Thursday 17/11/2011

The Periodic Table

The first periodic table was devised by Dmitri Mendeleev and published in 1869.

Mendeleev found he could arrange the 65 elements that were then known in a grid

or table so that each element had:

1. A higher atomic weight than the one on its left.

2. Similar chemical properties to other elements in the same column.

He realized that the table in front of him lay at the very heart of chemistry. In his

table he noted gaps - spaces where elements should be but none had yet been

discovered.

In 1913, Henry Moseley, who worked with Rutherford, showed that it is atomic

number (electric charge) which is most fundamental to the chemical properties of

any element. Mendeleev had believed chemical properties were determined by

atomic weight. Moseley correctly predicted the existence of new elements based on

atomic numbers.

Today the chemical elements are still arranged in order of increasing atomic

number (Z) as you go from left to right across the table. We call the horizontal rows

periods and the vertical rows groups.

We also know now that an element's chemistry is determined by the way its

electrons are arranged - its electron configuration.

The noble gases are found in group 18, on the far right of each period. The

reluctance of the noble gases to undergo chemical reactions indicates that the atoms

of these gases strongly prefer their own electron configurations - featuring a full

outer shell of electrons - to any other.

In contrast to the noble gases, the elements with the highest reactivity are those

with the greatest need to gain or lose electrons in order to achieve a full outer shell

of electrons.

Elements that sit in the same group (e.g. the alkali metals in Group 1) all have the

same number of outer electrons, leading to similar chemical properties.

Likewise the halogens in Group 17 also have similar properties to one another.

When halogens react, they gain an electron to form negatively charged ions. Each

ion has the same electron configuration as the noble gas in the same period. The

ions are therefore more chemically stable than the elements from which they

formed.

There is a progression from metals to non-metals across each period.

The block of elements in groups 3 - 12 contains the transition metals. These are

similar to one another in many ways: they produce colored compounds, have

variable valency and are often used as catalysts.

Then we come to the lanthanides (elements 58 - 71) and actinides (elements 90 -

103). The lanthanides are often called the rare earth elements, although in fact these

elements are not rare. The actinides include most of the well-known elements that

take part in or are produced by nuclear reactions. No element with atomic number

higher than 92 occurs naturally in large quantities. Tiny amounts of plutonium and

neptunium exist in nature as decay products of uranium. These elements, and higher

elements, are also produced artificially in nuclear reactors or particle accelerators

The periodic table of the chemical elements (also known as the periodic

table or periodic table of the elements) is a tabular display of the 118

known chemical elements organized by selected properties of their atomic

structures. Elements are presented by increasing atomic number, the number

of protons in an atom's atomic nucleus. While rectangular in general outline, gaps

are included in the horizontal rows (known as periods) as needed to keep elements

with similar properties together in vertical columns (known as groups), e.g. alkali

metals, alkali earths, halogens, noble gases. The following is the periodic table as

defined by the International Union of Pure and Applied Chemistry (IUPAC):

Group # 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

Period

1

1

H

2

He

2

3

Li

4

Be

5

B

6

C

7

N

8

O

9

F

10

Ne

3

11

Na

12

Mg

13

Al

14

Si

15

P

16

S

17

Cl

18

Ar

4

19

K

20

Ca

21

Sc

22

Ti

23

V

24

Cr

25

Mn

26

Fe

27

Co

28

Ni

29

Cu

30

Zn

31

Ga

32

Ge

33

As

34

Se

35

Br

36

Kr

5

37

Rb

38

Sr

39

Y

40

Zr

41

Nb

42

Mo

43

Tc

44

Ru

45

Rh

46

Pd

47

Ag

48

Cd

49

In

50

Sn

51

Sb

52

Te

53

I

54

Xe

6

55

Cs

56

Ba

*

72

Hf

73

Ta

74

W

75

Re

76

Os

77

Ir

78

Pt

79

Au

80

Hg

81

Tl

82

Pb

83

Bi

84

Po

85

At

86

Rn

7

87

Fr

88

Ra

**

104

Rf

105

Db

106

Sg

107

Bh

108

Hs

109

Mt

110

Ds

111

Rg

112

Cn

113

Uut

114

Uuq

115

Uup

116

Uuh

117

Uus

118

Uuo

* Lanthanides

(Lanthanoids)

57

La

58

Ce

59

Pr

60

Nd

61

Pm

62

Sm

63

Eu

64

Gd

65

Tb

66

Dy

67

Ho

68

Er

69

Tm

70

Yb

71

Lu

** Actinides

(Actinoids)

89

Ac

90

Th

91

Pa

92

U

93

Np

94

Pu

95

Am

96

Cm

97

Bk

98

Cf

99

Es

100

Fm

101

Md

102

No

103

Lr

Organizing principles

The main value of the periodic table is the ability to predict the chemical properties

of an element based on its location on the table. It should be noted that the

properties vary differently when moving vertically along the columns of the table

than when moving horizontally along the rows.

The layout of the periodic table demonstrates recurring ("periodic") chemical

properties. Elements are listed in order of increasing atomic number (i.e., the

number of protons in the atomic nucleus). Rows are arranged so that elements with

similar properties fall into the same columns (groups or families). According

to quantum mechanical theories of electron configuration within atoms, each row

(period) in the table corresponded to the filling of a quantum shell of electrons.

There are progressively longer periods further down the table, grouping the

elements into s-, p-, d- and f-blocks to reflect their electron configuration

Atomic number

By definition, each chemical element has a unique atomic number, the number

of protons in its nucleus. Different atoms of many elements have different numbers

of neutrons, which differentiates between isotopes of an element. For example, all

atoms of hydrogen have one proton, and no atoms of any other element have

exactly one proton. On the other hand, a hydrogen atom can have one or two

neutrons in its nucleus, or none at all, yet all of these cases are isotopes of

hydrogen, not instances of some other element. (A hydrogen atom with no neutrons

in addition to its sole proton is called protium, one with one neutron in addition to

its proton is called deuterium, and one with two additional neutrons, tritium.)

In the modern periodic table, the elements are placed progressively in each row

(period) from left to right in the sequence of their atomic numbers, with each new

row starting with the next atomic number following the last number in the previous

row. No gaps or duplications exist. Since the elements can be uniquely sequenced

by atomic number, conventionally from lowest to highest, sets of elements are

sometimes specified by such notation as "through", "beyond", or "from ... through",

as in "through iron", "beyond uranium", or "from lanthanum through lutetium". The

terms "light" and "heavy" are sometimes also used informally to indicate relative

atomic numbers (not densities), as in "lighter than carbon" or "heavier than lead",

although technically the weight or mass of atoms of an element (their atomic

weights or atomic masses) do not always increase monotonically with their atomic

numbers.

The significance of atomic numbers to the organization of the periodic table was

not appreciated until the existence and properties of protons and neutrons became

understood. Mendeleev's periodic tables instead used atomic weights, information

determinable to fair precision in his time, which worked well enough in most cases

to give a powerfully predictive presentation far better than any other comprehensive

portrayal of the chemical elements' properties then possible. Substitution of atomic

numbers, once understood, gave a definitive, integer-based sequence for the

elements, still used today even as new synthetic elements are being produced and

studied.

Periodicity of chemical properties

The primary determinant of an element's chemical properties is its electron

configuration, particularly the valence shell electrons. For instance, any atoms with

four valence electrons occupying p orbitals will exhibit some similarity. The type of

orbital in which the atom's outermost electrons reside determines the "block" to

which it belongs. The number of valence shell electrons determines the family, or

group, to which the element belongs.

Subshell S G F D P

Period

1 1s

2 2s

2p

3 3s

3p

4 4s

3d 4p

5 5s

4d 5p

6 6s

4f 5d 6p

7 7s

5f 6d 7p

8 8s 5g 6f 7d 8p

The total number of electron shells an atom has determines the period to which it

belongs. Each shell is divided into different subshells, which as atomic number

increases are filled in roughly this order (the Aufbau principle) (see table).[5]

Hence

the structure of the periodic table. Since the outermost electrons determine chemical

properties, those with the same number of valence electrons are generally grouped

together.

Progressing through a group from lightest element to heaviest element, the outer-

shell electrons (those most readily accessible for participation in chemical

reactions) are all in the same type of orbital, with a similar shape, but with

increasingly higher energy and average distance from the nucleus. For instance, the

outer-shell (or "valence") electrons of the first group, headed by hydrogen, all have

one electron in an s orbital. In hydrogen, that s orbital is in the lowest possible

energy state of any atom, the first-shell orbital (and represented by hydrogen's

position in the first period of the table). In francium, the heaviest element of the

group, the outer-shell electron is in the seventh-shell orbital, significantly further

out on average from the nucleus than those electrons filling all the shells below it in

energy. As another example, both carbon and lead have four electrons in their outer

shell orbitals.

Note that as atomic number (i.e., charge on the atomic nucleus) increases, this leads

to greater spin-orbit coupling between the nucleus and the electrons, reducing the

validity of the quantum mechanical orbital approximation model, which considers

each atomic orbital as a separate entity.[citation needed]

Groups

A group or family is a vertical column in the periodic table. Groups are considered

the most important method of classifying the elements. In some groups, the

elements have very similar properties and exhibit a clear trend in properties down

the group. Under the international naming system, the groups are numbered

numerically 1 through 18 from the left most column (the alkali metals) to the right

most column (the noble gases).[7]

The older naming systems differed slightly

between Europe and America (the table shown in this section shows the old

American Naming System).[8]

Some of these groups have been given trivial (unsystematic) names, such as

the alkali metals, alkaline earth metals, halogens, pnictogens, chalcogens, and noble

gases. However, some other groups, such as group 7, have no trivial names and are

referred to simply by their group numbers, since they display fewer similarities

and/or vertical trends.[7]

Modern quantum mechanical theories of atomic structure explain group trends by

proposing that elements within the same group generally have the same electron

configurations in their valence shell, which is the most important factor in

accounting for their similar properties.[1]

Elements in the same group show patterns in atomic radius, ionization energy,

and electronegativity. From top to bottom in a group, the atomic radii of the

elements increase. Since there are more filled energy levels, valence electrons are

found farther from the nucleus. From the top, each successive element has a lower

ionization energy because it is easier to remove an electron since the atoms are less

tightly bound. Similarly, a group has a top to bottom decrease in electronegativity

due to an increasing distance between valence electrons and the nucleus.[9]

Periods

A period is a horizontal row in the periodic table. Although groups are the most

common way of classifying elements, there are some regions of the periodic table

where the horizontal trends and similarities in properties are more significant than

vertical group trends. This can be true in the d-block (or "transition metals"), and

especially for the f-block, where the lanthanides and actinides form two substantial

horizontal series of elements.

Elements in the same period show trends in atomic radius, ionization

energy, electron affinity, and electronegativity. Moving left to right across a period,

atomic radius usually decreases. This occurs because each successive element has

an added proton and electron which causes the electron to be drawn closer to the

nucleus.[10]

This decrease in atomic radius also causes the ionization energy to

increase when moving from left to right across a period. The more tightly bound an

element is, the more energy is required to remove an electron. Electronegativity

increases in the same manner as ionization energy because of the pull exerted on

the electrons by the nucleus.[9]

Electron affinity also shows a slight trend across a

period. Metals (left side of a period) generally have a lower electron affinity than

nonmetals (right side of a period) with the exception of the noble gases.[11]

Blocks

Because of the importance of the outermost electron shell, the different regions of

the periodic table are sometimes referred to asperiodic table blocks, named

according to the subshell in which the "last" electron resides. The s-

block comprises the first two groups (alkali metals and alkaline earth metals) as

well as hydrogen and helium. The p-block comprises the last six groups which are

groups 13 through 18 in IUPAC (3A through 8A in American) and contains, among

others, all of the semimetals. The d-block comprises groups 3 through 12 in IUPAC

(or 3B through 8B in American group numbering) and contains all of the transition

metals. The f-block, usually offset below the rest of the periodic table, comprises

the lanthanides and actinides.[12]

Effective nuclear charge

Effective Nuclear Charge Diagram

The effective nuclear charge is the net positive charge experienced by

an electron in a multi-electron atom. The term "effective" is used because

the shielding effect of negatively charged electrons prevents higher orbital electrons

from experiencing the fullnuclear charge by the repelling effect of inner-layer

electrons. The effective nuclear charge experienced by the outer shell electron is

also called the core charge. It is possible to determine the strength of the nuclear

charge by looking at the oxidation number of the atom.

Calculating the effective nuclear charge

In an atom with one electron, that electron experiences the full charge of the

positive nucleus. In this case, the effective nuclear charge can be calculated

from Coulomb's law.

However, in an atom with many electrons the outer electrons are simultaneously

attracted to the positive nucleus and repelled by the negatively charged electrons.

The effective nuclear charge on such an electron is given by the following equation:

Zeff = Z − S

where

Z is the number of protons in the nucleus (atomic number), and

S is the average number of electrons between the nucleus and the electron in

question (the number of nonvalence electrons).

S can be found by the systematic application of various rule sets, the

simplest of which is known as "Slater's rules" (named after John C.

Slater). Douglas Hartree defined the effective Z of a Hartree-Fock orbital

to be:

where

<r>H is the mean radius of the orbital for hydrogen, and

<r>Z is the mean radius of the orbital for an electron configuration with

nuclear charge Z.

Note: Zeff is also often written Z*.

Example

Consider a sodium cation, a fluorine anion, and a

neutral neon atom. Each has 10 electrons, and the number of

nonvalence electrons is 2 (10 total electrons - 8 valence) but the

effective nuclear charge varies because each has a different

atomic number:

Zeff(F-) = 9 − 2 = 7 +

Zeff(Ne) = 10 − 2 = 8 +

Zeff(Na+) = 11 − 2 = 9 +

So, the sodium cation has the largest effective nuclear

charge, and thus the smallest atomic radius.

Values Shielding effect

The shielding effect describes the decrease in attraction between an electron and

the nucleus in any atom with more than one electron shell. It is also referred to as

the screening effect or atomic shielding.

Slater's rules

In quantum chemistry, Slater's rules provide numerical values for the effective

nuclear charge concept. In a many-electron atom, each electron is said to

experience less than the actual nuclear chargeowning to shielding or screening by

the other electrons. For each electron in an atom, Slater's rules provide a value for

the screening constant, denoted by s, S, or σ, which relates the effective and actual

nuclear charges as

The rules were devised semi-empirically by John C. Slater and published in

1930.[1]

Rules

Firstly,[1][4]

the electrons are arranged in to a sequence of groups in order of

increasing principal quantum number n, and for equal n in order of

increasing azimuthal quantum number l, except that s- and p- orbitals are kept

together.

[1s] [2s,2p] [3s,3p] [3d] [4s,4p] [4d] [4f] [5s, 5p] [5d] etc.

Each group is given a different shielding constant which depends upon the

number and types of electrons in those groups preceding it.

The shielding constant for each group is formed as the sum of the following

contributions:

1. An amount of 0.35 from each other electron within the same group

except for the [1s] group, where the other electron contributes only

0.30.

2. If the group is of the [s p] type, an amount of 0.85 from each electron

with principal quantum number (n) one less and an amount of 1.00 for

each electron with an even smaller principal quantum number

3. If the group is of the [d] or [f], type, an amount of 1.00 for each

electron inside it. This includes i) electrons with a smaller principal

quantum number and ii) electrons with an equal principal quantum

number and a smaller azimuthal quantum number (l)

In tabular form, the rules are summarized as:

Group

Other

electrons in

the same

group

Electrons in group(s)

with principal quantum

number n

andazimuthal quantum

number < l

Electrons in

group(s)

with principal

quantum

number n-1

Electrons in all

group(s)

with principal

quantum number <

n-1

[1s] 0.3 N/A N/A N/A

[ns,np] 0.35 N/A 0.85 1

[nd] or [nf] 0.35 1 1 1

Atomic Radius

The atomic radius of an element is half of the distance between the centers of two

atoms of that element that are just touching each other. Generally, the atomic radius

decreases across a period from left to right and increases down a given group. The

atoms with the largest atomic radii are located in Group I and at the bottom of

groups.

Moving from left to right across a period, electrons are added one at a time to the

outer energy shell. Electrons within a shell cannot shield each other from the

attraction to protons. Since the number of protons is also increasing, the effective

nuclear charge increases across a period. This causes the atomic radius to decrease.

Moving down a group in the periodic table, the number of electrons and filled

electron shells increases, but the number of valence electrons remains the same.

The outermost electrons in a group are exposed to the same effective nuclear

charge, but electrons are found farther from the nucleus as the number of filled

energy shells increases. Therefore, the atomic radii increase.

Ionization Energy

The ionization energy, or ionization potential, is the energy required to completely

remove an electron from a gaseous atom or ion. The closer and more tightly bound

an electron is to the nucleus, the more difficult it will be to remove, and the higher

its ionization energy will be. The first ionization energy is the energy required to

remove one electron from the parent atom. The second ionization energy is the

energy required to remove a second valence electron from the univalent ion to form

the divalent ion, and so on. Successive ionization energies increase. The second

ionization energy is always greater than the first ionization energy. Ionization

energies increase moving from left to right across a period (decreasing atomic

radius). Ionization energy decreases moving down a group (increasing atomic

radius). Group I elements have low ionization energies because the loss of an

electron forms a stable octet.

The property is alternately still often called the ionization potential, measured in

volts. In chemistry it is often referred to one mole of substance (molar ionization

energy or enthalpy) and reported in kJ/mol. In atomic physics the ionization energy

is typically measured in the unit electron volt (eV).

The ionization energy is different for electrons of different atomic or molecular

orbitals. More generally, the nth ionization energy is the energy required to strip off

the nth electron after the first n − 1 electrons have been removed. It is considered a

measure of the tendency of an atom or ion to surrender an electron, or the strength

of the electron binding; the greater the ionization energy, the more difficult it is to

remove an electron. The ionization energy may be an indicator of the reactivity of

an element. Elements with a low ionization energy tend to be reducing agents and

form cations, which in turn combine with anions to form salts.

Electron binding energy (BE), more accurately, is the energy required to release

an electron from its atomic or molecular orbital when adsorbed to a surface rather

than a free atom. Binding energy values are normally reported as positive values

with units of eV. The binding energies of 1s electrons are roughly proportional to

(Z-1)² (Moseley's law).

Values and trends

Main article: Molar ionization energies of the elements

Generally the (n+1)th ionization energy is larger than the nth ionization energy.

Always, the next ionization energy involves removing an electron from an

orbital closer to the nucleus. Electrons in the closer orbitals experience greater

forces of electrostatic attraction; thus, their removal requires increasingly more

energy.

Some values for elements of the third period are given in the following table:

Successive molar ionization energies in kJ/mol

(96.485 kJ/mol = 1 eV/particle)

Element First Second Third Fourth Fifth Sixth Seventh

Na 496 4,560

Mg 738 1,450 7,730

Al 577 1,816 2,881 11,600

Si 786 1,577 3,228 4,354 16,100

P 1,060 1,890 2,905 4,950 6,270 21,200

S 999.6 2,260 3,375 4,565 6,950 8,490 27,107

Cl 1,256 2,295 3,850 5,160 6,560 9,360 11,000

Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000

Large jumps in the successive molar ionization energies occur when

passing noble gas configurations. For example, as can be seen in the table above,

the first two molar ionization energies of magnesium (stripping the two 3s

electrons from a magnesium atom) are much smaller than the third, which

requires stripping off a 2p electron from the very stable neon configuration of

Mg2+

.

Periodic trend for ionization energy. Each period begins at a minimum for the

alkali metals, and ends at a maximum for the noble gases.

Ionization energy is also a periodic trend within the periodic table organization.

Moving left to right within aperiod or upward within a group, the first ionization

energy generally increases. As the atomic radiusdecreases, it becomes harder to

remove an electron that is closer to a more positively charged nucleus. ionization

enthalpy increases from left to right in a period and decreases from top to bottom

in a group.

Electron Affinity

Electron affinity reflects the ability of an atom to accept an electron. It is the energy

change that occurs when an electron is added to a gaseous atom. Atoms with

stronger effective nuclear charge have greater electron affinity. Some

generalizations can be made about the electron affinities of certain groups in the

periodic table. The Group IIA elements, the alkaline earths, have low electron

affinity values. These elements are relatively stable because they have

filled s subshells. Group VIIA elements, the halogens, have high electron affinities

because the addition of an electron to an atom results in a completely filled shell.

Group VIII elements, noble gases, have electron affinities near zero, since each

atom possesses a stable octet and will not accept an electron readily. Elements of

other groups have low electron affinities.

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