Chemistry part 3, Julia Burdge,2e (2009) pptx

Alpha a rays consist of positively charged particles, called a particles, that are deflected away from f3 ray ~ s /, IX rays / Photographic plate negatively charged particles e

2.1 The Atomic Theory

2.2 The Structure of the

Atom

• Discovery of the Electron

• Radioactivity

• The Proton and the Nucleus

• Nuclear Model of the Atom

• The Neutron

2.3 Atomic Number, Mass

Number, and Isotopes

2.4 The Periodic Table

2.5 The Atomic Mass Scale

and Average Atomic Mass

2.6 Molecules and Molecular

An inadequate supply of iron and the resulting shortage of hemoglobin can cause iron

deficiency anemia (IDA) Some of the symptoms of IDA are fatigue, weakness, pale color, poor appetite, headache, and light-headedness

•

Although IDA can be caused by loss of blood or by poor absorption of iron, the most common cause is insufficient iron in the diet Dietary iron comes from such sources as meat, eggs, leafy green vegetables, dried beans, and dried fruits Some breakfast cereals,

such as Total, are fOliified with iron in the form of iron metal, also known as elemental

or reduced iron The absorption of dietary iron can be enhanced by the intake of vitamin

C (ascorbic acid) When the diet fails to provide enough iron, a nutlitional supplement

may be necessary to prevent a deficiency Many supplements provide iron in the form of

a compound called ferrous sulfate

Elemental iron, ascorbic acid, and the iron in ferrous sulfate are examples of some of the atoms, molecules, and ions that are essential for human health

Iron absorption can be dim i nished by certain

d iso rders, such as Crohn's disease, and by some

Before you begin, you should review

• Significant figures [~ ~ Section 1.5]

• Dimensional analysis [~ ~ Section 1.6]

Meat, eggs, and some vegetables and cereals are good sources

of dietary iron When diet alone does not provide an adequate

supply, iron supplements can be taken

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33

34 CHAPTER 2 Atoms, Molecules, and Ions

Figure 2.1 (a) John Dalt on

(b) Dalton's sys tem of sy mbol s

The Atomic Theory

In the fifth century B.C the Greek philosopher Democritus proposed that all matter consists of very

small, indivisible particles, which he named atomos (meaning uncuttable or indivisible) Although Democritus' idea was not accepted by many of his contemporaries (notably Plato and Aristotle),

somehow it endured Experimental evidence from early scientific investigations provided support for the notion of "atomism" and gradually gave rise to the modern definitions of elements and compounds In 1808, an English scientist and school teacher, John Dalton (Figure 2.1), formulated

a precise definition of the indivisible building blocks of matter that we call atoms

Dalton's work marked the beginning of the modern era of chemistry The hypotheses about

the nature of matter on which Dalton's atomic theory is based can be summarized as follows:

1 Elements are composed of extremely small particles called atoms All atoms of a given ele

-ment are identical, having the same size, mass, and chemical properties The atoms of one element are different from the atoms of all other elements

2 Compounds are composed of atoms of more than one element In any given compound, the

same types of atoms are always present in the same relative numbers

3 A chemical reaction rearranges atoms in chemical compounds; it does not create or destroy

Figure 2.2 This represents a c hemi ca l reaction between the elements

oxygen and carbon (a) Oxygen does not exist as i so lated atoms under

ordinary conditions, but rather as molecules, each of which consists of two

oxygen atoms Note that the oxygen atoms ( red sphe re s) appear all to be

identical to one another ( hypothe sis 1 ) (b) Likewise, the carbon atoms ( black

spheres) all appear to be identical to one another Car bon also exists in the

form of molecule s that are more varied a nd complex than those of oxyge n

T h e carbo n ha s been represented as isolated atoms to simplify the figure (a)

+

(c) The compound CO2 forms when each carbon atom combines with two

oxygen atoms ( hypothe s i s 2) Finally, th e reaction results in the rearrangement

of the atoms, but a ll the atoms present before the reaction (lef t of the arrow)

are also present after the reaction ( right of the arrow) (hypothesis 3)

•

Who Was John Dalton?

John Dalton (1766-1844) was an English chemist,

mathemati-cian, and philosopher [Figure 2.l(a)] In addition to the atomic

theory, he also formulated several gas laws and gave the first

detailed description of the type of color blindness from which he

suffered This type of color blindness is known today as

"Dalton-ism." He was described by his friends as awkward and without

social grace, and in fact, he spent much of his time studying and

tutoring students

In addition to his other work, Dalton devised a system of symbols [Figure 2.1(b)], which he used to represent the elements

He also assigned relative weights to the elements, many of which

differ considerably from those that we use today

Dalton's only recreation is said to have been a nightly pint of ale and lawn bowling every Thursday afternoon Note

how closely the game balls resemble modern-day molecular

models!

Dalton's concept of an atom was far more detailed and specific than that of Democritus The

first hypothesis states that atoms of one element are different from atoms of all other elements

Dalton made no attempt to describe the structure or composition of atoms he had no idea what

an atom was really like He did realize, though, that the different properties shown by elements

uch as hydrogen and oxygen could be explained by assuming that hydrogen atoms were not the

ame as oxygen atoms

The second hypothesis suggests that, in order to form a certain compound, we not only need

atoms of the right kinds of elements, but specific numbers of these atoms as well This idea is an

extension of a law published in 1799 by Joseph Proust, a French chemist According to Proust's law

of definite proportions, different samples of a given compound always contain the same elements in

the same mass ratio Thus, if we were to analyze samples of carbon dioxide gas obtained from

differ-ent sources, such as the exhaust from a car in Mexico City or the air above a pine forest in northern

Y1aine, each sample would contain the same ratio by mass of oxygen to carbon Consider the

follow-ing results of the analysis of three samples of carbon dioxide, each from a different source:

Sample

123 g carbon dioxide 50.5 g carbon dioxide

88.6 g carbon dioxide

Mass of 0 (g)

89.4 36.7 64.4

In any sample of pure carbon dioxide, there are 2.66 g of oxygen for every gram of carbon present

This constant mass ratio can be explained by assuming that the elements exist in tiny particles of

fixed mass (atoms), and that compounds are formed by the combination of fixed numbers of each

type of particle

Dalton's second hypothesis also supports the law of multiple proportions According to this

law, if two elements can combine to form more than one compound with each other, the masses of

one element that combine with a fixed mass of the other element are in ratios of small whole num

-bers That is, different compounds made up of the same elements differ in the number of atoms of

each kind that combine For example, carbon combines with oxygen to form carbon dioxide and

arbon monoxide In any sample of pure carbon monoxide, there are 1.33 g of oxygen for every

36 CHAPTER 2 Atoms, Molecules, and Ions

Figure 2.3 An illustration of the

law of multiple proportions

ratio of 0 to C in carbon dioxide = 2.66 = 2: 1 ratio of 0 to C in carbon monoxide 1.33

For samples containing equal masses of carbon, the ratio of oxygen in carbon dioxide to oxygen

in carbon monoxide is 2: l Modern measurement techniques indicate that one atom of carbon combines with two atoms of oxygen in carbon dioxide and with one atom of oxygen in carbon monoxide This result is consistent with the law of multiple proportions (Figure 2.3)

Dalton's third hypothesis is another way of stating the law of conservation of mass, 1 which

is that matter can be neither created nor destroyed Because matter is made up of atoms that are unchanged in a chemical reaction, it follows that mass must be conserved as well Dalton's bril-liant insight into the nature of matter was the main stimulus for the rapid progress of chemistry during the nineteenth century

On the basis of Dalton's atomic theory, we can define an atom as the basic unit of an element that can enter into chemical combination Dalton imagined an atom that was both extremely small and indivisible However, a series of investigations that began in the 1850s and extended into the twentieth century clearly demonstrated that atoms actually possess internal structure; that is, they are made up of even smaller particles, which are called subatomic particles This research led to the discovery of electrons, protons, and neutrons

Discovery of the Electron

Many scientists in the l890s studied radiation, the emission and transmission of energy through

space in the form of waves Information gained from this research contributed greatly to our

under-standing of atomic structure One device used to investigate this phenomenon was a cathode ray tube, the forerunner of the tubes used in older televi sions and computer monitors (Figure 2.4)

A cathode ray tube consists of two metal plates sealed inside a glass tube from which most

of the air has been evacuated When the metal plates are connected to a high-voltage source, the negatively charged plate, called the cathode, emits an invisible ray The cathode ray is drawn to the positively charged plate, called the anode, where it passes through a hole and continues trav-eling to the other end of the tube When the ray strikes the specially phosphor-coated surface, it produces a bright light

Because consistent results are observed regardless of the composition of the cathode, ode rays were presumed to be a component of all matter Furthermore, because the path of the cathode rays could be deflected by magnetic and electric fields, as shown in Figure 2.4, they must

cath-I According to A lb ert Einstein, mass and energy are alternate aspects of a single ent it y called ma ss-e n ergy Chemical reac tion s u s uall y involve a gain or l oss of h eat and other forms of e n ergy Thus, whe n e ner gy is lost in a reaction , for example,

-ma ss is also lost Except for nuclear reaction s (see Chapter 20), however, changes of ma ss in chemical r eact ion s are far too

s mall to detect Therefore, for all pra ct i cal purposes ma ss is conserved

SECTION 2.2 The Structure of the Atom 37

behaves like a magnet and can interact with electric and magnetic fields through which it passes

bearing negative charges, it must consist of negatively charged particles We know these

picture to become distorted temporarily when a magnet is brought close to the screen

-tromagnetic theory to determine the ratio of electric charge to the mass of an individual electron

The number he calculated was 1.76 X 108 C/g, where C stands for coulomb, which is the derived

SI unit of electric charge (In SI base units, 1 C = 1 A s.) Thereafter, in a series of experiments

arried out between 1908 and 1917, American R A Millikan succeeded in measuring the charge

tiny drops of oil that picked up static charge from particles in the air He suspended the charged

drops in air by applying an electric field and followed their motions through a microscope (Figure

_.6) His work proved that the charge on each electron was exactly the same: - 1.6022 X 10-19 C "

calcu-lated the mass of an electron as follows:

Figure 2.4 A cathode ray tube with

an electric field perpendicular to the

direction of the cathode rays and an

exte rnal magnetic field The sy mbol s

Nand S denote the north and so uth

poles of the magnet The cathode rays will strike the end of the tube at point

A in the presence of a magnetic field and at point B in the pre se nce of an electric field (In the absence of any external field or when the effects of the e l ectric field and magnetic field

cancel each other-the cathode ray s

will not be deflected but will travel in a

s traight line and strike the middle of the

circular screen.)

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Cathode ray tube experiment

and second, respectively [ ~ S ec t io n 1 3,

T able 1 2]

droplets with multiple charges attached Th e

charges he determined were alway s multiples

of - 1.6022 x 10 - 19 C

(c)

Figure 2.5 (a) Although cathode ray s them se lve s are invisible, they cause a phosphorescent screen to glow when the y s trike it-making it possible

f or us to see their path (b), (c) Cathode rays are deflected by a magnet, indic at in g that they are electricall y c harged

38 CHAPTER 2 Atoms, Molecules, and Ions

Millikan's oil-drop experiment

In 1895, the German physicist Wilhelm Rontgen noticed that cathode rays caused glass and

met-als to emit yet another type of ray This highly energetic radiation penetrated matter, darkened

covered photographic plates, and caused a variety of substances to fluoresce (give off light) These rays were not deflected by a magnet, however, so unlike cathode rays, they could not contain charged particles Rontgen called them X rays because of their mysterious nature

Not long after Rontgen's discovery, Antoine Becquerel, a professor of physics in Paris, began to study the fluorescent properties of substances Purely by accident, he found that exposing thickly wrapped photographic plates (so no light could get in) to a uranium compound caused them

to darken, even without the stimulation of cathode rays Like X rays, the rays from the uranium

compound were highly energetic and could not be deflected by a magnet, but they differed from X

rays because they arose spontaneously One of Becquerel's students, Marie Curie, suggested the name radioactivity to describe this spontaneous emission of particles and/or radiation Today, we use the term radioactive to describe any element that spontaneously emits radiation

Three types of rays are produced by the breakdown, or decay, of radioactive substances such

as uranium Two of the three are deflected by oppositely charged metal plates (Figure 2.7) Alpha

(a) rays consist of positively charged particles, called a particles, that are deflected away from

f3 ray ~ s /,

IX rays

/

Photographic plate

negatively charged particles (electrons) and are therefore attracted by the positively charged plate The

opposite holds true for so-called Q' rays-they are actually positively charged particles and are drawn to the negatively charged plate Because y rays are not particles and have no charge, their path is unaffected by an external electric field

SECTION 2.2 The Structure of the Atom 39

the positively charged plate Beta (13) rays, or 13 particles, are electrons, so they are deflected away

from the negatively charged plate The third type of radioactive radiation consists of high-energy

gamma ('}') rays Like X rays, '}' rays have no charge and are unaffected by external electric or

magnetic fields

The Proton and the Nucleus

By the early 1900s, scientists knew that atoms contained electrons and that they were

electri-cally neutral In order to be neutral, an atom must contain equal amounts of positive and negative

charge Thomson proposed, therefore, that an atom could be thought of as a sphere of positively

charged matter in which negatively charged electrons were embedded uniformly, like the

choco-late chips in a scoop of mint chocolate chip ice cream This so-called plum-pudding model was the

accepted theory for a number of years

In 1910 the New Zealand physicist Ernest Ruthelford, who had studied with Thomson at

Cambridge University, decided to use a particles to probe the structure of atoms Together with

his associate Hans Geiger and an undergraduate named Ernest Marsden, Rutherford carried out

a series of experiments using very thin foils of gold and other metals as targets for a particles

from a radioactive source They observed that the majority of particles penetrated the foil either

ompletely un deflected or with only a small angle of deflection Every now and then, however, an

a particle was scattered (or deflected) at a large angle In some instances, the a particle actually

bounced back in the direction from which it had come! This was an extraordinarily surprising

find-ing In Thomson's model the positive charge of the atom was so diffuse that the relatively massive,

positively charged particles should all have passed through the foil with little or no deflection To

quote Rutherford's initial reaction when told of this discovery: "It was as incredible as if you had

fired a IS-inch shell at a piece of tissue paper and it came back and hit you." Figure 2.8 illustrates

the results of Rutherford's a-scattering experiment

the s cattering of a particles by a piece of gold foil The plum - pudding

model predicted that the a particle s

would all pa ss through the gold foil

go ld foil with little or no deflection ,

but a few are deflected at lar ge angles

the fo il back toward the so urce The

nucl ear model e x plain s the re s ult s of

Ruth erfo rd 's experiments

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Rutherford's experiment

40 CHAPTER 2 Atoms, Molecules, and Ions

Atomic radii are sometimes given in angstroms,

where 1 angstrom (A) = 1 X 10 - 10 m Using

angstroms, a typical atomic radius is about

1 A and a typic'l l nuclear radius is about

5 x 10 -5 A

Nuclear Model of the Atom

Rutherford later explained the results of the a-scattering experiment by proposing a new model

for the atom According to Ruthelford, most of the atom must be empty space This would explain

why the majority of a particles passed through the gold foil with little or no deflection The atom's positive charges, Rutherford proposed, were all concentrated in the nucleus, which is an extremely

dense central core within the atom Whenever an a particle came close to a nucleus in the scatter

-ing experiment, it experienced a large repulsive force and therefore a large deflection Moreover,

an a particle traveling directly toward a nucleus would be completely repelled and its direction

would be reversed

The positively charged particles in the nucleus are called protons In separate experiments,

it was found that each proton carried the same quantity of charge as an electron (just opposite in

sign) but had a mass of 1.67262 X 10 - 24 g Although this is an extremely small value, it is nearly

2000 times the mass of an electron

Based on these data, the atom was believed to consist of a nucleus that accounted for most

of the mass of the atom, but which occupied only a tiny fraction of its volume We express atomic

(and molecular) dimensions using the SI unit picometer (pm), where

1 pm = 1 X 1O -J? m

A 'tYPICai atoinic '[adiu's ' ls ' about"1 00 pm, whereas the radius of an atomic nucleus is only about

5 X 10 - 3 pm You can appreciate the relative sizes of an atom and its nucleus by imagining that if

an atom were the size of the New Orleans Superdome, the volume of its nucleus would be rable to that of a marble While the protons are confined to the nucleus of the atom, the electrons are distributed around the nucleus at relatively large distances from it

compa-The concept of atomic radius is useful experimentally, but you should not get the

impres-sion that atoms have well-defined boundaries or surfaces We will learn in Chapter 6 that the outer regions of atoms are relatively "fuzzy," not sharply defined

The Neutron

Rutherford's model of atomic structure left one major problem unsolved It was known that gen, the simplest atom, contained only one proton and that the helium atom contained two pro-tons Therefore, the ratio of the mass of a helium atom to that of a hydrogen atom should be 2: 1 (Because electrons are much lighter than protons, their contribution to atomic mass can be ignored.) In reality, however, the ratio is 4: 1 Rutherford and others postulated that there must

hydro-be another type of subatomic particle in the atomic nucleus, the proof of which was provided by James Chadwick, an English physicist, in 1932 When Chadwick bombarded a thin sheet of beryl-lium with a particles, a very high-energy radiation was emitted by the metal that was not deflected

by either electric or magnetic fields Although similar to 'Y rays, later experiments showed that the rays actually consisted of a third type of subatomic particle, which Chadwick named neutrons,

because they were electrically neutral particles having a mass slightly greater than that of protons The mystery of the mass ratio could now be explained A typical helium nucleus consists of two protons and two neutrons, whereas a typical hydrogen nucleus contains only a proton; the mass ratio, therefore, is 4: 1

Figure 2.9 shows the location of the elementary particles (protons, neutrons, and electrons)

in an atom There are other subatomic particles, but the electron, the proton, and the neutron are the three fundamental components of the atom that are important in chemistry Table 2.1 lists the masses and charges of these three elementary particles

Atomic Number, Mass Number, and Isotopes

All atoms can be identified by the number of protons and neutrons they contain The atomic ber (Z) is the number of protons in the nucleus of each atom of an element It also indicates the number of electrons in the atom because atoms are neutral and contain the same number of protons and electrons The chemical identity of an atom can be determined solely from its atomic number For example, the atomic number of nitrogen is 7 Thus, each nitrogen atom has seven pro-

num-tons and seven electrons Or, viewed another way, every atom in the universe that contains seven protons is a nitrogen atom

The mass number (A) is the total number of neutrons and protons present in the nucleus of

an atom of an element Except for the most common form of hydrogen, which has one proton and

SECTION 2.3 Atomic Number, Mass Number, and Isotopes 41

mass number (A) = number of protons (Z) + number of neutrons

The number of neutrons in an atom equals the difference between the mass number and the atomic

number, or (A - Z) For example, the mass number of fluorine is 19 and the atomic number is

9 (indicating 9 protons in the nucleus) Thus, the number of neutrons in an atom of fluorine is

19 - 9 = 10 The atomic number, number of neuu'ons, and mass number all must be positive

integers (whole numbers)

The accepted way to denote the atomic number and mass number of an atom of an element

•

Contrary to the first hypothesis of Dalton's atomic theory, atoms of a given element do not all have

the same mass Most elements have two or more isotopes, atoms that have the same atomic

num-ber (Z) but different mass numbers (A) For example, there are three isotopes of hydrogen, called

hy drogen (or protium ), deuterium , and tritium Hydrogen has one proton and no neutrons in its

nucleus, deuterium has one proton and one neutron, and tritium has one proton and two neutrons

Thus, to represent the isotopes of hydrogen, we write

Figure 2.9 The protons and

Electrons are distributed within the

sp her e surrou nding the nucleu s

Collecti vely, protons and neutrons are called

nucleons A nucleon is a particle within the nucleus

Because these symbols de Sig nate isotopes

by specifying nu mbers of nucleons, they are sometimes referred to as nuclear symbol s

42 CHAPTER 2 Atoms, Molecules, and Ions

Think About It Verify that the

number of proton s and the number

of neutrons for each example s um

to the mass number that is gi v en

In part (a), for examp l e, there are

17 protons and 18 neutron s , which

sum to give a mass number of 35 ,

the value given in the problem In

part (b), 17 protons + 20 neutrons

= 37 In part (c ) , 19 protons +

22 neutrons = 41 In part (d ) , 6

protons + 8 neutrons = 14

Similarly, the two common isotopes of uranium (Z = 92), which have mass numbers of 235 and

The first isotope, with 235 - 92 = 143 neutrons in its nucleus, is used in nuclear reactors and atomic

appli-cations With the exception of hydrogen, which has different names for each of its isotopes, the topes of other elements are identified by their mass numbers The two isotopes of uranium are called

3H and 235U are sufficient to specify the isotopes tritium and uranium-235, respectively

The chemical properties of an element are detelmined primarily by the protons and electrons

isotopes of the same element exhibit similar chemical properties, forming the same types of com

-pounds and displaying similar reactivities

Sample Problem 2.1 shows how to calculate the number of protons, neutrons, and electrons

of protons

Setup Number of proton s = Z , number of neutron s = A - Z, and number of electrons = number of protons Recall that the 14 in carbon-14 i s the mas s number

Solution

( a) The atomic number is 17 , s o there are 17 protons The mass number is 35, so the number of neutrons

is 35 - 17 = 1 8 The number of electron s equals the number of proton s , so there are 17 electrons

( b) Again, the atomic number i s 17 , s o there are 17 protons The ma s s number is 37, so the number

of neutron s is 37 - 17 = 20 The number of e l ectron s equal s the number of protons, so there are

17 electrons , too

(c) The atomic number of K ( potas s ium ) is 19, s o there are 19 protons The mass number i s 41, so there are 41 - 19 = 2 2 neutron s There are 19 electrons

(d) Carbon-14 can al s o be represented a s 14 c The atomic number of carbon is 6, so there are

6 proton s and 6 electron s There are 14 - 6 = 8 neutrons

Practice Problem A How many protons , neutrons , and electrons are there in an atom of (a) 19B, (b ) 36 Ar , ( c) ~~ Sr , and ( d ) carbon-II ?

Practice Problem B Give the correct s y mbols to id entify an atom that con t a in s (a) 4 protons,

4 e l ectron s , and 5 neutron s; ( b) 23 proton s , 23 e l ectrons, and 28 neutrons; (c) 54 protons,

54 electron s, and 70 neutron s; and (d ) 31 protons , 31 electrons, and 38 neutrons

2.3.1 Ho w many neutron s are there in an 2.3.2 What is the mass number of an oxygen

atom of 6o Ni? atom with nine neutrons in its nucleus?

How Are Atomic Masses Measured?

The most direct and most accurate method for determining atomic

and molecular masses is mass spectrometry In a mass

spectrome-ter, such as that depicted in Figure 2.10, a gaseous sample is

bom-barded by a stream of high-energy electrons Collisions between

the electrons and the gaseous atoms (or molecules) produce

posi-tively charged species, called ions, by dislodging an electron from

the atoms or molecules These positive ions (of mass m and charge

e) are accelerated as they pass through two oppositely charged

plates The emerging ions are deflected into a circular path by a

magnet The radius of the path depends on the charge-to-mass

ratio (i.e., elm) Ions with a small elm ratio trace a wider arc than

those having a larger elm ratio, so ions with equal charges but

dif-ferent masses are separated from one another The mass of each

ion (and hence its parent atom or molecule) is determined from the

magnitude of its deflection Eventually the ions arrive at the

detec-Electron beam

Sample gas

tor, which registers a current for each type of ion The amount of

of isotopes, such as neon-20 (natural abundance 90.48 percent)

and neon-22 (natural abundance 9.25 percent) When more

exper-imental accuracy is to a quantitative science like chemistry Early

abun-dance was so small Only 27 in 10,000 Ne atoms are neon-21

Magnet

Detecting screen

-v olume of a va ilable in fo rma t i o n about the s tructure and pr o perti es of elem e ntal s u b s tance s , led

t o the d eve lopment o f th e periodi c tabl e, a c hart in whi c h e lem e nt s ha v in g s imil a r ch e mi c al and

ph ys ical prop e rtie s are g roup e d t oge th e r Fi g ur e 2 1 2 s how s th e mod e rn p e riodic table in which the e lem e nt s a r e an a nged b y a t o mi c number (s hown ab o ve the el e ment s ymb o l ) in horizont a l row s

c all e d period s a nd in ve rtical co l umn s c alled group s orfamilies Elem e nt s in th e s am e g roup tend

to h av e s imilar ph ys ical a nd c h e mi c a l p ro perti es

Th e el e ment s c an be ca teg o ri z e d a s m e tal s, n o nmet a l s, o r met a lloid s A metal i s a go o d

co ndu c t o r o f h e at a nd e l ec tri ci t y, where as a nonm e tal i s u s u a ll y a poor conductor o f he a t and elec

-tricit y A metalloid ha s prop e ltie s that ar e interm e diate between tho se of metal s and nonme t al s

F i g ur e 2 1 2 s h ows that th e m aj orit y of kno w n e l e ment s a re metal s; on l y 17 e lem e nt s are n o

nmet-a l s, a nd fewe r than 10 e lement s ar e m e tall o ids A lthough m o st s o ur ce s, incl u ding thi s text , de s

-ign a t e th e e lement s B , Si , G e, A s, S b , a nd T e a s m e tall o id s, s our ces var y f o r t h e e lem e nt s Po and

A t In thi s te x t , we cla ss i fy b o th P o and A t as met a lloid s Fr o m l e ft to right a cro ss a n y peri o d , th e

phy s ical a nd c h e mi ca l pr o p e r t i es of the e lem e nt s c han ge gradua ll y from m e talli c to n o nm e ta ll ic

El e ment s ar e o ften r e f e n'ed t o c o ll ec ti ve l y b y th e ir period i c table g ro u p number ( Group lA ,

Gr o up 2 A, an d s o o n ) Fo r co n v enienc e , h ow e ver , s om e element group s h a v e b ee n given s pecial

-Figure 2.12 Th e mod e m p e ri o di c tab l e T h e e lem e nt s ar e a rr a n g ed acc ordin g to a tomi c number , w hi c h i s s h ow n a b ove ea ch e l em e nt 's sy mb o l

With th e exc epti o n of h y d rog e n ( H ), n o nm eta l s a p pe ar at th e fa r ri g ht o f the t a bl e Th e t wo ro ws of me tal s bene a th th e main b o d y o f th e tab l e are se t

apart t o k ee p the tabl e from be i n g to o wi d e Ac t u a ll y, l a nthanum ( 5 7) sh o u l d fo ll ow barium (5 6), and act i nium (8 9 ) sh o u l d fo ll ow radium ( 8 8) Th e 1-1 8

g roup d es i g nati o n h as be e n r e c o mm e nd e d b y th e Int e rn a tion a l U ni o n o f Pure a nd A ppli e d C h e mi s tr y ( IUPA C) but i s n o t ye t in wid e u se In thi s te x t , w e

ge n e rall y u s e the s t a n d ard U S n o tati o n f or g rou p numb e r s (lA-8A and lB- 8B) No nam es h ave ye t been ass i g n e d t o elem e nt s 112-11 6 and 118 As of

th i s w ritin g, t h e sy nth es i s of e l e m e nt 11 7 h as n o t ye t b ee n r epo rt e d