Periodic table

The periodic table is a tabular display of the chemical elements, organized on a basis of their properties. Elements are presented in increasing atomic number; while rectangular in general outline, gaps are included in the rows or periods to keep elements with similar properties together, such as the halogens and the noble gases, in columns or groups, forming distinct rectangular areas or blocks.^[1] Because the periodic table accurately predicts the properties of various elements and the relations between properties, its use is widespread within chemistry, providing a useful framework for analysing chemical behavior, as well as in other sciences.

Although precursors exist, the current table is generally credited to Dmitri Mendeleev, who developed it in 1869 to illustrate periodic trends in the properties of the then-known elements;^[2] the layout has been refined and extended as new elements have been discovered and new theoretical models developed to explain chemical behavior.^[3] Mendeleev's presentation also predicted some properties of then-unknown elements expected to fill gaps in his arrangement; these predictions were proved right when said elements were discovered and found to have properties close to the predictions.

All elements from atomic numbers 1 to 118 have been isolated. Of these, all up to plutonium exist naturally in significant quantities; the rest have only been artificially synthesised in laboratories, along with numerous synthetic radionuclides of naturally occurring elements. Production of elements beyond ununoctium is being pursued, with the question of how the periodic table may need to be modified to accommodate these elements being a matter of ongoing debate.

Contents 1 History 2 Contents 3 Organization 3.1 Periodic trends 3.1.1 Groups 3.1.2 Periods 3.1.3 Blocks 3.2 Uncertainties after element 118 4 Alternatives 5 See also 6 References 7 Bibliography 8 External links

History

In 1789, Antoine Lavoisier published a list of 33 chemical elements. Although Lavoisier grouped the elements into gases, metals, nonmetals, 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.^[5] 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 elements with similar properties often shared the same valency.^[6]

English chemist John Newlands produced a series of papers in 1864 and 1865 that described his own classification of the elements: he noted that 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.^[7][8] This Law of Octaves, however, was ridiculed by his contemporaries, and the Chemical Society refused to publish his work.^[9] Nonetheless, Newlands was able to draft an atomic table and use it to predict the existence of missing elements, such as germanium. The Chemical Society only acknowledged the significance of his discoveries some five years after they credited Mendeleev.

Russian chemistry professor Dmitri Ivanovich Mendeleev and German chemist 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.^[10] 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.^[11] Mendeleev was not the first chemist to do so, but he was the first to be recognized as using the trends in his periodic table to predict the properties of those missing elements, such as gallium and germanium.^[12] 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 listed the elements in order of increasing atomic number.^[13]

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.^[14]

In the years following publication of Mendeleev's periodic table, the gaps he identified were filled as chemists discovered additional naturally occurring elements. It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939.^[15] However, plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring primordial element in 1971.^[16]

The production of various transuranic elements has expanded the periodic table significantly, the first of these being neptunium, synthesized in 1939.^[17] Because many of the transuranic elements are highly unstable and decay quickly, they are challenging to detect and characterize when produced, and there have been controversies concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority, and hence naming rights. The most recently named element is copernicium (number 112), named on 19 February 2010;^[18] the most recently accepted discoveries are ununquadium (114) and ununhexium (116), both accepted on 1 June 2011.^[19]

Contents

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

This common arrangement of the periodic table separates the lanthanides (lanthanoids) and actinides (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 chemical properties Alkali metals Alkaline earth metals Inner transition metals Transition metals Post-transition metals Other nonmetals Halogens Noble gases Lanthanides Actinides
Atomic number colors show state of matter at standard conditions (0 °C and 1 atm): Solids Liquids Gases Unknown	Borders show natural occurrence: Primordial From decay Synthetic

All versions of the periodic table only include chemical elements, not mixtures, compounds, or subatomic particles, and each isotope of a given element is represented in the same cell. In the standard periodic table, the elements are listed in order of increasing atomic number (the number of protons in the nucleus of an atom). A new row (period) is started when a new shell has its first electron. Columns (groups) are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. oxygen and selenium are in the same column because they both have 4 electrons in the outermost p-subshell). The periods are longer further down in the periodic table, and the groups get longer on the right (although the alkali metals, the largest group, is on the far left, and the alkaline earth metals, another large group, are next to the alkali metals). In general, elements with similar chemical properties fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.^[1]

As of 2012, the periodic table contains 118 confirmed chemical elements, of which 112 have been recognized by the International Union of Pure and Applied Chemistry (IUPAC) and named. 94 of these occur naturally, of which 91 occur regularly, and of which 83 are primordial. The other 11 elements only occur in decay chains of primordial elements. All elements from americium to copernicium, while not occurring naturally in the universe, have been officially recognized by the IUPAC as being discovered, while elements 113 to 118 have reportedly been synthesized in laboratories and are currently known only by their systematic element names, based off their atomic numbers.^[20] No elements past 118 have been discovered as of 2012.

In printed or other formally presented periodic tables, each element is provided a formatted cell that usually provides some of the basic properties of the element. Atomic number, element symbol, and name are almost always included, and atomic weights, densities, melting and boiling points, crystal structure as a solid, origin, abbreviated electron configuration, electronegativity, and most common valence numbers are often included as well.^[21]

By definition, each chemical element has a unique atomic number representing the number of protons in its nucleus, but most elements have differing numbers of neutrons among different atoms; these are referred to as isotopes. For example, all atoms of carbon have six protons and usually have six neutrons as well, but about 1% have seven neutrons, and a very small amount have eight neutrons; carbon has three different naturally occuring isotopes. Isotopes are never separated in the periodic table; they are always grouped together under a single element. Elements with no stable isotopes have the atomic masses of their most stable isotopes listed in parentheses.^[22]

In presentations of the periodic table, the lanthanides and the actinides are customarily shown as two additional rows below the main body of the table,^[1] with placeholders or else a selected single element of each series (either lanthanum or lutetium, and either actinium or lawrencium, respectively) shown in a single cell of the main table, between barium and hafnium, and radium and rutherfordium, respectively. This convention is entirely a matter of aesthetics and formatting practicality; a rarely used wide-formatted periodic table inserts the lanthanide and actinide series in their proper places, as parts of the table's sixth and seventh rows (periods).

Many presentations of the periodic table show a dark stair-step diagonal line along the metalloids, with metals to the left of the line and non-metals to the right.^[1] Various other groupings of the chemical elements are sometimes also highlighted on a periodic table, such as transition metals, poor metals, and metalloids. Other informal groupings of the elements exist, such as the platinum group and the noble metals, but are rarely addressed in periodic tables.

Organization

In the modern periodic table, the elements are placed progressively in each period from left to right in the sequence of their atomic numbers, with a new row started after a noble gas. The first element in the next row is always an alkali metal with an atomic number one greater than that of the noble gas (e.g. after krypton, a noble gas with the atomic number 36, a new row is started by rubidium, an alkali metal with the atomic number 37). No gaps currently exist because all elements between hydrogen and ununoctium (element 118) have been discovered. Since the elements are sequenced by atomic number, sets of elements are sometimes specified by terms such as "through" (e.g. through iron), "beyond" (e.g. beyond uranium), or "from ... through" (e.g. 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. For instance cobalt, element 27, is on average heavier than nickel, element 28.

Hydrogen and helium are often placed in different places than their electron configurations would indicate; Hydrogen is usually placed above lithium, in accordance with its electron configuration, but is sometimes placed above fluorine, or even carbon, as it also behaves similar to them. Helium is almost always placed above neon, as they are very similar chemically.^[1]

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.

Periodic trends

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.^[1]

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 (hypothetical)	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 roughly in the order depicted in the table at hand (according to the Aufbau principle; see right).^[23] 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.^[1]

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).^[24] 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.^[1]

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.

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).^[25] The older naming systems differed slightly between Europe and America (the table shown in this section shows the old American Naming System).^[26]

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.^[25]

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.^[27]

Periods

A period is a horizontal row in the periodic table. Although groups are the most common way of classifying elements, there are regions where horizontal trends are more significant than vertical group trends, such as the f-block, where the lanthanides and actinides form two substantial horizontal series of elements.^[28]

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.^[29] 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.^[30]

Blocks

(American Group Numbering System).]]

Because of the importance of the outermost electron 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 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.^[31]

Uncertainties after element 118

Element 118 completes the seventh period of the periodic table. Since the properties of any additional elements are still unknown, it is unclear whether they will continue the pattern of the currently accepted periodic table as an additional period (Period 8), or require further adaptations or adjustments to the currently known patterns. Glenn T. Seaborg expected the next 50 elements to form an eighth period, including a two-element s-block for elements 119 and 120, a g-block (the first) for the next 18 elements (121-138), filling a g-shell of electrons, and the 30 additional elements continuing the current p-, d-, and f-blocks.^[32][33] However, some physicists, including Pekka Pyykkö, have theorized that these additional elements will deviate from the Madelung energy-ordering rule, which predicts how electron shells are filled, and thus affect the appearance of the present periodic table.^[34]

Alternatives

While the iconic format presented above is widely used,^[1] other alternative periodic tables exist, including not only various rectangular formats, but also circular or cylindrical versions in which the rows (periods) flow from one into another, without the arbitrary breaks required at the margins of the usual printed or screen-formatted versions. Alternative periodic tables are developed often to highlight or emphasize different chemical or physical properties of the elements which are not as apparent in traditional periodic tables. Some tables aim to emphasize both the nucleon and electronic structure of atoms. This can be done by changing the spatial relationship or representation each element has with respect to another element in the table. Other tables aim to emphasize the chemical element isolations by humans over time.

A common alternate layout is Charles Janet's Left Step Periodic Table, which organizes elements according to orbital filling. The modern version, known as the ADOMAH Periodic Table, helps with writing electron configurations; the table is oriented 90˚ from the traditional periodic table, with the s-block moved to the end, after the noble gases.^[35]

Another of the most common alternative layouts is Theodor Benfey's periodic table, where elements are arranged in a spiral with hydrogen at the center and spiraling outward, with the transition metals, lanthanides, and actinides as peninsulas.^[36]

Three dimensional periodic tables exist as well, such as Paul Giguere's periodic table, which has four billboards, each representing a block, with elements on the front and back. Hydrogen and helium are omitted.

See also

	Book: Periodic table
Wikipedia books are collections of articles that can be downloaded or ordered in print.

References

Bibliography

• Ball, Philip (2002). The Ingredients: A Guided Tour of the Elements. Oxford University Press. ISBN 0-19-284100-9. 

• Bouma, J. (1989). "An Application-Oriented Periodic Table of the Elements". J. Chem. Ed. 66 (9): 741. Bibcode 1989JChEd..66..741B. doi:10.1021/ed066p741. 

• Hjørland, Birger (2011). "The periodic table and the philosophy of classification". Knowledge Organization 38 (1): 9–21. http://ucla.academia.edu/EricScerri/Papers/432740/Forum%5FThe%5FPhilosophy%5Fof%5FClassification. Retrieved 2011-03-13. 

• Kean, Sam (2010). The Disappearing Spoon - and other true tales from the Periodic Table. London: Black Swan. ISBN 978-0-552-77750-6. 

• Levi, Primo (1984). The Periodic Table [1975]. London: Penguin Books. ISBN 978-0-141-39944-7. 

• Mazurs, E.G (1974). Graphical Representations of the Periodic System During One Hundred Years. Alabama: University of Alabama Press. 

• Moore, John (2003). Chemistry For Dummies. New York: Wiley Publications. p. 111. ISBN 978-0-7645-5430-8. OCLC 51168057. 

• Scerri, Eric (2007). The periodic table: its story and its significance. Oxford: Oxford University Press. ISBN 0-19-530573-6. 

External links

• Interactive periodic table
• Video periodic table
• WebElements
• IUPAC periodic table
• 118 elements: The Periodic Table of Videos made by Brady Haran, featuring Martyn Poliakoff and others, at the University of Nottingham.
• A catalog of various forms of the periodic table
• The periodic table in more than 200 languages, including dead languages

H																															He
Li	Be																									B	C	N	O	F	Ne
Na	Mg																									Al	Si	P	S	Cl	Ar
K	Ca															Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
Rb	Sr															Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
Cs	Ba	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
Fr	Ra	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn	Uut	Uuq	Uup	Uuh	Uus	Uuo
Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Post-transition metals Metalloids Other nonmetals Halogens Noble gases Unknown chem. properties
Large version

Periodic tables Layouts Standard Large table Inline f-block Vertical Text only Metals and nonmetals Blocks Valences Extension beyond the 7th period Large extended table Electron configurations Electronegativities Alternatives Crystal structure Discovery periods List of elements by Name etymology (symbol) Discovery Oxidation state Abundance (in humans) Nuclear stability Hardness Atomic properties Production Groups 1 (Alkali metals) 2 (Alkaline earth metals) 3 4 5 6 7 8 9 10 11 12 13 (Boron group) 14 (Carbon group) 15 (Pnictogens) 16 (Chalcogens) 17 (Halogens) 18 (Noble gases) Other element categories Periods Metals Transition metals Metalloids Nonmetals Lanthanides Actinides Rare earth elements Platinum group metals (PGMs) Post-transition metals Refractory metals Blocks s-block p-block d-block f-block g-block Periods 1 2 3 4 5 6 7 8 Book:Periodic table Category:Periodic table Portal:Chemistry

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