related quantities, and you can write the following expression

(where the fractional abundance of an isotope is the percent abundance of the isotope

divided by 100)

Atomic weight = 121.760

=  (fractional abundance of



121



(fractional abundance of



123



Sb)(120.904) + 



© Phil Degginger/Alamy







A sample of the metalloid antimony. The element has two stable

isotopes, 121Sb and 123Sb.



Sb)(122.904)



or

121.760 = x(120.904) + y(122.904)

where x = fractional abundance of 121Sb and y = fractional abundance of 123Sb. Because you know

that the fractional abundances of the isotopes must equal 1, x + y = 1, and you can solve the

equations simultaneously for x and y.

Solution Because y = fractional abundance of 123Sb = 1 − x, you can make a substitution

for y.

121.760 = x(120.904) + (1 − x)(122.904)

Expanding this equation, you have

121.760 = 120.904x + 122.904 − 122.904x

Finally, solving for x, you find

121.760 − 122.904 = (120.904 − 122.904)x

x = 0.5720

The fractional abundance of 121Sb is 0.5720 and its percent abundance is 57.20% . This means that

the percent abundance of 123Sb must be 42.80%.

Think about Your Answer The result confirms your initial inference that the lighter isotope is

the more abundant of the two.

Check Your Understanding

Neon has three stable isotopes, one with a small abundance. What are the abundances of the

other two isotopes?

20



Ne, mass = 19.992435; percent abundance = ?



21



Ne, mass = 20.993843; percent abundance = 0.27%



22



Ne, mass = 21.991383; percent abundance = ?



rEvIEW & cHEcK FOr SEctIOn 2.4

1.



Which is the more abundant isotope of copper, 63Cu or 65Cu?

(a)



2.



63



Cu



(b)



65



Cu



Which of the following is closest to the observed abundance of 71Ga, one of two stable gallium isotopes (69Ga and 71Ga)?

(a)



kotz_48288_02_0050-0109.indd 57



60%



(b) 40%



(c)



20%



(d) 70%



11/18/10 2:05 PM



c h a p t er2 Atoms, Molecules, and Ions



case study



Using Isotopes: Ötzi, the Iceman of the Alps



In 1991 a hiker in the Alps on

the Austrian-Italian border

found the well-preserved remains of an

approximately 46-year-old man, now nicknamed “The Iceman,” who lived about

5200 years ago (page 1). Studies using isotopes of oxygen, strontium, lead, and argon,

among others, have helped scientists paint a

detailed picture of the man and his life.

The 18O isotope of oxygen can give information on the latitude and altitude in which

a person was born and raised. Oxygen in

biominerals such as teeth and bones comes

primarily from ingested water. The important fact is that there is a variation in the

amount of 18O water (H218O) that depends

on how far inland the watershed is found

and on its altitude. As rain clouds move

inland, water based on 18O will be “rained

out” before H216O. The lakes and rivers on

the northern side of the Alps are known to

have a lower 18O content than those on

the southern side of the mountains. On the

northern side precipitation originates in the

cooler, and more distant, Atlantic Ocean.

On the southern side, the precipitation

comes from the closer and warmer

Mediterranean Sea. The 18O content of the

teeth and bones of the Iceman was found to

be relatively high and characteristic of the

watershed south of the Alps. He had clearly

been born and raised in that area.



The relative abundance of isotopes of

heavier elements also varies slightly from

place to place and in their incorporation

into different minerals. Strontium, a member of Group 2A along with calcium, is incorporated into teeth and bones. The ratio of

strontium isotopes, 87Sr/86Sr, and of lead

isotopes, 206Pb/204Pb, in the Iceman’s teeth

and bones was characteristic of soils from a

narrow region of Italy south of the Alps,

which established more clearly where he

was born and lived of his life.

The investigators also looked for food

residues in the Iceman’s intestines. Although

a few grains of cereal were found, they also

located tiny flakes of mica believed to have

broken off stones used to grind grain and

that were therefore eaten when the man ate

the grain. They analyzed these flakes using

argon isotopes, 40Ar and 39Ar, and found

their signature was like that of mica in an

area south of the Alps, thus establishing

where he lived in his later years.

The overall result of the many isotope

studies showed that the Iceman lived thousands of years ago in a small area about

10–20 kilometers west of Merano in northern Italy.

For details of the isotope studies, see

W. Müller, et al., Science, Volume 302,

October 13, 2003, pages 862–866.



© Handout/Reuters/Corbis



58



ƯtzitheIceman. A well-preserved mummy

of a man who lived in northern Italy about

5000 years ago.



Questions:

1. How many neutrons are there in atoms

of 18O? In each of the two isotopes of

lead?

2. 14C is a radioactive isotope of carbon that

occurs in trace amounts in all living

materials. How many neutrons are in a

14

C atom?

3. The ratio 87Sr/86Sr in the Iceman study

was in the range of 0.72. How does this

compare with the ratio calculated from

average abundances (87Sr = 7.00% and

86

Sr = 9.86%)?

Answers to these questions are available in

Appendix N.



2.5 ThePeriodicTable

•AboutthePeriodicTable



For more

information on the periodic table we

recommend the following:

• American Chemical Society (pubs

.acs.org/cen/80th/elements.html).

• www.ptable.com

• J. Emsley: Nature’s Building Blocks—

An A–Z Guide to the Elements, New

York, Oxford University Press, 2001.

• E. Scerri, The Periodic Table, New

York, Oxford University Press, 2007.



Module1:ThePeriodicTable covers

concepts in this section.



kotz_48288_02_0050-0109.indd 58



The periodic table of elements is one of the most useful tools in chemistry. Not only

does it contain a wealth of information, but it can also be used to organize many of

the ideas of chemistry. It is important to become familiar with its main features and

terminology.



DevelopingthePeriodicTable

Although the arrangement of elements in the periodic table is now understood on

the basis of atomic structure [▶ Chapters 6 and 7], the table was originally developed from many experimental observations of the chemical and physical properties

of elements and is the result of the ideas of a number of chemists in the 18th and

19th centuries.

In 1869, at the University of St. Petersburg in Russia, Dmitri Ivanovitch Mendeleev (1834–1907) was pondering the properties of the elements as he wrote a

textbook on chemistry. On studying the chemical and physical properties of the elements, he realized that, if the elements were arranged in order of increasing atomic

mass, elements with similar properties appeared in a regular pattern. That is, he saw

a periodicity or periodic repetition of the properties of elements. Mendeleev organized the known elements into a table by lining them up in horizontal rows in order

of increasing atomic mass (page 50). Every time he came to an element with properties similar to one already in the row, he started a new row. For example, the elements Li, Be, B, C, N, O, and F were in a row. Sodium was the next element then



11/22/10 9:14 AM



2.5 The Periodic Table







A CLOSER LOOK

by Eric R. Scerri, UCLA



The Story of the Periodic Table



John C. Kotz



Dmitri Mendeleev was probably the greatest scientist produced by

Russia. The youngest of 14 children, he was

taken by his mother on a long journey, on

foot, in order to enroll him into a university.

However, several attempts initially proved

futile because, as a Siberian, Mendeleev was

barred from attending certain institutions.

His mother did succeed in enrolling him in a

teacher training college, thus giving

Mendeleev a lasting interest in science education, which contributed to his eventual

discovery of the periodic system that essentially simplified the subject of inorganic

chemistry.



Statue of Dmitri Mendeleev and a periodic

table. This statue and mural are at the Institute

of Metrology in St. Petersburg, Russia.



After completing a doctorate, Mendeleev

headed to Germany for a postdoctoral fellowship and then returned to Russia, where

he set about writing a book aimed at summarizing all of inorganic chemistry. It was

while writing this book that he identified

the organizing principle with which he is

now invariably connected—the periodic

system of the elements.

More correctly, though, the periodic system was developed by Mendeleev, as well as

five other scientists, over a period of about

10 years, after the Italian chemist Cannizzaro

had published a consistent set of atomic

weights in 1860. It appears that Mendeleev

was unaware of the work of several of his

co-discoverers, however.

In essence, the periodic table groups

together sets of elements with similar properties into vertical columns. The underlying

idea is that if the elements are arranged in

order of increasing atomic weights, there are

approximate repetitions in their chemical

properties after certain intervals. As a result

of the existence of the periodic table, students and even professors of chemistry were

no longer obliged to learn the properties of

all the elements in a disorganized fashion.

Instead, they could concentrate on the properties of representative members of the

eight columns or groups in the early shortform periodic table, from which they could

predict properties of other group members.



Mendeleev is justly regarded as the

leading discoverer of the periodic table

since he continued to champion the finding and drew out its consequences to a far

greater extent than any of his contemporaries. First, he accommodated the 65 or so

elements that were known at the time into

a coherent scheme based on ascending

order of atomic weight while also reflecting chemical and physical similarities.

Next, he noticed gaps in his system, which

he reasoned would eventually be filled by

elements that had not yet been discovered.

In addition, by judicious interpolation

between the properties of known elements, Mendeleev predicted the nature of

a number of completely new elements.

Within a period of about 20 years, three of

these elements—subsequently called gallium, scandium, and germanium—were

isolated and found to have almost the

exact properties that Mendeleev had predicted.

What is not well known is that about half

of the elements that Mendeleev predicted

were never found. But given the dramatic

success of his early predictions, these later

lapses have largely been forgotten.



Eric Scerri, The Periodic Table: Its Story and

Its Significance, Oxford University Press,

New York, 2007.



known; because its properties closely resembled those of Li, Mendeleev started a

new row. As more and more elements were added to the table, new rows were begun, and elements with similar properties (such as Li, Na, and K) were placed in the

same vertical column.

An important feature of Mendeleev’s table—and a mark of his genius—was that

he left an empty space in a column when he believed an element was not known but

should exist and have properties similar to the elements above and below it in his

table. He deduced that these spaces would be filled by undiscovered elements. For

example, he left a space between Si (silicon) and Sn (tin) in Group 4A for an element he called eka-silicon. Based on the progression of properties in this group,

Mendeleev was able to predict the properties of the missing element. With the discovery of germanium (Ge) in 1886, Mendeleev’s prediction was confirmed.

In Mendeleev’s table the elements were ordered by increasing mass. A glance at

a modern table, however, shows that, if some elements (such as Ni and Co, Ar and

K, and Te and I) were ordered by mass and not chemical and physical properties,

they would be reversed in their order of appearance. Mendeleev recognized these

discrepancies and simply assumed the atomic masses known at that time were

inaccurate—not a bad assumption based on the analytical methods then in use. In

fact, his order is correct and what was wrong was his assumption that element properties were a function of their mass.



kotz_48288_02_0050-0109.indd 59



59



• Mendeleev and Atomic Numbers

Mendeleev developed the periodic table based on atomic masses. The concept of atomic numbers was not known

until after the development of the

structure of the atom in the early 20th

century.



11/18/10 2:05 PM



60



c h a p t er 2   Atoms, Molecules, and Ions



In 1913 H. G. J. Moseley (1887–1915), a young English scientist working with

Ernest Rutherford (1871–1937), bombarded many different metals with electrons

in a cathode-ray tube (page 340) and examined the x-rays emitted in the process.

Moseley realized the wavelength of the x-rays emitted by a given element was related

in a precise manner to the positive charge in the nucleus of the atoms of

Transition Metals

Group 2B



Group 2A



Magnesium—Mg



Titanium—Ti



Vanadium—V



Chromium—Cr



Manganese—Mn



Iron—Fe



Cobalt—Co



Nickel—Ni



Copper—Cu



Zinc—Zn



Mercury—Hg



Group 1A



1A



Lithium—Li



1



H



2



Li Be



3



Na Mg



4



K



Group 8A, Noble Gases



Main Group Metals

Transition Metals

Metalloids

Nonmetals



2A



3B



7B



8A



2B



3A



4A



5A



6A



7A



B



C



N



O



F



Al Si



P



S



Cl Ar



Ne



4B



5B



6B



Ca Sc Ti



V



Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr



8B



1B



He



Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te



5



Rb Sr



6



Cs Ba La Hf Ta



7



Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo



Y



W Re Os Ir



I



Xe



Pt Au Hg Tl Pb Bi Po At Rn

Neon—Ne



Potassium—K

Group 4A



Photos © Cengage Learning/Charles D. Winters



Group 3A



Boron—B



Carbon—C



Group 5A



Tin—Sn



Group 6A



Group 7A



Sulfur—S

Nitrogen—N2



Bromine—Br

Aluminum—Al



Silicon—Si



Lead—Pb



FIGURE 2.4   Some of the 118 known elements.



kotz_48288_02_0050-0109.indd 60



Selenium—Se



Phosphorus—P



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the element and that this provided a way to experimentally determine the atomic

number of a given element. Indeed, once atomic numbers could be determined,

chemists recognized that organizing the elements in a table by increasing atomic

number corrected the inconsistencies in Mendeleev’s table. The law of chemical

periodicity is now stated as the properties of the elements are periodic functions of atomic

number.



1

2

3

4

5

6

7

Periods



Features of the Periodic Table

A



The main organizational features of the periodic table are the following:

•



•



61



2.5 The Periodic Table







Elements are arranged so that those with similar chemical and physical properties lie in vertical columns called groups or families. The periodic table commonly used in the United States has groups numbered 1 through 8, with each

number followed by the letter A or B. The A groups are often called the main

group elements and the B groups are the transition elements.

The horizontal rows of the table are called periods, and they are numbered

beginning with 1 for the period containing only H and He. For example, sodium, Na, in Group 1A, is the first element in the third period. Mercury, Hg,

in Group 2B, is in the sixth period (or sixth row).



The periodic table can be divided into several regions according to the properties of the elements. On the table inside the front cover of this book, elements that

behave as metals are indicated in purple, nonmetals are indicated in yellow, and elements called metalloids appear in green. Elements gradually become less metallic as

one moves from left to right across a period, and the metalloids lie along the metalnonmetal boundary. Some elements are shown in Figure 2.4.

You are probably familiar with many properties of metals from your own experience (Figure 2.5a). At room temperature and normal atmospheric pressure metals

are solids (except for mercury), can conduct electricity, are usually ductile (can be

drawn into wires) and malleable (can be rolled into sheets), and can form alloys

(mixtures of one or more metals in another metal). Iron (Fe) and aluminum (Al)

are used in automobile parts because of their ductility, malleability, and low cost

relative to other metals. Copper (Cu) is used in electric wiring because it conducts

electricity better than most other metals.

The nonmetals lie to the right of a diagonal line that stretches from B to Te in

the periodic table and have a wide variety of properties. Some are solids (carbon,

sulfur, phosphorus, and iodine). Five elements are gases at room temperature (hydrogen, oxygen, nitrogen, fluorine, and chlorine). One nonmetal, bromine, is a

liquid at room temperature (Figure 2.5b). With the exception of carbon in the form



Bismuth



Copper



Bromine, Br2



Iodine, I2



3 4 5 6 7 8



B

3 4 5 6 7



8



1 2



Groups or Families



• Periods and Groups in the Periodic

Table One way to designate periodic

groups is to number them 1 through

18 from left to right. This method is

generally used outside the United

States. The system predominant in

the United States labels main group

elements as Groups 1A–8A and transition elements as Groups 1B–8B.

This book uses the A/B system.



Main Group Metals

Transition Metals



Metalloids

Nonmetals



Forms of

silicon



Photos © Cengage Learning/Charles D. Winters



Molybdenum



A



1 2



(a)

Metals Molybdenum (Mo, wire), bismuth (Bi, center

(a) Metals



(b)

(b) Nonmetals

Nonmetals Only about 18



(c)

(c) Metalloids

Metalloids Only 6 elements are gener-



object), and copper (Cu) are metals. Metals can be generally

drawn into wires and conduct electricity.



elements can be classified as

nonmetals. Here are orange-red

liquid bromine and purple solid

iodine.



ally classified as metalloids or semimetals.

This photograph shows solid silicon in

various forms, including a wafer that holds

printed electronic circuits.



FIGURE2.5 Metals, nonmetals, and metalloids.



kotz_48288_02_0050-0109.indd 61



11/18/10 2:05 PM