Chapter 3 Atoms and - Cengage Learning4ltrpress.cengage.com/chem/demo/docs/chem_ch3_sample.pdf3.2...

W hy does an element or compound have the properties it has? Why does one element or compound undergo a change that another element or compound will not un-

dergo? Inanimate matter behaves the way it does because of the nature of its parts. The use of atoms to represent these parts dates back to about 400 B.C., when the Greek philosopher Leucippus and his student Democritus (460–370 B.C.) argued for a limit to the divisibility of mat-ter, which was counter to the prevailing view of Greek philosophers that matter is endlessly divisible. Democritus used the Greek word atomos, which literally means “uncuttable,” to describe the ultimate particles of matter, particles that could not be divided further. How-ever, it wasn’t until John Dalton (1766–1844) introduced his atomic theory in 1803 that the importance of using atoms to explain proper-ties of matter was recognized.

Dalton’s atomic theory and the development of the periodic table by Mendeleev in 1869 led to the rapid growth of chemistry as a sci-ence. In particular, the infl uence of the location and number of elec-trons in atoms on the properties of elements has become one of the essential ideas of chemistry. In this chapter, we will use current knowl-edge about the atom together with the periodic table as the basis for our understanding of the chemical view of matter.

3.1 John Dalton’s Atomic Theory

John Dalton, drawing from his own quantitative experiments and those of earlier scientists, proposed the following in 1803:

1. All matter is made up of indivisible and indestructible particles called atoms.

2. All atoms of a given element are identical, both in mass and in properties. By contrast, atoms of different elements have different masses and different properties.

chemChapter 3

Atoms and the Periodic Table

sections

3.1 John Dalton’s Atomic Theory

3.2 Structure of the Atom

3.3 Modern View of the Atom

3.4 Where Are the Electrons in Atoms?

3.5 Development of the Periodic Table

3.6 The Modern Periodic Table

3.7 Periodic Trends

3.8 Properties of Main-Group Elements

30 C h a p t e r 3 : A t o m s a n d t h e P e r i o d i c T a b l e© 2011 Cengage Learning. All Rights Reserved. May not be scanned, copied or duplicated, or posted to a publicly accessible website, in whole or in part.

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3. Compounds form when atoms of different elements combine in ratios of small whole numbers.

4. Elements and compounds are composed of defi nite arrangements of atoms. Chemical change occurs when the atomic arrays are rearranged.

John Dalton’s atomic theory was accepted because it could be used to explain several scientifi c laws that Dalton and other scientists of the time had established. Two of these laws are (1) the law of conserva-tion of matter and (2) the law of defi nite proportions.

Some years before Dalton proposed his atomic theory, Antoine Lavoisier (1743–1794) had carried out a series of experiments in which the reactants were carefully weighed before a chemical reaction and the products were carefully weighed afterward. He found no change in mass when a reaction occurred,

© Amanda Hall/Robert Harding World Imagery/Getty Images / © Jose Luis Pelaez, Inc./Image Source



proposed that this was true for every reaction, and called his proposal the law of conservation of matter: Matter is neither lost nor gained during a chemical reaction.Others verifi ed his results, and the law became accepted. The second and fourth points in Dalton’s theory imply the same thing: If each kind of atom has a particular characteristic mass, and if there are exactly the same numbers of each kind of atom before and after a reaction, the masses before and after must also be the same.

Another chemical law known in Dalton’s time had been proposed by Joseph Louis Proust (1754–1826) as a result of his analyses of minerals. Proust found that a particular compound, once purifi ed, always contained the same elements in the same ratio by mass. One such study, which Proust made in 1799, involved copper car-bonate. Proust discovered that regardless of how copper carbonate was prepared in the laboratory or how it was isolated from nature, it always contained the same pro-portions by mass—fi ve parts copper, four parts oxygen, and one part carbon. Careful analyses of this and other compounds led Proust to propose the law of defi nite proportions: In a compound, the constituent elements are always present in a defi nite proportion by mass. For

example, pure water, a compound, is always made up of 11.2% hydrogen and 88.8% oxygen by weight. Pure table sugar, another compound, always con-tains 42.1% carbon, 6.5% hydrogen, and 51.4% oxy-gen by weight. The source of the pure substance is irrelevant.

3.2 Structure of the Atom

Scientists now have experimental evidence for the exis-tence of more than 60 subatomic particles. However, only three are important to our understanding of the chemical view of matter: protons and neutrons, found in the nucleus of the atom, and electrons, found outside the nucleus (Table 3.1). The mass and charge of electrons and protons were determined with experiments that used electric and magnetic fi elds. Electrons and protons have the same quantity of charge but different signs. The mass of the positively charged proton is about the same as the mass of the neutron. Both have masses 1800 times the mass of the negatively charged electron. Although stud-ies in the early 1900s indicated the presence of a third type of particle with no charge but a mass similar to that of the proton, the lack of charge meant that electric and magnetic fi elds could not be used to detect this type of particle. As a result, the neutron was not discovered until 1932, when James Chadwick (1891–1974) detected neu-trons with experiments using emissions from radioactive elements.

Even though protons and electrons had been identi-fi ed as subatomic particles, the arrangement of these par-ticles within the atom was not known. An atom was thought to be a uniform sphere of protons within which electrons circulated in rings. Only after the discovery of natural radioactivity did Ernest Rutherford (1871–1937) carry out experiments that led to the idea of a nucleus as a tiny core of the atom.

Natural Radioactivity

Henri Becquerel (1852–1908) discovered natu-ral radioactivity in natural uranium and ra-

dium ores in 1896 (Section 13.1). In 1898, Marie Sklodowska Curie (1867–1934)—a student of Becquerel and the fi rst person honored with two Nobel Prizes—and her husband Pierre discovered two radioactive elements, radium and polonium. In 1899, Marie Curie suggested that atoms of radio-

active substances disintegrate when they emit these unusual rays. She named this phe-

nomenon radioactivity. A given radioactive element gives off exactly the same type of radio-

active particles or rays regardless of whether it is © S

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Law of conservation of matter: Matter is neither lost

nor gained during a chemical

reaction

Law of definite proportions: In a

compound, the constituent

elements are always present in a

definite proportion by weight

Proton: A positively charged

subatomic particle found in the

nucleus

Neutron: An electrically

neutral subatomic particle found

in the nucleus

Nucleus: Small central core

of the atom; contains the

protons and neutrons

Electron: A negatively

charged subatomic particle found

in the space around the nucleus

Radioactivity:

Spontaneous decomposition of

unstable atomic nuclei



found in its pure state or combined with other elements. About 25 elements exist only in radioactive forms. Marie Curie’s suggestion that atoms disintegrate contradicted Dalton’s idea that atoms are indivisible.

Ernest Rutherford (1871–1937) began studying the radiation emitted from radioactive elements soon after ex-periments by the Curies and others had shown that three types of radiation may be spontaneously emitted by radio-active elements (Section 13.1). These are referred to as alpha (α) particles, beta (β) particles, and gamma (γ) rays. They behave differently when they pass between electrically charged plates, as shown in Figure 3.1. Alpha and beta rays are defl ected, while gamma rays pass straight through undefl ected. This implies that alpha and beta rays are electrically charged particles, because particles with a charge would be attracted or repelled by the charged plates. Even though an alpha particle has an electrical charge (+2) twice as large as that of a beta particle (–1), alpha particles are defl ected less; hence, alpha particles must be heavier than beta particles. Careful studies by Rutherford showed

that alpha particles are equivalent to helium atoms that have lost two electrons and thus have a +2 charge (He2+). Beta particles are negatively charged particles identical to electrons. Gamma rays have no detect-able charge or mass—they behave like rays of light.

The Nucleus of the Atom

Rutherford’s experiments with alpha particles led him to consider using them in experiments on the structure of the atom. In 1909, he suggested to two of his co-workers that they bombard a piece of gold foil with alpha particles. Hans Geiger (1882–1945), a German physicist, and Ernest Marsden

Figure 3.1 Separation of alpha and beta particles and gamma rays by an electric fi eld. Alpha particles are defl ected toward the negative plate, beta particles are attracted toward the positive plate, and

gamma rays are not defl ected. Additional studies showed that alpha particles are high-energy helium nuclei, beta

particles are high-energy electrons, and gamma rays are high-energy electromagnetic radiation.

Electrically

charged plates

Radioactive

material

Alpha particle: A

positively charged particle

emitted by certain radioactive

isotopes

Beta particle: An electron

ejected at high speeds from the

nuclei of certain radioactive

isotopes

Gamma ray: A high-

energy electromagnetic radiation

emitted from radioactive

isotopes

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Table 3.1 Summary of Properties of Electrons, Protons, and Neutrons

Relative Charge Mass (g) Mass (amu)*Approximate Relative Mass

(amu)* Location

Electron –1 0.000911 × 10–24 0.00055 ,0 Outside the nucleus

Proton +1 1.672623 × 10–24 1.00727 ,1 Nucleus

Neutron 0 1.674929 × 10–24 1.00867 ,1 Nucleus

*amu is the abbreviation for atomic mass unit (page 37); 1 amu = 1.6605 × 10–24 g

Marie

Sklodowska

Curie



(1889–1970), an undergraduate student, set up the appa-ratus diagrammed in Figure 3.2 and observed what hap-pened when alpha particles hit the thin gold foil. Most passed straight through, but Geiger and Marsden were amazed to fi nd that a few alpha particles were defl ected through large angles, and some came almost straight back. Rutherford later described this unexpected result by saying, “It was about as incredible as if you had fi red a 15-inch [artillery] shell at a piece of paper and it came back and hit you.”

What allowed most of the alpha particles to pass through the gold foil in a rather straight path? According to Rutherford’s interpretation, the atom is mostly empty space and therefore offers little resistance to the alpha particles (Figure 3.3).

What caused a few alpha particles to be defl ected? According to Rutherford’s model of the atom, all of the positive charge and most of the mass of the atom must be concentrated in a very small volume at the center of the atom. He named this part of the atom, which contained most of the mass of the atom and all of the positive charge, the nucleus. When an alpha particle passes near the nucleus, the positive charge of the nucleus repels the positive charge of the alpha particle; the path of the smaller alpha particle is consequently defl ected. The closer an alpha particle comes to a target nucleus, the more it is defl ected. Those alpha particles that meet a nucleus head-on bounce back toward the source as a re-sult of the strong positive–positive repulsion, since the

alpha particles do not have enough energy to penetrate the nucleus.

Rutherford’s calculations, based on the observed de-fl ections, indicated that the nucleus is a very small part of an atom. The diameter of an atom is about 100,000 times greater than the diameter of its nucleus.

Truly, Rutherford’s model of the atom was one of the most dramatic interpretations of experimental evidence to come out of this period of signifi cant discoveries.

3.3 Modern View of the Atom

Early experiments on the structure of the atom clearly showed that the three primary constituents of atoms are electrons, protons, and neutrons. The nucleus or core of the atom is made up of protons with a positive electrical charge and neutrons with no charge. The electrons, with a negative electrical charge, are found in the space around the nucleus (Figure 3.4). For an atom, which always has no net electrical charge, the number of negatively charged electrons around the nucleus equals the number of posi-tively charged protons in the nucleus.

Atoms are extremely small—far too small to be seen with even the most powerful optical microscopes. The di-ameters of most atoms range from 1 × 10–8 cm to

5 × 10–8 cm. For example, the diameter of a gold atom is 3 × 10–8 cm. To visualize how small this is, consider that it would take approximately 517 million gold atoms to run the length (15.5 cm) of a dollar bill. If this weren’t hard enough to imagine, remember that Ruther-ford’s experiments provided ev-idence that the diameter of the nucleus is 100,000 times smaller than the diameter of the atom. For example, if an atom were scaled upward in size so that the nucleus was the size of a small marble, the atom would be the size of Houston’s Reliant Astro-dome, and most of the space in between would be empty. Be-cause the nucleus carries most

Figure 3.2 Rutherford gold foil experiment. A beam of positively charged alpha particles was

directed at a very thin piece of gold foil. A luminescent screen was used to detect particles passing through or

deflected by the foil. Most particles passed straight through. Some were deflected slightly, and a few were deflected

back toward the source. (Rutherford actually used a movable luminescent screen instead of the circular screen shown.)

Undeflected particles

Goldfoil

Deflectedparticles

ZnS fluorescentscreen

Source of narrow beamof fast-moving particles



of the mass of the atom (the neutrons and protons) in such a small volume, a small matchbox full of atomic nuclei would weigh more than 2.5 billion tons. The interior of a collapsed star is made up of nuclear material estimated to be nearly this dense.

Scientists have been able to obtain computer-enhanced images of the outer surface of atoms (Figure 3.5) using the scanning tunneling microscope (STM) and the atomic force microscope (AFM).

Atomic Number

The atomic number of an element indicates the number of protons in the nucleus of the atom. All atoms of the same element have the same number of protons in the nucleus. In the periodic table (Figure 3.11), the atomic number for each element is given above the element’s symbol. There is a different atomic number for each ele-ment, beginning with the atomic number 1 for hydrogen. Phosphorus, for example, has an atomic number of 15; thus, the nucleus of a phosphorus atom contains 15 protons.

It is the number of pro-tons, and nothing more, that determines the identity of an atom. In a neutral atom,

Atomic number: The

number of protons in the nucleus

of an element

Figure 3.3 Rutherford’s interpretation of the gold foil experiment done by Geiger and Marsden. Each circle represents an atom, and the dots

represent their nuclei. The gold foil was about 1000 atoms thick.

A few particles collide head-on withnuclei and are deflected back toward the source

Atoms in a thinsheet of gold Nucleus

Electrons occupy spaceoutside nucleus

Some particles aredeflected considerably

Most particles passstraight through or aredeflected very little

Beam of particles

Figure 3.4 Model of atom. All atoms consist of one or more

protons (positively charged) and usually at least as many neutrons (no

charge) packed into an extremely small nucleus. Electrons (negatively

charged) are arranged in a cloud around the nucleus.

Region occupiedby electrons

Approximately 10 10m (0.1 nm)

Approximately 10 15 m



the number of protons is equal to the number of elec-trons; therefore, the atomic number also gives the num-ber of electrons, all of which exist outside the nucleus in an atom of an element.

Mass Number

The mass number of a particular atom is the total num-ber of neutrons and protons present in the nucleus of an atom. Because the atomic number gives the number of protons in the nucleus, the difference between the mass number and the atomic number equals the number of neutrons in the nucleus.

A notation frequently used for showing the mass number and atomic number of an atom places subscripts and superscripts to the left of the symbol.

19F9 Symbol of elementMass numberAtomic number

The subscript giving the atomic number is optional because the element sym-bol tells you what the atomic number must be.

For example, the fl uorine atom can have the notation 9

19F or simply 19F. For an atom of fl uorine,

919F, the number of

protons is 9, the number of electrons is also 9, and the number of neutrons is 19 – 9 = 10.

EXAMPLE 3.1 Atomic Composition

DHow many protons, neutrons, and electrons are

in an atom of gold (Au) with a mass number

of 197?

SOLUTION

The atomic number of an element gives the number of protons and electrons. If the element is known, its atomic number can be obtained from the periodic table (on the periodic table review card). The atomic number of gold is 79. Gold has 79 protons and 79 electrons. The number of neutrons is obtained by subtracting the atomic number from the mass number:

197 – 79 = 118 neutrons

TRY IT 3.1

How many protons, neutrons, and electrons are there in a

2859Ni atom?

Figure 3.5 Scanning tunneling microscope (STM). (a) When an electric current passes through a tungsten needle with a narrow tip (atom’s width) into the atoms

on the surface of the sample being examined, the electron flow between the tip and the surface changes in relation to the electron clouds around the atoms. By adjusting the position

of the needle to maintain a constant current, the positions of the atoms are measured. (b) An STM image depicting a field of xenon (Xe) atoms. The sample of matter being imaged is

600 billionths of an inch (16.8 nm) wide.

Atom at tip

of probe

Scanning

direction

Surface

atoms

Electrons leak

between probe

and surface.

Tip moves up and

down to keep

current fixed.

(b)(a)

Mass number: The

number of neutrons and protons

in the nucleus of an atom

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To represent isotopes with words

instead of symbols, the mass number

is added to the name—for example,

neon-20, neon-21, and neon-22.

Isotopes

When a natural sample of almost any element is analyzed, the

element is found to be composed of atoms with different mass numbers. Atoms of the same ele-ment having different

mass numbers are called isotopes of that element.

The element neon is a good example to consider. A natural sample of neon gas is found to be a mixture of three isotopes of neon:

1020Ne

1021Ne

1022Ne

The fundamental difference between isotopes is the dif-ferent number of neutrons per atom. All atoms of neon have 10 electrons and 10 protons. About 90.92% of the atoms have 10 neutrons, 0.26% have 11, and 8.82% have 12. Because they have different numbers of neutrons, it follows that they must have different masses. Note that all of the isotopes have the same atomic number because they are all neon. An analogy may be helpful here. Con-sider three brand new, identical automobiles that have had different numbers of 50 lb. weights hidden in their trunks. The automobiles are identical in every way ex-cept for their different masses.

Over 100 elements are known, and more than 1000 isotopes have been identifi ed, many of them produced artifi cially (see Section 13.5). Some elements have many

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The Isotopes, whose mascot Orbit is

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based on Rutherford’s model of the

atom.

isotopes; tin, for example, has ten natural isotopes. Hy-drogen has three isotopes, which are the only known iso-topes generally referred to by different names:

11H is

commonly called hydrogen or protium, 12H is called deu-

terium, and 13H is called tritium. Tritium is radioactive.

EXAMPLE 3.2 Isotopes

DCarbon has seven known isotopes. Three of

these have six, seven, and eight neutrons, re-

spectively. Write the complete chemical notation for

these three isotopes, giving mass number, atomic num-

ber, and symbol.

SOLUTION

The atomic number of carbon is 6. The mass number for the three isotopes is 6 plus the number of neutrons, which equal 12, 13, and 14, respectively.

612C

613C

614C

TRY IT 3.2

Silver has two natural isotopes, one with 60 neutrons and the other with 62 neutrons. Give the complete chemical notation for these isotopes.

Atomic Mass and Atomic Weight

Although Dalton knew nothing about subatomic particles, he proposed that atoms of different elements have differ-ent masses. Eventually, it was found that an oxygen atom is about 16 times heavier than a hydrogen atom. This fact, however, does not tell us the mass of either atom. These are relative masses in the same way that a grapefruit may weigh twice as much as an orange. This information gives neither the mass of the grapefruit nor that of the orange. Nevertheless, if a specifi c number is assigned as the mass of any particular atom, this fi xes the numbers assigned to the masses of all other atoms. The present atomic mass scale, adopted by scientists worldwide in 1961, is based on assigning the mass of a particular isotope of the car-bon atom, the carbon-12 isotope, as exactly 12 atomic mass units (amu).

The atomic masses are given in the periodic table below the symbol for the ele-ment, as illustrated at the top of the next page for copper.

Isotopes: Atoms of an

element with different mass

numbers due to different

numbers of neutrons

Atomic mass unit (amu): The unit for elements’

relative atomic masses

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spectrum and is obtained by passing sunlight or light from an incandescent lightbulb through a glass prism. When we see a rainbow, we are looking at the continuous visible spectrum that forms when raindrops act as prisms and disperse the sunlight. The different colors of light have different wavelengths. Red light has longer wave-lengths than does blue light, but in the continuous spec-trum the colors merge from one to another with no break in the spectrum. White light is a combination of all the colors of different wavelengths.

If a high voltage is applied to an element in the gas phase in a partially evacuated tube, the atoms absorb en-ergy and are said to be “excited.” The excited atoms emit light. An example of this is a neon advertising sign in which excited neon atoms emit orange-red light (Figure 3.6a). When light from such a source passes through a prism, a different type of spectrum is obtained, one that is not continuous but has characteristic lines at specifi c wavelengths (Figure 3.6b). This type of spectrum is called a line emission spectrum.

Visible light is only a small portion of the electro-magnetic spectrum (Figure 3.7). Ultraviolet radiation, the type that leads to sunburn and some forms of skin cancer, has wavelengths shorter than those of visible light; X-rays and gamma rays (the latter emitted from radioactive atoms) have even shorter wavelengths. Infra-red radiation, the type that is sensed as heat from a fi re, has longer wavelengths than visible light. Longer still are the wavelengths of the types of radiation in a microwave oven and in television and radio transmissions. Although the example of a line spectrum shown in Figure 3.6b is only for the visible region, excited atoms of elements also emit characteristic wavelengths in other regions of the electromagnetic spectrum, as demonstrated by the experiments described in the next section.

Bohr Model of the Atom

In 1913, Niels Bohr introduced his model of the hydro-gen atom. He proposed that the single electron of the hy-drogen atom could occupy only certain energy levels. He referred to these energy levels as orbits and represented the energy difference between any two adjacent orbits as a single quantum of energy. When the hydrogen elec-tron absorbs a quantum of energy, it moves to a higher energy level. When this electron returns to the lower, more stable energy level, the quantum of energy is emit-ted as a specifi c wavelength of light.

In Bohr’s model, each allowed orbit is assigned an integer, n, known as the principal quantum number. The values of n for the orbits range from 1 to infi nity. Any

29

Cu63.546

Atomic number

Symbol

Atomic weight

The periodic table entry for copper. (The complete periodic

table and its signifi cance are discussed in Section 3.6.)

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These values are average masses, which take into account the relative abundances of the different isotopes as found in nature. This average is often referred to as the atomic weight of the element. For example, boron has two natu-rally occurring isotopes,

510B and

511B, with natural percent

abundances of 19.91% and 80.09%. The masses in atomic mass units are 10.0129 and 11.0093, respectively. The atomic weight listed in the periodic table is the average mass of a natural sample of atoms, expressed in atomic mass units. The average mass of 10.81 for B is closer to the mass of

511B because of its higher abundance.

3.4 Where Are the Electrons in Atoms?

Experiments on the interactions of light with matter pro-vide important information about the energy and location of electrons in atoms.

Continuous and Line Spectra

We are familiar with the spectrum of colors that makes up visible light. The spectrum of white light is a display of separated colors. This type of spectrum is re-ferred to as a continuous

Atomic weight: The

average atomic mass of an

element’s isotopes weighted by

percentage abundance

Continuous spectrum:A spectrum that contains

radiation distributed over all

wavelengths

Quantum: The smallest

increment of energy



atom with its electrons in their normal, lowest energy level is said to be in the ground state. Energy must be supplied to move the electron farther away from the nu-cleus because the positive nucleus and the negative elec-tron attract each other. When the electron of a hydrogen atom occupies a higher energy level with n greater than 1, the atom has more energy than in its ground state and is said to be in an excited state. The excited state is an unstable state, and the extra energy is emitted when the electron returns to the ground state. According to Bohr, the light forming the lines in the brightline emission spectrum of hydrogen comes from electrons moving to-

ward the nucleus after having fi rst been excited to energy levels farther from the nucleus. Because the energy lev-els have only certain energies, the emitted light has only certain wavelengths.

With brilliant imagination, Bohr applied a little alge-bra and some classic mathematical equations of physics to his tiny solar system model of the hydrogen atom and was able to calculate the wavelengths of the lines in the hydrogen spectrum. By 1900, scientists had measured the wavelengths of lines for hydrogen in the ultraviolet, visible, and infrared regions, and Bohr’s calculated val-ues agreed with the measured values. Niels Bohr had tied

the unseen (the interior of the atom) with the seen (the observable lines in the hy-drogen spectrum)—a fan-tastic achievement. The concepts of quantum num-ber and energy level are valid for all atoms and molecules.

The Bohr model was accepted almost immedi-ately after its presentation, and Bohr was awarded the

Figure 3.6 Neon. (a) A partially evacuated tube that contains neon gas gives a reddish-orange glow when high voltage is applied. (b) The line emission spectrum

of neon is obtained when light from a neon source passes through a prism.

(b)(a)

Figure 3.7 The electromagnetic spectrum. Visible light (enlarged section) is but a small part of the entire

spectrum. The energy of electromagnetic radiation increases from the radio wave end to the gamma ray end. The frequency of

electromagnetic radiation is related to the wavelength by vλ = c, where v = frequency;

λ = wavelength; and c = speed of light, 3.00 × 108 meters (m)/second (s). The higher

the frequency, the lower the wavelength and the larger the energy. The energy (E) of a

photon (or quantum) of light is given by the expression Ephoton

= hv, where h is Planck’s

constant (6.6262 × 10–34 J • s).

Rays(high energy)

X-rays UV

Visible spectrum

Wavelength (nanometers)

400

10 16 10 14 10 12 10 10 10 8

500 600 700

IR Microwave

Radio waves

FM AM Long radio waves(low energy)

Longer wavelength(lower energy)

Shorter wavelength(higher energy)

10 6 10 4 10 2 100 102 104 106 108 Wavelength (meters)

Frequency (sec 1)1024 1022 1020 1018 1016 1014 1012 1010 108 106 104 102 100

Increasing energy ( )Increasing frequency ( )Increasing wavelength ( )

Ground state: The

condition of an atom in which all

electrons are in their normal,

lowest energy level

Excited state: An unstable,

higher energy state of an atom

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It is important to understand the relationships among l, n, and EAs lT, nc EcAs lc, nT ETWavelength and frequency (or energy) are inversely related.



that each energy level could hold only a characteristic number of electrons and

that only those electrons in the high-est occupied level were involved when one atom combined with an-

other. These outermost electrons came to be known as valence elec-

trons. The number of valence electrons an atom has is important in determining how the atom combines with other atoms.

Energy Level (n) 2n2

Maximum Number of Electrons

1 2(1)2 2

2 2(2)2 8

3 2(3)2 18

4 2(4)2 32

5 2(5)2 50

6 2(6)2 72

For example, look at phosphorus (P). The Bohr ar-rangement of electrons is 2–8–5. This means that the stable state of the phosphorus atom has electrons in three energy levels. The one closest to the nucleus has two electrons, the second energy level has eight electrons, and the highest energy level has fi ve electrons. The en-ergy level with fi ve electrons is farthest from the nucleus; thus, P has fi ve valence electrons, and these are the elec-trons that are most available for interactions with valence electrons of other atoms in chemical reactions. This in-teraction results in the formation of bonds, which will be discussed in Chapter 5.

Beyond the Bohr Atom

The simple Bohr model of the atom, in which each shell is capa-ble of holding 2n2 electrons, can be used to explain many properties of atoms and their associated electrons. However, more sophisticated treatments of atoms depend on the fact that electrons moving at very high speeds, up to 2 million meters per second (or roughly 1 percent the speed of light), exhibit

he Bohr

r model of thech shell is capa-electrons, can be

many properties of ssociated electrons.

sophisticated ms dependlectrons

high milliond (or t the hibit

Nobel Prize in physics in 1922 for his contribution to the understanding of the hydrogen atom. However, his model gave only approximate agreement with line spectra of atoms having more than one electron. Later mod-els of the atom have been more suc-cessful by considering electrons as having both particle and wave characteristics. This led to mathematical treatment of the loca-tions of electrons as probabilities instead of as the precise locations envisioned by Bohr. Thus, in the modern view of the atom, we can picture a space around the nucleus occupied by each electron, somewhat like a cloud of electrical charge of a particular energy. We just don’t know where within the cloud a particular electron is at any given instant. Bohr’s concept of the main energy levels represented by the quantum number n remains valid, however; for our purposes, this concept is all that we need to discuss the location of electrons in atoms.

Atom Building Using the Bohr Model

Recall that the atomic number is the number of electrons (or protons) per atom of an element. Imagine building atoms by adding one electron to the appropriate energy level as another proton is added to the nucleus. As part of his theory, Bohr proposed that only a fi xed number of electrons could be accommodated in any one level, and he calculated that this number was given by the formula 2n2, where n equals the number of the energy level. For the lowest energy level (fi rst level), n equals 1, and the maximum number of electrons allowed is 2(1)2, or 2. For the second energy level, the maximum number of elec-trons is 2(2)2, or 8. By using 2n2, the maximum number of electrons allowed for levels 3, 4, and 5 are 18, 32, and 50, respectively. A general overriding rule to the preced-ing numbers is that the highest energy level can have no more than eight electrons, except for transition elements, for a stable atom. Table 3.2 lists the Bohr electron ar-rangements for the fi rst 20 elements.

Electrons in the highest occupied energy level listed for the elements in Table 3.2 are at the greatest stable dis-

tance from the nucleus. These are the most impor-tant electrons in the study of chemistry because they are the ones that interact when atoms react with each other. G. N. Lewis fi rst proposed

Valence electrons: The

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properties of waves. Although it is appropriate at some level of understanding to envision electrons as particles moving around the nucleus like planets moving around the sun, electron locations around the nucleus can be

more accurately described by using wave the-ory. This theory defi nes electron locations in

terms of regions of probability known as or-bitals. A full treatment of highly mathemat-

ical wave theory, developed by an Austrian scientist named Erwin Schrodinger, is well beyond the scope of this text, but his results

are nonetheless useful.In essence, Schrodinger con-

cluded that the volume of space in

which each pair of electrons could most likely be found defi nes an or-bital and that each orbital can ac-commodate a maximum of two electrons. Furthermore, he con-cluded that Bohr’s planetary shells can contain more than one such or-bital and the orbitals within a given shell may be of different types. Quantum theory shows that the two electrons residing in the fi rst Bohr shell occupy a single, spherical or-bital called a 1s orbital (Figure 3.8).

The second Bohr shell can ac-commodate a total of eight electrons (2 × 22), but quantum theory shows that these eight electrons occupy two different types of orbitals. This sec-ond energy shell thus contains two subshells comprised of a single spherical 2s orbital and three sepa-rate, dumbbell-shaped orbitals ori-ented at right angles with respect to each other and known as 2p orbitals. All of these orbitals, like all atomic orbitals regardless of type, are cen-tered about the nucleus of the atom.

The 2s orbital is higher in energy than the 1s orbital, but it is slightly lower in energy than the three 2p orbitals. The 2s orbital holds two electrons and the three 2p orbitals each hold two electrons for a total of eight electrons in the second energy shell.

A continuation of this treatment showed that the third energy level, which can accommodate a maximum of 2 × 32 or 18 electrons, is actually composed of three sub-shells. The lowest energy orbital in this shell is the spheri-cal 3s orbital, capable of holding two electrons. The next-lowest energy orbitals in this shell are three dumb-bell-shaped 3p orbitals (Figure 3.8) capable of holding a total of six electrons. Lastly, the third energy shell con-tains fi ve orbitals known as 3d orbitals. These fi ve 3d

Table 3.2 Electron Arrangements of the First 20 Elements*

Number of Electrons in Each Energy Level

ElementAtomic

Number 1st 2nd 3rd 4th

Hydrogen (H) 1 1 e

Helium (He) 2 2 e

Lithium (Li) 3 2 e 1 e

Beryllium (Be) 4 2 e 2 e

Boron (B) 5 2 e 3 e

Carbon (C) 6 2 e 4 e

Nitrogen (N) 7 2 e 5 e

Oxygen (O) 8 2 e 6 e

Fluorine (F) 9 2 e 7 e

Neon (Ne) 10 2 e 8 e

Sodium (Na) 11 2 e 8 e 1 e

Magnesium (Mg) 12 2 e 8 e 2 e

Aluminum (Al) 13 2 e 8 e 3 e

Silicon (Si) 14 2 e 8 e 4 e

Phosphorus (P) 15 2 e 8 e 5 e

Sulfur (S) 16 2 e 8 e 6 e

Chlorine (Cl) 17 2 e 8 e 7 e

Argon (Ar) 18 2 e 8 e 8 e

Potassium (K) 19 2 e 8 e 8 e 1 e†

Calcium (Ca) 20 2 e 8 e 8 e 2 e

* Valence electrons are shown in blue.† The preceding discussion, Beyond the Bohr Atom (page 40), more clearly explains why the nineteenth and twentieth electrons do not add to the third energy level as might have been predicted using the 2n2 guide.

Orbital: A region of three-

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shown in Figure 3.9 and place the appropri-ate number of electrons in each orbital or set of orbitals in the order specifi ed. Remember that there are 2 elec-trons in the 2s orbitals, 6 electrons in the 2p or-bitals, 10 electrons in the 3d orbitals, and 14 electrons in the 4forbitals.

Using this method, we can write the elec-tron confi guration of each atom as shown in Table 3.3, where super-scripts are used to indi-cate the number of

electrons in the orbitals. All electrons that occupy a sub-shell with the same main shell number (for example, where the second shell is signifi ed by using a 2 in front of s and p orbitals) are valence electrons.

This discussion helps explain why the last electron added when going from

18Ar to

19K goes in the fourth

energy level (as shown in Table 3.2) and not in the third energy level, as one might have predicted using the sim-

3s

x y

z

x y

z

x y

z

3p x 3p y 3p z

3dz2 3dxz 3dyz 3dxy 3dx

2 – y

2

2p x 2p y 2p z

1s

2s

Figure 3.8 Atomic orbitals. Boundary surface diagrams for electron densities of 1s, 2s, 2p, 3s, 3p, and 3d orbitals. For

the p orbitals, the subscript letter on the orbital notation (x, y, z) indicates the cartesian axis along which the orbital lies.

Figure 3.9 Subshell filling order. Subshells in atoms are filled

in order of increasing energy, as this diagram shows. The order of filling is

1s 1 2s 1 2p 1 3s 1 3p 1 4s 1 3d and so on.

Shell

8 8s

7s

6s

5s

4s

3s

2s

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5p

4p

3p

6d

5d

4d

5f

4f

3d

2p

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7

6

5

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3

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1

orbitals can accommodate a total of ten electrons. Sum-marizing, we see that the third shell contains a single 3sorbital (2-electron capacity), three 3p orbitals (6-electron capacity), and fi ve 3d orbitals (10-electron capacity) for a maximum of 18 electrons in the third shell.

The fourth energy shell contains, after the 4s orbital, three 4p orbitals, fi ve 4d orbitals, and seven 4f orbitals (not pictured) that can accommodate a maximum of 14 electrons. Thus, the fourth shell or energy level can ac-commodate a total of 2 × 42 (or 32) electrons. The trend continues in this fashion until no additional orbitals are required to accommodate the electrons.

Although the Bohr model correctly predicts the number of electrons in each energy shell according to the 2n2 calculation with n equalling the number of the energy shell, the quantum wave model more accurately describes the location probability of the electrons in three-dimensional space.

As one progresses from one atom to the atom of next higher atomic number in the periodic table, electrons are added singly to the lowest energy orbital in a given sub-shell that does not already have an electron. Once each orbital in a subshell has a single electron, additional elec-trons are then added to the singly-occupied orbitals in that subshell until they are each doubly occupied. At that point, additional electrons are then added to the orbitals with the next highest energy until one exhausts the num-ber of electrons associated with the atom.

Figure 3.9 provides an order-of-fi lling chart that al-lows the exact specifi cation of orbitals occupied by elec-trons in a given atom. Simply follow the arrowed pathway



ple Bohr model. This happens because according to the fi lling model, electrons are added to the 4s orbital before being added to the 3d orbital. Following this model, the next electron added to arrive at

20Ca goes into the 4s or-

bital; the 3d orbitals are fi lled next, before the 4p orbitals. If one were to follow the fi lling method across the entire periodic table, the transition metals (Section 3.6) would occupy the region of the periodic table, where the d and f orbitals come into play. Figure 3.10 presents a graphic, “building-up” version of the periodic table.

In summary, quantum theory’s more sophisticated treatment of electrons allows us to more accurately de-scribe the positions of electrons around a given nucleus. This can help us understand chemical bonding in a more detailed way.

3.5 Development of the Periodic Table

So far, in exploring the chemical view of matter, you have seen that everything is made of atoms. Atoms of elements combined in specifi c ratios are present in chem-ical compounds. All matter consists of elements, com-pounds, or mixtures of elements or compounds. You have also been introduced to the subatomic composition of atoms—the electrons, protons, and neutrons. The next part of the story is to see how atomic structure and the properties of elements and compounds are related to each other. Fortunately, the chemistry of the elements can be organized and classifi ed in a way that helps both chem-ists and nonchemists. The periodic table is the single most important classifi cation system in chemistry because it summarizes, correlates, and predicts a wealth of chemical information. Chemists consult it every day during every possible kind of work. It can simply be a reminder of the symbols and names of the elements, of which elements have similar properties, and of where each element lies on the continuum of atomic numbers. It can also be an inspiration in the search for new com-pounds or mixtures that will fulfi ll a specifi c need. Mem-orizing the periodic table is no more necessary than

Table 3.3 Electron Confi gurations of the First 20 Elements*

ElementAtomic

Number Element Atomic Number

Hydrogen (H) 1 1s1 Sodium (Na) 11 1s22s22p6 3s1

Helium (He) 2 1s2 Magnesium (Mg) 12 1s22s22p6 3s2

Lithium (Li) 3 1s2 2s1 Aluminum (Al) 13 1s22s22p6 3s23p1

Beryllium (Be) 4 1s2 2s2 Silicon (Si) 14 1s22s22p6 3s33p2

Boron (B) 5 1s2 2s22p1 Phosphorus (P) 15 1s22s22p6 3s23p3

Carbon (C) 6 1s2 2s22p2 Sulfur (S) 16 1s22s22p6 3s23p4

Nitrogen (N) 7 1s2 2s22p3 Chlorine (Cl) 17 1s22s22p6 3s23p5

Oxygen (O) 8 1s2 2s22p4 Argon (Ar) 18 1s22s22p6 3s23p6

Fluorine (F) 9 1s2 2s22p5 Potassium (K) 19 1s22s22p63s23p6 4s1

Neon (Ne) 10 1s2 2s22p6 Calcium (Ca) 20 1s22s22p63s23p6 4s2

* Valence electrons are shown in blue.

Physicist Max Planck

proposed in 1900 that energy comes

in discrete bundles called quanta. Something

that can have only certain values with

nothing in between is referred to

as quantized.



textbook on chemistry. He had the properties of each element written on cards, with a sepa-rate card for each element. While he was shuffl ing the cards trying to gather his thoughts before writing his manuscript, Mendeleev real-ized that if the elements were arranged in the order of their atomic weights, there was a trend in properties that repeated itself several times. He arranged the elements into groups that had similar properties and used the resulting periodic chart to predict the properties and places in the chart of as yet undiscov-ered elements.

Thus, the periodic law and the periodic table were born, although only 63 ele-ments had been discovered by 1869 (for example, the noble gases were not discovered until after 1893), and the clar-ifying concept of the atomic number was not known until 1913. Mendeleev’s idea and textbook achieved great suc-cess, and he rose to a position of prestige and fame while he continued to teach at the Uni-versity of St. Petersburg.

Mendeleev aided the dis-covery of new elements by pre-dicting their properties with remarkable accuracy, and he even suggested the geographi-cal regions in which minerals containing the elements could be found. The properties of a missing element were predicted by consideration of the proper-

ties of its neighboring elements in the table. For example, for the element we now know as germanium, which falls below silicon in the modern periodic table (Figure 3.11), Mendeleev predicted a gray element of atomic weight 72 with a density of 5.5 g/cm3. Germanium, once discovered, proved to be a gray element of atomic weight 72.59 with a density of 5.36 g/cm3.

Figure 3.10 In this “building-up” version of the periodic table, the lightest elements are at the bottom. Electrons fill subshells from bottom to top in order of energy as the atomic number of the atom

increases. The numbers across the top give the number of electrons in each subshell. The ground-state electron

configurations of most elements are apparent from their positions in the table. Those that are known to differ from

expectation are indicated explicitly.

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The empty spaces in the table and Mendeleev’s pre-dictions of the properties of missing elements stimulated a fl urry of prospecting for elements in the 1870s and 1880s, and eight more were discovered by 1886.

Mendeleev found that a few elements did not fi t under other elements with similar chemical properties when arranged according to increasing atomic weight. Eventually, it was found that the atomic weight is not the property that governs the similarities and differences among the elements. This was discovered in 1913 by H. G. J. Moseley (1888–1915), a young scientist work-ing with Ernest Rutherford. Moseley found that the wavelengths of X-rays emitted by a particular element are related in a precise way to the atomic number of that element. He quickly realized that other atomic properties may be similarly related to atomic number and not, as Mendeleev had believed, to atomic weight.

By building on the work of Mendeleev and others, and by using the concept of the atomic number, we are now able to state the modern periodic law: When ele-ments are arranged in the order of their atomic numbers, their chemical and physical properties show repeatable, or periodic, trends. Other familiar periodic phenomena include the changing of the seasons and the orbits of the planets, which are periodic with time. To illustrate this idea on a simple level, a shingle roof, which has the same pattern over and over, is periodic.

Thus, to build up a periodic table according to the pe-riodic law, the elements are lined up in a horizontal row in the order of their atomic numbers, as in Figure 3.11. At an element with similar properties to one already previously placed in the row, a new row is started. Each column then contains elements with similar properties. Some chemical

and physical properties of the fi rst 20 elements are summarized in Table 3.4. Periodic law: When

elements are arranged in the

order of their atomic numbers,

their chemical and physical

properties show repeatable, or

periodic, trends

Periodic table: An

arrangement of elements by

atomic number

Group: A vertical column of

elements in the periodic table

Period: A horizontal row of

elements in the periodic table

Metal: An element that

conducts electric current

Do you see any trends and similarities among the elements in Table 3.4? For example, lithium (Li) is a soft metal with low density that is very reactive. It combines with chlorine gas to form lithium chloride with the formula LiCl. The other elements in Table 3.4 that have properties similar to those of lithium are sodium (Na) and potassium (K). Ac-cording to the periodic law, lithium, sodium, and potassium should be in the same group—and they are. Look for simi-larities among other elements listed in Table 3.4 and check your grouping with that shown in the periodic table in Fig-ure 3.11.

3.6 The Modern Periodic Table

In the modern periodic table (Figure 3.11), the elements are arranged in order of their atomic numbers so that ele-ments with similar chemical and physical properties fall together in vertical columns. These vertical columns are called groups. The fi rst two and last six groups (Groups 1, 2, and 13 to 18) are the representative or main-group elements. The middle ten groups (Groups 3 to 12) are the transition elements that link the two areas of representa-tive elements. The inner transition elements are the lan-thanide series and the actinide series. They are placed at the bottom of the periodic table because the similarity of properties within the two series would require their placement between lanthanum and hafnium, and between actinium and rutherfordium, respectively, which would make the table inconveniently wide.

The horizontal rows are called periods. These periods or rows are related to energy levels for electrons in atoms (Figures 3.10 and 3.11). The length of a row is linked to the maximum number of electrons, 2n2, that can fi t into an energy level. The periods are not equal in size because the maximum number of electrons per energy level increases

as the distance of the energy level from the nucleus in-creases. Periods one through seven have 2, 8, 8, 18, 18, 32, and 23 (incomplete) elements, respectively.

Larger periods, seen as the atoms of elements get larger, are similar to the longer rows and more seats per row in a stadium as you proceed up

from fi eld level in the stands.Most of the elements are metals. These

are found in Groups 1, 2, and parts of

and physical properties ofthe fi rst 20 elements aresummarized in Table 3.4.

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Table 3.4 Some Properties of the First 20 Elements

Element Atomic

Number Description

Compound Formation*

With Cl (or Na) With O (or Mg)

Hydrogen (H) 1 Colorless gas, reactive HCl H2O

Helium (He) 2 Colorless gas, unreactive None None

Lithium (Li) 3 Soft metal, low density, very reactive LiCl Li2O

Beryllium (Be) 4 Harder metal than Li, low density, less reactive than Li BeCl2

BeO

Boron (B) 5 Both metallic and nonmetallic, very hard, not very reactive BCl3 B

2O

3

Carbon (C) 6 Brittle nonmetal, unreactive at room temperature CCl4 CO

2

Nitrogen (N) 7 Colorless gas, nonmetallic, not very reactive NCl3 N

2O

5

Oxygen (O) 8 Colorless gas, nonmetallic, reactive Na2O, Cl

2O MgO

Fluorine (F) 9 Greenish-yellow gas, nonmetallic, extremely reactive NaF, ClF MgF2, OF

2

Neon (Ne) 10 Colorless gas, unreactive None None

Sodium (Na) 11 Soft metal, low density, very reactive NaCl Na2O

Magnesium (Mg) 12 Harder metal than Na, low density, less reactive than Na MgCl2

MgO

Aluminum (Al) 13 Metal as hard as Mg, less reactive than Mg AlCl3 Al

2O

3

Silicon (Si) 14 Brittle nonmetal, not very reactive SiCl4 SiO

2

Phosphorus (P) 15 Nonmetal, low melting point, white solid, reactive PCl3 P

2O

5

Sulfur (S) 16 Yellow solid, nonmetallic, low melting point, moderately reactive Na2S, SCl

2MgS

Chlorine (Cl) 17 Yellow-green gas, nonmetallic, extremely reactive NaCl MgCl2, Cl

2O

Argon (Ar) 18 Colorless gas, unreactive None None

Potassium (K) 19 Soft metal, low density, very reactive KCl K2O

Calcium (Ca) 20 Harder metal than K, low density, less reactive than K CaCl2

CaO

* The chemical formulas shown are lowest ratios. The molecular formula for AlCl3 is Al

2Cl

6 , and that for P

2O

5 is P

4O

10.

Groups 13 to 16 (red in Figure 3.11). All of the transi-tion elements are metals (yellow in Figure 3.11). Char-acteristic physical properties of metals include malleability (ability to be beaten into thin sheets such as aluminum foil), ductility (ability to be stretched or drawn into wire such as copper), and good conduction of heat and electricity.

Eighteen elements are nonmetals (in green in Figure 3.11), and except for hydrogen they are found on the right side of the periodic table. Hydrogen is shown above Group 1 because its atoms have one electron. However, hydrogen is a nonmetal and probably should be in a group by itself, although you may see H in both Group 1 and Group 17 in some periodic tables. Hydrogen forms compounds with formulas similar to those of the Group

1 elements but with vastly different properties. For ex-ample, compare NaCl (table salt) with HCl (a strong acid), or compare Na

2O (an active metal oxide) with H

2O

(water). Hydrogen also forms compounds similar to those of the Group 17 elements: NaCl and NaH (sodium hydride), and CaBr

2 and CaH

2 (calcium hydride).

The physical and chemical properties of nonmetals are opposite those of metals. For example, nonmetals are insulators; that is, they are extremely poor conductors of heat and electricity.

Elements that border the staircase in Figure 3.11—those between met-als and nonmetals—are

Nonmetal: An element that

does not conduct electrical

current

Insulator: A poor conductor

of heat and electricity



six metalloids (in blue). Their properties are inter-mediate between those of metals and nonmetals. For example, silicon (Si), ger-manium (Ge), and arsenic (As) are semiconductorsand are the elements that form the basic components of computer chips. Semi-conductors conduct elec-tricity less well than metals,

such as silver and copper, but better than insulators, such as sulfur. The six noble gases in Group 18 have little ten-dency to undergo chemical reactions and at one time were considered inert (totally unreactive). The classifi cations of metals, nonmetals, and metalloids will enable you to pre-dict the kinds of compounds formed between elements.

EXAMPLE 3.3 Periodic Table

DFor the elements with atomic numbers 17, 33,

and 82, give the names and symbols and iden-

tify the elements as metals, metalloids, or nonmetals.

SOLUTION

Chlorine (Cl) is the element with atomic number 17. It is in Group 17. Chlorine and all the other elements in Group 17 are nonmetals.

Arsenic (As) is the element with atomic number 33. It is in Group 15. Because it lies along the line between metals and nonmetals, arsenic is a metalloid.

Lead (Pb) is the element with atomic number 82. It is in Group 14. Like other elements at the bottom of Groups 13 to 16, lead is a metal.

TRY IT 3.3

List the main-groups in the periodic table that (a) consist entirely of metals, (b) consist entirely of nonmetals, and (c) include metalloids. Identify the numbers of valence elec-trons in atoms from groups listed under (a), (b), and (c).

3.7 Periodic Trends

Why do elements in the same group in the periodic table have similar chemical behavior? Why do metals and

Metalloid: An element with

properties intermediate between

those of metals and nonmetals

Semiconductor: A material

with electrical conductivity

intermediate between those of

metals and insulators

Noble gas: An element in

Group 18 of the periodic table

Table 3.5 Lewis Dot Symbols for Atoms

1 2 13 14 15 16 17 18

H· He:

Li· ·Be· ·B·

· ·C··· ·N

···· :O

···· :F

····· :Ne

···· :

Na· ·Mg· ·Al·

· ·Si·· · ·P

···· :S

···· :Cl

····· :Ar

···· :

K· ·Ca·

nonmetals have different properties? G. N. Lewis was seeking answers to these questions during his develop-ment of the concept of valence electrons. He assumed that each noble gas atom had a completely fi lled outer-most shell, which he regarded as a stable confi guration because of the lack of reactivity of noble gases. He also assumed that the reactivity of other elements was infl u-enced by their numbers of valence electrons.

Lewis Dot Symbols

Lewis used the element’s symbol to represent the atomic nucleus together with all but the outermost electrons. The valence electrons, which are the outermost elec-trons of an atom, are represented by dots. The dots are placed around the symbol one at a time until they are used up or until all four sides are occupied; any remain-ing electron dots are paired with the ones already there. Lewis dot symbols for atoms of the fi rst 20 elements are shown in Table 3.5. Notice that all atoms of elements in a given main-group have the same number of valence electrons. The importance of valence electrons in the study of chemistry cannot be overemphasized. The iden-tical number of valence electrons primarily accounts for the similar properties of elements in the same group. The chemical view of matter is primarily concerned with what valence electrons are doing in the course of chemical reactions.

Lewis dot symbols will be used extensively in Chap-ter 5 in the discussions of bonding.

Atomic Properties

From left to right across each period, metallic character gives way to nonmetallic character (Figure 3.11). The elements with the most metallic character are at the lower-left part of the periodic table near cesium (Cs). The elements with the most nonmetallic character are at the upper-right portion of the periodic table near fl uorine.



The six metalloid elements (in blue) that begin with boron and move down like a staircase to as-tatine (At) roughly separate the metals and the nonmetals.

Atomic radii show periodic-ity by increasing with atomic number down the main-groups of the periodic table (Figure 3.12). Why do atoms get larger from the top to the bottom of a group? The larger atoms simply have more energy levels occu-pied by electrons than do the smaller atoms.

Atomic radii decrease across a period from left to right (Figure 3.12). You may see a paradox in adding electrons and getting smaller atoms, but protons also are being added. The increasing positive charge of the nucleus pulls electrons closer to the nu-cleus and causes contraction of the atom.

We can also use the trends in size of atomic radii to predict trends in reactivity. The valence electrons in larger atoms are far-ther from the nucleus. The larger the atom, the easier it is to remove the valence electrons because the attractive forces between protons in the nucleus and valence elec-trons decrease with increasing size of the atom.

In ionic compounds, atoms have gained and lost electrons to form ions, which have positive or negative charges. Metal atoms lose valence electrons to form pos-itive ions. The larger the metal atom, the greater the ten-dency to lose valence electrons and the more reactive the metal. Therefore, we would predict that the most reactive metal in Figure 3.12 is cesium (Cs), the metal with the largest radius, and this is correct. The resulting increasein reactivity of metals down a given group in the periodic table is dramatically illustrated by lithium, sodium, and potassium—the fi rst three metals in Group 1. Their atoms increase in size in this order down the group. Each ele-ment reacts with water—lithium quietly and smoothly, sodium more vigorously, and potassium much more quickly. The reactions of both sodium and potassium

Figure 3.12 Atomic radii of the main-group elements (in picometers). A picometer (pm)

is 1 × 10–12 m (100 nm). Remember that the radius is one-half of the diameter.

Li(152)

Be(111)

B(88)

C(77)

N(70)

O(66)

F(64)

Ne(70)

Na(186)

Mg(160)

Al(143)

Si(117)

P(110)

S(104)

Cl(99)

Ar(94)

K(231)

Ca(197)

Ga(130)

Ge(122)

As(121)

Se(117)

Br(114)

Kr(109)

Rb(244)

Sr(215)

In(162)

Sn(140)

Sb(141)

Te(137)

I(133)

Xe(130)

Cs(262)

Ba(217)

Tl(171)

Pb(175)

Bi(146)

Po(165)

At(140)

Rn(140)

H(37)

He(50)

Decreasing radius

Incr

ea

sin

g r

ad

ius

give off enough heat to ignite the hydrogen gas produced by the reaction, but as shown in Figure 3.13, potassium reacts with explosive violence. For elements at the bot-tom of Group 1, exposure to moist air produces a vigor-ous explosion.

Nonmetal atoms gain electrons from metals to form negative ions. The smaller the nonmetal atom, the higher the reactivity of the nonmetal. For example, fl uorine atoms are the smallest of Group 17 elements, and fl uo-rine is the most reactive nonmetal. It reacts with all other elements except three noble gases—helium, neon, and argon. The reaction of Group 17 elements with hy-drogen illustrates how the

Ion: An atom or group of

atoms with a positive or negative

charge



group as the atomic radius increases. Therefore, calcium should be more reactive than beryllium.

(c) Phosphorus and arsenic are Group 15 nonmet-als; thus, phosphorus, which has a smaller atomic radius, is predicted to be more reactive than arsenic.

TRY IT 3.4A

Which element in each pair has the larger radius—that is, which is the larger atom in each pair: (a) Ca or Ba, (b) S or Se, (c) Si or S, or (d) Ga or Br?

TRY IT 3.4B

Which element in each pair is more reactive: (a) Mg or Sr, (b) Cl or Br, or (c) Rb or Cs?

3.8 Properties of Main-Group Elements

Elements in a periodic table group have similar proper-ties but not the same properties. Some properties, as al-ready illustrated for atomic radii and reactivity, increase or decrease in a predictable fashion from top to bottom of a periodic group.

Elements in a group generally react with other ele-ments to form similar compounds, a fact accounted for by their identical numbers of valence electrons. For ex-ample, the formula for the compound of Li and Cl is LiCl; thus, you can expect there to be a compound of Rb and Cl with the formula RbCl and a compound of Cs and Cl with the formula of CsCl. Likewise, the formula Na

2O

is known; therefore, a compound with the formula Na2S

predictably exists, because oxygen and sulfur are in the same group. In general, elements in the same group of the periodic table form some of the same types of com-pounds. In fact, several main-groups have group names because of the distinctive and similar properties of the group elements.

The Group 1 elements (Li, Na, K, Rb, Cs, and Fr) are called the alkali metals. The name alkali derives from an old word meaning “ashes of burned plants.” All the

Figure 3.13 Reaction of alkali metals with water: (a) sodium, and

(b) potassium.

Cha

rles

D. W

inte

rs /

Cha

rles

D. W

inte

rs

(a)

reactivity of nonmetals decreases down the group. Fluo-rine reacts explosively with hydrogen, but the reaction with hydrogen is less violent for chlorine and is very slow for iodine.

Why are there repeatable patterns of properties across the periods in the periodic table? Again, it is be-cause there is a repeatable pattern in atomic structure, and properties depend on atomic structure. Each period begins with one valence electron for atoms of the ele-ments in Group 1. Each period builds up to eight valence electrons, and the period ends. This pattern repeats across periods two through six. As more elements are synthesized by chemists using nuclear accelerators (Sec-tion 13.6), it is possible that period seven will be com-pleted someday.

EXAMPLE 3.4 Atomic Radii and Reactivity

DWhich element in each pair is more reactive:

(a) O or S, (b) Be or Ca, or (c) P or As?

SOLUTION

(a) Oxygen and sulfur are Group 16 nonmetals. Gener-ally, the smaller the atomic radius of a nonmetal, the more reactive the nonmetal is. Because oxygen atoms are smaller than sulfur atoms, oxygen should be more reac-tive than sulfur.

(b) Beryllium (Be) and calcium (Ca) are Group 2 metals. Reactivity of metals in a given group increases down the

(b)



Recall that the only

other diatomic elements are

H2, N

2, and O

2.

alkali metals are soft enough to be cut with a knife. None are found in nature as free elements, because all combine rapidly and completely with virtually all of the nonmet-als and, as illustrated in Figure 3.13, with water. These elements form positive ions (cations) with a +1 charge. Francium, the last member of Group 1, is found only in trace amounts in nature, and all of its 21 isotopes are naturally radioactive.

The Group 2 elements (Be, Mg, Ca, Sr, Ba, and Ra) are called the alkaline earth metals. Compared with the alkali metals, the alkaline earth metals are harder, are more dense, and melt at higher temperatures. These ele-ments form cations with a +2 charge. Because the va-lence electrons of alkaline earths are held more tightly, they are less reactive than their alkali metal neighbors. All of the alkaline earth metals react with oxygen to form an oxide MO, where M is the alkaline earth.

2 M(s) + O2(g) 2 MO(s)

The halogens (F, Cl, Br, I, and At) are Group 17 elements. In the elemental state, each of these elements exists as diatomic molecules (X

2). Fluorine (F

2) and

chlorine (Cl2) are gases at room temperature, whereas

bromine (Br2) is a liquid and iodine (I

2) is a solid. This

illustrates a typical trend within a group—an increase in melting point and boiling point in going down a group. All isotopes of astatine (At) are naturally radio-active and disintegrate quickly. The name halogencomes from a Greek word and means “salt producing.” The best-known salt containing a halogen is sodium chloride (NaCl), table salt. However, there are many other halogen salts, such as calcium fl uoride (CaF

2), a

natural source of fl uorine; potassium iodide (KI), an ad-ditive to table salt that prevents goiter; and silver bro-mide (AgBr), the active photosensitive component of photographic fi lm. The halogen salts contain negatively charged ions known as anions.

The noble gases (He, Ne, Ar, Kr, Xe, and Rn) are Group 18 elements. They are all colorless gases composed of single atoms at room tempera-ture. They are referred to as “noble” because they lack chemical reactivity and generally do not react with “common” elements. Neon is the gas that glows orange-red in tubes of neon lights. Other gases and color-tinted tubes are used to give different col-ors. Radon (Rn) is natu-rally radioactive. Problems associated with indoor radon pollution are discussed in Section 13.7.

Alkali metal: An element

in Group 1 of the periodic table

Cation: An atom or group of

bonded atoms that has a positive

charge

Alkaline earth metal:

An element in Group 2 of the

periodic table

Halogen: An element in

Group 17 of the periodic table

Anion: An atom or group of

bonded atoms that has a

negative charge

Can you answer the following questions?

What are the milestones in the development

of atomic theory?

What is the experimental evidence for

the existence of subatomic particles within

atoms?

What are the three basic subatomic

particles of the atom?

What are isotopes?

Where are electrons in atoms, and how are

they arranged?

How was the periodic table developed?

Why do elements in the same group in

the periodic table have similar chemical

properties?

How does the activity of an element relate

to its position in the periodic table?

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applying your knowledge1. What is the law of conservation of matter? Give an

example of the law in action.

2. State the law of definite proportions. Give an example to

illustrate what it means.

3. Which early scientist is given credit for much of the

development of the atomic theory of atoms, molecules,

elements, and compounds?

4. What kinds of evidence did Dalton have for atoms that the

early Greeks (Democritus, Leucippus) did not have?

5. How does Dalton’s atomic theory explain these?

(a) The law of conservation of matter

(b) The law of definite proportions

6. Describe the most significant observation to be derived

from Lord Ernest Rutherford’s experiments in which he

bombarded gold foil with alpha particles.

7. The fact that carbon dioxide always contains 27.27%

carbon and 72.73% oxygen by mass is a manifestation of

the:

(a) Law of mass action

(b) Law of definite proportions

(c) Law of conservation of matter

(d) Periodic law

8. What is meant by the term subatomic particles? Give two

examples.

9. Give short definitions for the following terms:

(a) Atomic number (b) Mass number

(c) Atomic weight (d) Isotope

(e) Natural abundance (f) Atomic mass unit

10. There are more than 1000 kinds of atoms, each with a

different weight, yet only 116 elements were known as

this book was being written. How does one explain this in

terms of subatomic particles?

11. Dmitri Mendeleev was the last born of 17 children in his

family. Can you write the atomic symbol and give the total

number of electrons and the number of valence electrons

for each of the first 17 elements in his periodic table?

12. Which of the following pairs are isotopes? Explain your

answers.

(a) 50Ti and 50V (b) 12C and 14C

(c) 40Ar and 40K

13. What do all the atoms of an element have in common?

14. If 3717Cl is used to designate a certain isotope of chlorine,

what would 3517Cl represent?

(a) An isotope with two fewer neutrons

(b) An isotope with two fewer protons

(c) A different element

(d) An isotope with two more electrons

(e) An isotope with one less proton and one less neutron

15. A common isotope of Li has a mass of 7. The atomic

number of Li is 3. How can this information be used to

determine the number of protons and neutrons in the

nucleus?

16. The element iodine (I) occurs naturally as a single isotope

of atomic mass 127; its atomic number is 53. How many

protons and how many neutrons does it have in its

nucleus?

17. Which of the following elements would need only to gain

two electrons to have a full outer shell (energy level) of

electrons equivalent to the noble gas of the next highest

atomic number?

(a) 7N (b)

20Ca (c)

50Sn (d)

35Br (e)

34Se

18. Complete the following table so that it contains

information about five different atoms.

Number of Protons

Number of Neutrons

Number of Electrons

Atomic Number

Mass Number

(a) 32 73

(b) 14 14

(c) 28 59

(d) 48 64

(e) 115 77

19. Identify the elements in question 18.

20. What number is most important in identifying an atom?

21. The atomic weight listed in the periodic table for

magnesium is 24.305 amu. Someone said that there

wasn’t a single magnesium atom on the entire earth with

a weight of 24.305 amu. Is this statement correct? Why or

why not?

22. Complete the following table:

IsotopeAtomic

No.MassNo.

No. ofProtons

No. ofNeutrons

No. ofElectrons

Bromine-81 81

Boron-11 535Cl 1752Cr 52

Ni-60

Sr-90

Lead-206

23. The longest wavelength for visible light is approximately

700 nm. What is the frequency for this type of light? Recall:

c 5 ln; c 5 3.00 3 108 m/s.

24. Krypton is the name of Superman’s home planet and also

that of an element. Look up the element krypton and list

its symbol, atomic number, atomic weight, and electron

arrangement.

25. If the frequency of a given type of electromagnetic

radiation increases, the:

(a) speed of light increases

(b) energy decreases

(c) wavelength increases

(d) energy increases



26. Some ultraviolet (UV) radiation has a wavelength of

350 nm. What is this equivalent to according to the

following analysis?

350 nm 31029 m

1nm3

102 cm

1 m5 ______

(a) 350 3 1027 cm (b) 350 3 109 m

(c) 35 cm (d) 350 3 10211 mm

27. Infrared (IR) radiation (wavelengths in the range of

1024 cm) is significantly less energetic than UV radiation.

This means that:

(a) Infrared radiation has a higher frequency than UV

radiation

(b) Ultraviolet radiation has a longer wavelength than IR

radiation

(c) Infrared radiation has a lower frequency than UV

radiation

(d) Infrared radiation has a shorter wavelength than UV

radiation

28. Which equation represents the maximum number of

electrons that can occupy a given energy level (shell) in an

element if n is an integer representing the energy level?

(a) n2 (b) 2n2 (c) 2n (d) n (e) n/2

29. Write the placement of electrons in their ground-state

energy levels according to the Bohr theory for atoms

having 6, 10, 13, and 20 electrons.

30. Write the Bohr electron notation for atoms of the elements

sodium through argon.

31. How many valence electrons do atoms of each of the

elements in question 30 contain?

32. Ultraviolet light has a wavelength of approximately

200 nm. What is the frequency for this form of UV light?

Recall: c 5 ln; c 5 3.00 3 108 m/s.

33. State the periodic law.

34. Give definitions for the following terms:

(a) Group (b) Period

(c) Chemical properties (d) Transition element

(e) Inner transition element

(f) Representative element

35. How did the discovery of the periodic law lead to the

discovery of elements?

36. Describe the relative locations of metals, nonmetals, and

metalloids in the periodic table.

37. How do metals differ from nonmetals?

38. Identify each of the following elements as either a metal,

nonmetal, or metalloid.

(a) Nitrogen (b) Arsenic

(c) Argon (d) Calcium

(e) Uranium

39. Answer the following questions about the periodic table.

(a) How many periods are there?

(b) How many representative groups or families are there?

(c) How many groups consist of all metals?

(d) How many groups consist of all nonmetals?

(e) Is there a period that consists of all metals?

40. What do the electron structures of alkali metals have in

common?

41. The diameter of an atom is approximately 10210 m. The

diameter of an atomic nucleus is about 10215 m. What is the

ratio of the atomic diameter to the nuclear diameter?

42. Give the number of valence electrons for each of the

following:

(a) Ba (b) Al

(c) P (d) Se

(e) Br (f) K

43. Give the symbol for an element that has

(a) 3 valence electrons (b) 4 valence electrons

(c) 7 valence electrons (d) 1 valence electron

44. Draw the Lewis dot symbols for Be, Cl, K, As, and Kr.

45. From their position in the periodic table, predict which will

be more metallic.

(a) Be or B (b) Be or Ca

(c) As or Ge (d) As or Bi

46. Which atom in the following pairs is more metallic?

(a) Li or F (b) Li or Cs

(c) Be or Ba (d) C or Pb

(e) B or Al (f) Na or Ar

47. What general electron arrangement is conducive to

chemical inactivity?

48. Use the information in the periodic table to locate the

following details.

(a) The nuclear charge on cadmium (Cd)

(b) The atomic number of As

(c) The atomic mass (or mass number) of an isotope of Br

having 46 neutrons

(d) The number of electrons in an atom of Ba

(e) The number of protons in an isotope of Zn

(f) The number of protons and neutrons in an isotope of

Sr, atomic mass (or mass number) of 88

(g) An element forming compounds similar to those of Ga

49. Complete the following table.

Atomic Number

Name of Element

Number of Valence

Electrons PeriodMetal or

Nonmetal

6

12

17

37

42

54

50. Write the symbols of the halogen family in the order of

increasing size of their atoms.

51. Why does Cs have larger atoms than Li?

52. How does the atomic radius for a metal atom relate to the

reactivity of the metal?

53. How does the atomic radius for a nonmetal atom relate to

reactivity?



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