Periodic table
From Wikipedia, the free encyclopedia
The periodic table of the chemical elements (also periodic table of the elements or just the periodic table) is a tabular display of
the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in
1869, who intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been
refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain
chemical behavior.[1]
The periodic table is now ubiquitous within the academic discipline of chemistry, providing a useful framework to classify, systematize,
and compare all of the many different forms of chemical behavior. The table has found many applications in chemistry, physics, biology,
and engineering, especially chemical engineering. The current standard table contains 118 elements to date. (elements 1–118).
Contents
1 Structure
2 Classification
2.1 Groups
2.2 Periods
2.3 Blocks
2.4 Other
3 Periodicity of chemical properties
3.1 Trends of groups
3.2 Trends of periods
4 History
5 Gallery
6 See also
7 Notes
8 References
9 Further reading
10 External links
Structure
Group # 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Period
1 1H
2
He
2 3Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
3 11Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
4 19K
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 37Rb
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 55Cs
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 87Fr
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
* Lanthanoids 57La
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
** Actinoids 89Ac
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
This common arrangement of the periodic table separates the lanthanoids and actinoids (the f-block) from other elements. The wide
periodic table incorporates the f-block. The extended periodic table adds the 8th and 9th periods, incorporating the f-block and adding
the theoretical g-block.
Element categories in the periodic table
Metals Metalloids Nonmetals Unknown
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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
chemical
properties
Alkali
metals
Alkaline earth
metals
Inner transition
elements Transition
elements
Other
metals
Other
nonmetals Halogens
Noble
gases
Lanthanides Actinides
Atomic number colors show state at standard temperature and
pressure (0 °C and 1 atm)
Solids Liquids Gases Unknown
Borders show natural occurrence
Primordial From decay Synthetic (Undiscovered)
Other alternative periodic tables exist.
Some versions of the table show a dark stair-step line along the metalloids. Metals are to the left of the line and non-metals to the
right.[2]
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.
In printed tables, each element is usually listed with its element symbol and atomic number; many versions of the table also list the
element's atomic mass and other information, such as its abbreviated electron configuration, electronegativity and most common valence
numbers.
As of 2010, the table contains 118 chemical elements whose discoveries have been confirmed. Ninety-four are found naturally on Earth,
and the rest are synthetic elements that have been produced artificially in particle accelerators. Elements 43 (technetium), 61
(promethium) and all elements greater than 83 (bismuth), beginning with 84 (polonium) have no stable isotopes. The atomic mass of
each of these element's isotope having the longest half-life is typically reported on periodic tables with parentheses.[3] Isotopes of
elements 43, 61, 93 (neptunium) and 94 (plutonium), first discovered synthetically, have since been discovered in trace amounts on
Earth as products of natural radioactive decay processes.
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.
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). Hence the structure of the table. Since the outermost electrons determine
chemical properties, those with the same number of valence electrons are 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.
The elements ununtrium, ununquadium, ununpentium, etc. are elements that have been discovered, but so far have not received a trivial
name yet. There is a system for naming them temporarily.
Classification
Groups
Main article: Group (periodic table)
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. These groups tend to
be given trivial (unsystematic) names, e.g., the alkali metals, alkaline earth metals, halogens, pnictogens, chalcogens, and noble gases.
Some other groups in the periodic table display fewer similarities and/or vertical trends (for example Group 14), and these have no trivial
names and are referred to simply by their group numbers.
Periods
Main article: Period (periodic table)
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This diagram shows the periodic table blocks.
Periodic trend for ionization energy. Each period
begins at a minimum for the alkali metals, and ends
at a maximum for the noble gases.
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.
Blocks
Main article: Periodic table block
Because of the importance of the outermost shell, the different regions of the
periodic table are sometimes referred to as periodic 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 (groups 13 through 18) and
contains, among others, all of the semimetals. The d-block comprises groups 3
through 12 and contains all of the transition metals. The f-block, usually offset
below the rest of the periodic table, comprises the rare earth metals.
Other
The chemical elements are also grouped together in other ways. Some of these
groupings are often illustrated on the periodic table, such as transition metals, poor
metals, and metalloids. Other informal groupings exist, such as the platinum group and the noble metals.
Periodicity of chemical properties
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.
Trends of groups
Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group have
the same electron configurations in their valence shell, which is the most important factor in accounting for their similar properties.
Elements in the same group also show patterns in their 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 will also see a top to bottom decrease in electronegativity due to an increasing distance
between valence electrons and the nucleus.
Trends of periods
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. 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. Similarly,
electronegativity will increase in the same manner as ionization energy because of
the amount of pull that is exerted on the electrons by the nucleus. 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.
History
Main article: History of the periodic table
In 1789, Antoine Lavoisier published a list of 33 chemical elements. Although Lavoisier grouped the elements into gases, metals,
non-metals, and earths, chemists spent the following century searching for a more precise classification scheme. In 1829, Johann
Wolfgang Döbereiner observed that many of the elements could be grouped into triads (groups of three) based on their chemical
properties. Lithium, sodium, and potassium, for example, were grouped together as being soft, reactive metals. Döbereiner also observed
that, when arranged by atomic weight, the second member of each triad was roughly the average of the first and the third.[4] This
became known as the Law of triads.[citation needed] German chemist Leopold Gmelin worked with this system, and by 1843 he had
identified ten triads, three groups of four, and one group of five. Jean Baptiste Dumas published work in 1857 describing relationships
between various groups of metals. Although various chemists were able to identify relationships between small groups of elements, they
had yet to build one scheme that encompassed them all.[4]
German chemist August Kekulé had observed in 1858 that carbon has a tendency to bond with other elements in a ratio of one to four.
Methane, for example, has one carbon atom and four hydrogen atoms. This concept eventually became known as valency. In 1864,
fellow German chemist Julius Lothar Meyer published a table of the 49 known elements arranged by valency. The table revealed that
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Portrait of Dmitri Mendeleev
elements with similar properties often shared the same valency.[5]
English chemist John Newlands published a series of papers in 1864 and 1865 that described his attempt at classifying the elements:
When listed in order of increasing atomic weight, similar physical and chemical properties recurred at intervals of eight, which he likened
to the octaves of music.[6][7] This law of octaves, however, was ridiculed by his contemporaries.[8]
Russian chemistry professor Dmitri Ivanovich Mendeleev and Julius Lothar Meyer independently
published their periodic tables in 1869 and 1870, respectively. They both constructed their tables in a
similar manner: by listing the elements in a row or column in order of atomic weight and starting a new
row or column when the characteristics of the elements began to repeat.[9] The success of Mendeleev's
table came from two decisions he made: The first was to leave gaps in the table when it seemed that the
corresponding element had not yet been discovered.[10] Mendeleev was not the first chemist to do so,
but he went a step further by using the trends in his periodic table to predict the properties of those
missing elements, such as gallium and germanium.[11] The second decision was to occasionally ignore
the order suggested by the atomic weights and switch adjacent elements, such as cobalt and nickel, to
better classify them into chemical families. With the development of theories of atomic structure, it
became apparent that Mendeleev had inadvertently listed the elements in order of increasing atomic
number.[12]
With the development of modern quantum mechanical theories of electron configurations within atoms, it
became apparent that each row (or period) in the table corresponded to the filling of a quantum shell of
electrons. In Mendeleev's original table, each period was the same length. However, because larger atoms have more electron
sub-shells, modern tables have progressively longer periods further down the table.[13]
In the years that followed after Mendeleev published his periodic table, the gaps he left were filled as chemists discovered more chemical
elements. The last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939.[14]
The periodic table has also grown with the addition of synthetic and transuranic elements. The first transuranic element to be discovered
was neptunium, which was formed by bombarding uranium with neutrons in a cyclotron in 1939.[15]
Gallery
See also
Alternative periodic tables
Abundance of the chemical elements
Atomic electron configuration table
Discoveries of the chemical elements
Extended periodic table
History of the periodic table
IUPAC's systematic element names
Periodic group
Chemical elements in East Asian
languages
Table of chemical elements
Table of nuclides
Periodic Matrix Sets
Photovoltaic effect
Notes
^ IUPAC article on periodic table (http://www.iupac.org/didac
/Didac%20Eng/Didac01/Content/S01.htm)
1.
^ Science Standards of Learning Curriculum Framework
(http://www.doe.virginia.gov/VDOE/Instruction/Science/ScienceCF-
PS.doc)
2.
^ Dynamic periodic table (http://www.ptable.com/)3.
^ a b Ball, p. 1004.
^ Ball, p. 1015.
^ Newlands, John A. R. (1864-08-20). "On Relations Among the
Equivalents" (http://web.lemoyne.edu/~giunta
/EA/NEWLANDSann.HTML#newlands3) . Chemical News 10:
94–95. http://web.lemoyne.edu/~giunta
/EA/NEWLANDSann.HTML#newlands3.
6.
^ Newlands, John A. R. (1865-08-18). "On the Law of Octaves"
(http://web.lemoyne.edu/~giunta
7.
/EA/NEWLANDSann.HTML#newlands4) . Chemical News 12: 83.
http://web.lemoyne.edu/~giunta
/EA/NEWLANDSann.HTML#newlands4.
^ Bryson, Bill (2004). A Short History of Nearly Everything. London:
Black Swan. pp. 141–142. ISBN 9780552151740.
8.
^ Ball, pp. 100–1029.
^ Pullman, p. 22710.
^ Ball, p. 10511.
^ Atkins, p. 8712.
^ Ball, p. 11113.
^ Adloff, Jean-Pierre; Kaufman, George B. (2005-09-25). Francium
(Atomic Number 87), the Last Discovered Natural Element
(http://chemeducator.org/sbibs/s0010005/spapers/1050387gk.htm) .
The Chemical Educator 10 (5). Retrieved on 2007-03-26.
14.
^ Ball, p. 12315.
References
Atkins, P. W. (1995). The Periodic Kingdom. HarperCollins Publishers, Inc.. ISBN 0-465-07265-8.
Ball, Philip (2002). The Ingredients: A Guided Tour of the Elements. Oxford University Press. ISBN 0-19-284100-9.
Brown, Theodore L.; LeMay, H. Eugene; Bursten, Bruce E. (2005). Chemistry: The Central Science (10th ed.). Prentice Hall.
ISBN 0-13-109686-9.
Periodic table - Wikipedia, the free encyclopedia http://en.wikipedia.org/wiki/Periodic_table
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Pullman, Bernard (1998). The Atom in the History of Human Thought. Translated by Axel Reisinger. Oxford University Press.
ISBN 0-19-515040-6.
Further reading
Bouma, J. (1989). "An Application-Oriented Periodic Table of the Elements". J. Chem. Ed. 66: 741. doi:10.1021/ed066p741
(http://dx.doi.org/10.1021%2Fed066p741) .
Eric Scerri (2007). The periodic table: its story and its significance. Oxford: Oxford University Press. ISBN 0-19-530573-6.
Mazurs, E.G (1974). Graphical Representations of the Periodic System During One Hundred Years. Alabama: University of Alabama
Press.
External links
Interactive periodic table (http://www.ptable.com/)
WebElements (http://www.webelements.com/)
IUPAC periodic table (http://www.iupac.org/reports/periodic_table/index.html)
A video for each one of the elements. (http://www.periodicvideos.com) Made by Brady Haran, featuring Martyn Poliakoff and others,
at the University of Nottingham.
pni:ﻞﺒﯿﭨ کڈﺎﯾﺮﯿﭘ
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