Periodic Table

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Periodic table

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Articles

Overview

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Groups

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Periods

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Overview

Periodic table

Theperiodic table of the chemical elements (also periodic table of the elementsor 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).

Structure

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 * Lanthanoids 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 ** Actinoids 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 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 **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 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.

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Element categories in the periodic table

Metals Metalloids Nonmetals Unknown

chemical properties Alkali metals Alkaline earth metals Inner transition elements Transition elements Other metals Other nonmetals Halogens Noble gases Lanthanides Actinides

Solids Liquids Gases Unknown

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 (groupsor 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-blocksto 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.

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

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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). 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

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.

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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.

Blocks

This diagram shows the periodic table blocks.

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.

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Trends of periods

Periodic trend for ionization energy. Each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases.

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

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 intotriads (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. 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 asvalency. 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.[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]

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Portrait of Dmitri Mendeleev

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]

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Gallery

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.

• Pullman, Bernard (1998).The Atom in the History of Human Thought . Translated by Axel Reisinger. Oxford University Press. ISBN 0-19-515040-6.

External links

• Interactive periodic table[16] • WebElements[17]

• IUPAC periodic table[18]

• A video for each one of the elements.[19]Made by Brady Haran, featuring Martyn Poliakoff and others, at the University of Nottingham.

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References

[1] IUPAC article on periodic table (http://www.iupac.org/didac/Didac Eng/Didac01/Content/S01.htm)

[2] Science Standards of Learning Curriculum Framework (http://www.doe.virginia.gov/VDOE/Instruction/Science/ScienceCF-PS.doc) [3] Dynamic periodic table (http://www.ptable.com/)

[4] Ball, p. 100 [5] Ball, p. 101

[6] Newlands, John A. R. (1864-08-20). "On Relations Among the Equivalents" (http://web.lemoyne.edu/~giunta/EA/NEWLANDSann. HTML#newlands3).Chemical News10: 94 – 95. .

[7] Newlands, John A. R. (1865-08-18). "On the Law of Octaves" (http://web.lemoyne.edu/~giunta/EA/NEWLANDSann. HTML#newlands4).Chemical News12: 83. .

[8] Bryson, Bill (2004). A Short History of Nearly Everything. London: Black Swan. pp. 141 – 142. ISBN 9780552151740. [9] Ball, pp. 100 – 102

[10] Pullman, p. 227 [11] Ball, p. 105 [12] Atkins, p. 87 [13] Ball, p. 111

[14] 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.

[15] Ball, p. 123

[16] http://www.ptable.com/ [17] http://www.webelements.com/

[18] http://www.iupac.org/reports/periodic_table/index.html [19] http://www.periodicvideos.com

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History

Thehistory of the periodic tablereflects over a century of growth in the understanding of chemical properties, and culminates with the publication of the first actual periodic table by Dmitri Mendeleev in 1869.[1]While Mendeleev built upon earlier discoveries by such scientists as Antoine-Laurent de Lavoisier, the Russian scientist is generally given sole credit for development of the actual periodic table itself.

The table itself is a visual representation of the periodic law which states that certain properties of elements repeat

periodically when arranged by atomic number. The table arranges elements into vertical columns (Groups) and horizontal rows (Periods) to display these commonalities.

A modern periodic table with (colored) discovery periods.

Elemental ideas from

ancient times

People have known about some chemical elements such as gold, silver and copper from antiquity, as these can all be discovered in nature in native form and are relatively simple to mine with primitive tools.[2] However, the notion that there were a limited number of elements from which everything was composed originated with the Greek philosopher Aristotle. About 330 B.C Aristotle proposed that everything is made up of a mixture of one or more of four "roots" (originally put forth by the Sicilian philosopher

Empedocles), but later renamed elements by Plato. The four elements were earth, water, air and fire. While the

concept of an element was thus introduced, Aristotle's and Plato's ideas did nothing to advance the understanding of the nature of matter.

Age of Enlightenment

Hennig Brand was the first person recorded to have discovered a new element. Brand was a bankrupt German merchant who was trying to discover the Philosopher's Stone — a mythical object that was supposed to turn inexpensive base metals into gold. He experimented with distilling human urine until in 1649[3]he finally obtained a glowing white substance which he named phosphorus. He kept his discovery secret, until 1680 when Robert Boyle rediscovered it and it became public. This and related discoveries raised the question of what it means for a substance to be an "element".

In 1661 Boyle defined an element as a substance that cannot be broken down into a simpler substance by a chemical reaction. This simple definition actually served for nearly 300 years (until the development of the notion of subatomic particles), and even today is taught in introductory chemistry classes.

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Antoine-Laurent de Lavoisier

Antoine Laurent de Lavoisier

Lavoisier's Traité Élémentaire de Chimie (Elementary Treatise of Chemistry, 1789, translated into English by Robert Kerr) is considered to be the first modern chemical textbook. It contained a list of elements, or substances that could not be broken down further, which included oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur. It also forms the basis for the modern list of elements. His list, however, also included light and caloric, which he believed to be material substances. While many leading chemists of the time refused to believe Lavoisier's new revelations, the Elementary Treatise was written well enough to convince the younger generation. However, as Lavoisier's descriptions only classified elements as metals or non-metals, it fell short of a complete analysis.

Johann Wolfgang Döbereiner

In 1817, Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. He found that some elements formed groups of three with related properties. He termed these groups "triads". Some triads classified by Döbereiner are:

1. chlorine, bromine, and iodine 2. calcium, strontium, and barium 3. sulfur, selenium, and tellurium 4. lithium, sodium, and potassium

In all of the triads, the atomic weight of the second element was almost exactly the average of the atomic weights of the first and third element.[4]

Classifying Elements

By 1869[3], a total of 63[3]elements had been discovered. As the number of known elements grew, scientists began to recognize patterns in the way chemicals reacted and began to devise ways to classify the elements.

Alexandre-Emile Béguyer de Chancourtois

Alexandre-Emile Béguyer de Chancourtois, a French geologist, was the first person to notice the periodicity of the elements — similar elements seem to occur at regular intervals when they are ordered by their atomic weights. He devised an early form of periodic table, which he called the telluric helix. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois saw that elements with similar properties lined up vertically. His chart included some ions and compounds in addition to elements. His paper was published in 1862, but used geological rather than chemical terms and did not include a diagram; as a result, it received little attention until the work of Dmitri Mendeleev.[5]

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John Newlands

John Newlands was an English chemist who in 1865 classified

John Newlands was an English chemist who in 1865 classified[6][6]the 56 elements that had been discovered at thethe 56 elements that had been discovered at the time into 11 groups which were based on similar physical properties.

time into 11 groups which were based on similar physical properties.

J. A. R. Newlands' law of octaves J. A. R. Newlands' law of octaves

Newlands noted that many pairs of Newlands noted that many pairs of similar elements existed which differed similar elements existed which differed by some multiple of eight in atomic by some multiple of eight in atomic weight. However, his

weight. However, his law of octaveslaw of octaves,, likening this periodicity of eights to likening this periodicity of eights to thethe musical scale, was ridiculed by his musical scale, was ridiculed by his contemporaries. It was not until the contemporaries. It was not until the following century, with Gilbert N. following century, with Gilbert N. Lewis' valence bond theory (1916) and Lewis' valence bond theory (1916) and

Irving Langmuir's octet theory of chemical bonding

Irving Langmuir's octet theory of chemical bonding[7][7] [8][8](1919) that the importance of the periodicity of eight would(1919) that the importance of the periodicity of eight would be accepted.

be accepted.

Dmitri Mendeleev

Mendeleev's 1869 periodic table Mendeleev's 1869 periodic table

Dmitri Ivanovich Mendeleev Dmitri Ivanovich Mendeleev

Dmitri Mendeleev, a Siberian-born Dmitri Mendeleev, a Siberian-born Russian chemist, was the first scientist Russian chemist, was the first scientist to make a periodic table much like the to make a periodic table much like the one we use today. Mendeleev arranged one we use today. Mendeleev arranged the elements in a table ordered by the elements in a table ordered by atomic weight, corresponding to atomic weight, corresponding to relative molar mass as defined today. It relative molar mass as defined today. It is sometimes said that he played is sometimes said that he played "chemical solitaire" on long train rides "chemical solitaire" on long train rides using cards with various facts of using cards with various facts of known elements.

known elements.[9][9]On March 6, 1869,On March 6, 1869, a formal presentation was made to the a formal presentation was made to the Russian Chemical Society, entitled Russian Chemical Society, entitledTheThe Dependence Between the Properties of Dependence Between the Properties of the Atomic Weights of the Elements the Atomic Weights of the Elements.. His table was published in an obscure His table was published in an obscure Russian journal but quickly Russian journal but quickly republished in a German journal, republished in a German journal,

Zeitschrift für Chemie

Zeitschrift für Chemie (Eng.,(Eng., "Chemistry Magazine"), in 1869. It "Chemistry Magazine"), in 1869. It stated:

stated:

1.

1. The The elemeelements, nts, if arif arrangeranged acd accordincordingg to their atomic weights, exhibit an to their atomic weights, exhibit an apparent periodicity of properties. apparent periodicity of properties. 2.

2. Elements which aElements which are similar as regards re similar as regards to their chemical propeto their chemical properties have atomic werties have atomic weights which are eitheights which are either of nearlyr of nearly the same value (e.g., Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs).

the same value (e.g., Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs). 3.

3. The arrangement of The arrangement of the elements, or of the elements, or of groups of elements in groups of elements in the order of their atomic the order of their atomic weights, corresponds weights, corresponds toto their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among other series in that of Li, Be, Ba, C, N, O, and Sn.

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4. The elements which are the most widely diffused have small atomic weights.

5. The magnitude of the atomic weight determines the character of the element, just as the magnitude of the molecule determines the character of a compound body.

6. We must expect the discovery of many yet unknown elements – for example, elements analogous to aluminium and silicon – whose atomic weight would be between 65 and 75.

7. The atomic weight of an element may sometimes be amended by a knowledge of those of its contiguous

elements. Thus the atomic weight of tellurium must lie between 123 and 126, and cannot be 128. (This was based on the position of tellurium between antimony and iodine whose atomic weight is 127. However Moseley later explained the position of these elements without revising the atomic weight values —see below.)

8. Certain characteristic properties of elements can be foretold from their atomic weights.

This version of Mendeleev's periodic table from 1891. It is lacking the noble gases

Scientific benefits of Mendeleev's table

• Mendeleev predicted the discovery of other elements and left space for these new elements, namely eka-silicon (germanium), eka-aluminium (gallium), and eka-boron (scandium). Thus, there was no disturbance in the periodic table. • He pointed out that some of the then

current atomic weights were incorrect. • He provided for variance from atomic

weight order.

Shortcomings of Mendeleev's table

• His table did not include any of the noble gases, which were discovered later. These were added by Sir William Ramsay as Group 0, without any disturbance to the basic concept of the periodic table.

• There was no place for the isotopes of the various elements, which were discovered later.

Lothar Meyer

Unknown to Mendeleev, Lothar Meyer was also working on a periodic table. Although his work was published in 1864, and was done independently of Mendeleev, few historians regard him as an equal co-creator of the periodic table. For one thing, Meyer's table only included 28 elements. Furthermore, Meyer classified elements not by atomic weight, but by valence alone. Finally, Meyer never came to the idea of predicting new elements and correcting atomic weights. Only a few months after Mendeleev published his periodic table of all known elements (and predicted several new elements to complete the table, plus some corrected atomic weights), Meyer published a virtually identical table. While a few people consider Meyer and Mendeleev the co-creators of the periodic table, most agree that, by itself, Mendeleev's accurate prediction of the qualities of the undiscovered elements lands him the larger share of credit. In any case, at the time Mendeleev's predictions greatly impressed his contemporaries and were eventually found to be correct. An English chemist, William Odling, also drew up a table that is remarkably similar to that of Mendeleev, in 1864.

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Refinements to the periodic table

Henry Moseley

In 1914 Henry Moseley found a relationship between an element's X-ray wavelength and its atomic number (Z), and therefore resequenced the table by nuclear charge rather than atomic weight. Before this discovery, atomic numbers were just sequential numbers based on an element's atomic weight. Moseley's discovery showed that atomic numbers had an experimentally measurable basis.

Thus Moseley placed argon (Z=18) before potassium (Z=19) based on their X-ray wavelengths, despite the fact that argon has a greater atomic weight (39.9) than potassium (39.1). The new order agrees with the chemical properties of these elements, since argon is a noble gas and potassium an alkali metal. Similarly, Moseley placed cobalt before nickel, and was able to explain that tellurium occurs before iodine without revising the experimental atomic weight of tellurium (127.6) as proposed by Mendeleev.

Moseley's research also showed that there were gaps in his table at atomic numbers 43 and 61 which are now known to be Technetium and Promethium, respectively, both radioactive and not naturally occurring. Following in the footsteps of Dmitri Mendeleev, Henry Moseley also predicted new elements.

Glenn T. Seaborg

During his Manhattan Project research in 1943 Glenn T. Seaborg experienced unexpected difficulty isolating Americium (95) and Curium (96). He began wondering if these elements more properly belonged to a different series which would explain why the expected chemical properties of the new elements were different. In 1945, he went against the advice of colleagues and proposed a significant change to Mendeleev's table: the actinide series. Seaborg's actinide concept of heavy element electronic structure, predicting that the actinides form a transition series analogous to the rare earth series of lanthanide elements, is now well accepted in the scientific community and included in all standard configurations of the periodic table. The actinide series are the second row of the f-block (5f series) and comprise the elements from Actinium to Lawrencium. Seaborg's subsequent elaborations of the actinide concept theorized a series of superheavy elements in a transactinide series comprising elements 104 through 121 and a superactinide series inclusive of elements 122 through 153.

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Main discovery periods

The history of the periodic table is also a history of the discovery of the chemical elements. IUPAC[10]suggest five "main discovery periods":

• Before 1800 (36 elements): discoveries during and before the Enlightenment.

• 1800-1849 (+22 elements): impulse from scientific (empirical processes systematization and modern atomic theory) and industrial revolutions.

• 1850-1899 (+23 elements): the age of classifying elements received an impulse from the spectrum analysis. Boisbaudran, Bunsen, Crookes, Kirchhoff, and others "hunting emission line signatures".

• 1900-1949 (+13 elements): impulse from the old quantum theory, the consolitated periodic table, and quantum mechanics.

• 1950-1999 (+15 elements): "atomic bomb" and Particle physics issues, for atomic numbers 97 and above.

The periodic table as a cultural icon

Throughout the 20th century, the periodic table grew in ubiquity. Its presence on a classroom wall tells the movie-viewing audience that they are viewing a science classroom. It is often provided to students taking standardized tests as a necessary tool to complete chemical problems.

In 1998, a 35-by-65 foot periodic table was constructed at the Science Museum of Virginia and is a Guinness World Record.[11]

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External links

• History of the Development of the Periodic Table of Elements[12] • Development of the periodic table[13]

• Classification of the elements[14] • The path to the periodic table[15]

• Web page listing several scholarly and semi-popular articles on various aspects of the periodic system and underlying theoretical concepts. Some are downloadable![16]

• Periodic Table Database[17]

References

[1] IUPAC article on periodic table (http://www.iupac.org/didac/Didac Eng/Didac01/Content/S01.htm)

[2] Scerri, E. R. (2006).The Periodic Table: Its Story and Its Significance; New York City, New York; Oxford University Press. [3] "A Brief History of the Development of Periodic Table" (http://www.wou.edu/las/physci/ch412/perhist.htm). .

[4] Leicester, Henry M. (1971).The Historical Background of Chemistry; New York City, New York; Dover Publications. [5] Annales des Mines history page (http://www.annales.org/archives/x/chancourtois.html).

[6] in a letter published in the Chemical News in February 1863, according to the Notable Names Data Base (http://www.nndb.com/people/ 480/000103171/)

[7] Irving Langmuir,“The Structure of Atoms and the Octet Theory of Valence”, Proceedings of the National Academy of Science, Vol. V, 252, Letters (1919) - online at (http:// dbhs.wvusd.k12.ca.us/webdocs/Chem-History/Langmuir-1919.html)

[8] Irving Langmuir,“The Arrangement of Electrons in Atoms and Molecules”, Journal of the American Chemical Society, Vol. 41, No, 6, pg. 868 (June 1919) - beginning and ending of the paper are transcribed online at (http://dbhs.wvusd.k12.ca.us/webdocs/Chem-History/ Langmuir-1919b.html); the middle is missing

[9] Physical Science, Holt Rinehart & Winston (January 2004), page 302 ISBN 0-03-073168-2 [10] http://old.iupac.org/reports/periodic_table/index.html

[11] "Richmond in the Record Books" (http://www.richmond.com/museums-galleries/6300). August 18, 2004. . [12] http://www.bpc.edu/mathscience/chemistry/history_of_the_periodic_table.html [13] http://www.chemsoc.org/viselements/pages/history.html [14] http://members.optushome. com.au/scottsofta/Pintro.htm [15] http://www.chemheritage.org/EducationalServices/chemach/ppt/ppt.html [16] http://www.chem.ucla.edu/dept/Faculty/scerri/index.html [17] http://www.meta-synthesis.com/webbook/35_pt/pt_database.php

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Alternative periodic tables

Theodor Benfey's periodic table Alternative periodic tables are

tabulations of chemical elements differing significantly in their organization from the traditional depiction of the Periodic System.[1] [2] Several have been devised, often purely for didactic reasons, as not all correlations between the chemical elements are effectively captured by the standard periodic table. A 1974 review of the tables then known is considered a definitive work on the topic.[3]

Aims

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 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.

Major alternatives

Janet's Left Step Periodic Table [4] (1928) is considered to be the most significant alternative to the traditional depiction of the periodic system. It organizes elements according to orbital filling and is widely used by physicists. Its modern version, known as ADOMAH Periodic Table[5], (2006) is helpful for writing electron configurations.[6]

f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 p1 p2 p3 p4 p5 p6 s1 s2 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 Uue Ubn

The Janet Periodic Table (with current element symbols).[7]

In Theodor Benfey's periodic table (1960), the elements form a two-dimensional spiral, starting from hydrogen, and folding their way around two islands, the transition metals, and lanthanides and actinides. A superactinide island is already slotted in. The Chemical Galaxy (2004) is organized in a similar way.

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In the extended periodic table, suggested by Glenn T. Seaborg in 1969, yet unknown elements are included up to atomic number 218. Helium is placed in the group 2 elements.

Mendeleev's 1869 periodic table

The oldest periodic table is the short form table of Dmitri Mendeleev, which shows secondary chemical kinships. For example, the alkali metals and the coinage metals (copper, silver, gold) are in the same column because both groups tend to have a valence of one. This format is still used by many, as shown by this contemporary Russian short form table [8] which includes all elements and element names to date. Timmothy Stowe's physicist's periodic

table is three-dimensional with the three axes representing the principal quantum number, orbital quantum number, and orbital magnetic quantum number. Helium is again a group 2 element.

Paul Giguere's 3-D periodic table consists of 4 billboards with the elements written on the front and the back. The first billboard has the group 1 elements on the front and the group 2 elements at the back, with hydrogen and helium omitted altogether. At a 90° angle the second billboard contains the groups 13 to 18 front and back. Two more billboard each making 90° angles contain the other elements.

In the research field of superatoms, clusters of atoms have properties of single atoms of another element. It is suggested to extend the periodic table with a second layer to be occupied with these cluster compounds. The latest addition to this multi-story table is the aluminum cluster ion Al−₇, which behaves like a multivalent germanium atom.[9]

Ronald L. Rich has proposed a periodic table where elements appear more than once when appropriate.[10]He notes that hydrogen shares properties with group 1 elements based on valency, with group 17 elements because hydrogen is a non-metal but also with the carbon group based on similarities in chemical bonding to transition metals and a similar electronegativity. In this rendition of the periodic table carbon and silicon also appear in the same group as titanium and zirconium. An interesting new chemists' table ("Newlands Revisited") that solves many of the problems of position of hydrogen, helium and the lanthanides, etc., was published by EG Marks and JA Marks in 2010.[11]

External links

• Janet's Left Step Periodic Table[12]

• Correction to Physicist periodic table offered by Jeries Rihani[13]as Meitnerium occupies the position that Hassium should have.

• A Wired Article on Alternate Periodic Tables[14] • A Selection of Periodic Tables[15]

• http://periodicspiral.com/arranges the periodic table in a (hexagonal) spiral. • Rotaperiod.com[16]A new periodic table.

• Note[17]on the T-shirt topology of the Z-spiral.

• New Periodic Table of the Elements[18]this is in a square-triangular periodic arrangement. • Periodic Table based onelectron configurations[5]

• Database of Periodic Tables[17] • Periodic Fractal of the Elements[19]

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• Bob Doyle Periodic Table of the Elements[20]A regrouping by properties that suggests a maximum of 120 elements.

• Earth Scientist's Periodic Table[21]

References

[1] E.R. Scerri.The Periodic Table, Its Story and Its Significance.Oxford University Press, New York, 2007.

[2] Henry Bent. New Ideas in Chemistry from Fresh Energy for the Periodic Law.AuthorHouse, 2006. ISBN 9781425948627

[3] Mazurs, E. G. Graphical Representations of the Periodic System During One Hundred Years. Alabama; University of Alabama Press,1974. ISBN 0-8173-3200-6.

[4] http://www.meta-synthesis.com/webbook/35_pt/pt_database.php?PT_id=152 [5] http://www.perfectperiodictable.com/userguide

[6] Philip J. Stewart:Charles Janet: unrecognized genius of the periodic system.Foundations of Chemistry. January, 2009. ISSN 1386-4238 (Print) ISSN 1572-8463 (Online), Vol.12, No.1 April, 2010;

[7] WebElements (http://www.webelements.com/nexus/Janet_Periodic_Table) : The Janet Periodic Table. [8] http://flerovlab. jinr.ru/flnr/dimg/Periodic_Table.jpg

[9] Beyond The Periodic Table Metal clusters mimic ch emical properties of atomsIvan Amato Chemical & Engineering News November 21,

2006Link (http://pubs.acs.org/cen/news/84/i48/8448notw8.html) [10] Rich, Ronald L. J. Chem. Educ.200582 1761

[11] Foundations of Chemistry 2010, 12: 85-93 (http://www.springerlink.com/content/q7j4670426845322/fulltext.pdf) Newlands revisited: a display of the periodicity of the chemical elements for chemists.

[12] http://www.meta-synthesis.com/webbook/35_pt/pt.html#j [13] http://jeries.rihani.com/references.html [14] http://www.wired.com/wired/archive/13.10/start.html?pg=11 [15] http://www.meta-synthesis.com/webbook/35_pt/pt.html [16] http://www.rotaperiod.com [17] http://arxiv.org/abs/physics/0603026 [18] http://www.egregoralfa.republika.pl/english/newtable.html [19] http://www.superliminal.com/pfractal.htm [20] http://www.wizworld.com/dox/doyle_periodic_table.gif [21] http://www.gly.uga.edu/railsback/PT.html

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Element

Achemical elementis a pure chemical substance consisting of one typeof atom distinguished by its atomic number, which is the number of protons in its nucleus.[1] Common examples of elements are iron, copper, silver, gold, hydrogen, carbon, nitrogen, and oxygen. In total, 118 elements have been observed as of March 2010, of which 92 occur naturally on Earth. Of these, oxygen is the most abundant element in the earth's crust. 80 elements have stable isotopes, namely all elements with atomic numbers 1 to 82, except elements 43 and 61 (technetium and promethium). Elements with atomic numbers 83 or higher (bismuth and above) are inherently unstable, and undergo radioactive decay. The elements from atomic number 83 to 92 have no stable nuclei, but are nevertheless found in nature, either surviving as remnants of the primordial stellar nucleosynthesis that produced the elements in the solar system, or else produced as short-lived daughter-isotopes through the natural decay of uranium and thorium.[2] All chemical matter consists of these elements. New elements of higher atomic number are discovered from time to time, as products of artificial nuclear reactions. When two distinct elements are chemically combined, the result is termed a compound.

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History

Mendeleev's 1869 periodic table

Ancient philosophy posited a set of classical elements to explain patterns in nature.

Elements originally referred to earth, water ,

air and fire rather than the chemical elements of modern science.

The term 'elements' (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE, in his dialogue Timaeus, which includes a discussion of the composition of inorganic and organic bodies and is a speculative treatise on chemistry. Plato believed the elements introduced a century earlier by Empedocles were composed of small polyhedral forms: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).[3] [4]

Aristotle, c. 350 BCE, also used the term

stoicheia and added a fifth element called

aether, which formed the heavens. Aristotle defined an element as:

Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.[5]

In 1661, Robert Boyle showed that there were more than just four classical elements as the ancients had assumed.[6] The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric.[7]By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that cannot be decomposed into any simpler substance.[6]Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. In 1913, Henry Moseley discovered that the physical basis of the atomic number of the atom was its nuclear charge, which eventually led to the current definition. The current definition also avoids some ambiguities due to isotopes and allotropes.

By 1919, there were seventy-two known elements.[8]In 1955, element 101 was discovered and named mendelevium in honor of Mendeleev, the first to arrange the elements in a periodic manner. In October 2006, the synthesis of element 118 was reported; the synthesis of element 117 was reported in April 2010.[9]

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Description

The lightest elements are hydrogen and helium, both theoretically created by Big Bang nucleosynthesis during the first 20 minutes of the universe[10] in a ratio of around 3:1 by mass (approximately 12:1 by number of atoms). Almost all other elements found in nature, including some further hydrogen and helium created since then, were made by various natural or (at times) artificial methods of nucleosynthesis, including occasionally breakdown activities such as nuclear fission, alpha decay, cluster decay, and cosmic ray spallation.

As of 2010, there are 118 known elements (in this context, "known" means observed well enough, even from just a few decay products, to have been differentiated from any other element).[11] [12] Of these 118 elements, 94 occur naturally on Earth. Six of these occur in extreme trace quantities: technetium, atomic number 43; promethium, number 61; astatine, number 85; francium, number 87; neptunium, number 93; and plutonium, number 94. These 94 elements, and also possibly element 98 californium, have been detected in the universe at large, in the spectra of stars and also supernovae, where short-lived radioactive elements are newly being made.

The remaining 24 elements, not found on Earth or in astronomical spectra, have been derived artificially. All of the elements that are derived solely through artificial means are radioactive with very short half-lives; if any atoms of these elements were present at the formation of Earth, they are extremely likely to have already decayed, and if present in novae, have been in quantities too small to have been noted. Technetium was the first purportedly non-naturally occurring element to be synthesized, in 1937, although trace amounts of technetium have since been found in nature, and the element may have been discovered naturally in 1925. This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally occurring trace elements.

Lists of the elements are available by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as Ionization energies of the elements. The most convenient presentation of the elements is in the periodic table, which groups elements with similar chemical properties together.

Atomic number

The atomic number of an element, Z , is equal to the number of protons that defines the element. For example, all carbon atoms contain 6 protons in their nucleus; so the atomic number "Z" of carbon is 6. Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known as isotopes of the element.

The number of protons in the atomic nucleus also determines its electric charge, which in turn determines the electrons of the atom in its non-ionized state. This in turn (by means of the Pauli exclusion principle) determines the atom's various chemical properties. So all carbon atoms, for example, ultimately have identical chemical properties because they all have the same number of protons in their nucleus, and therefore have the same atomic number. It is for this reason that atomic number rather than mass number (or atomic weight) is considered the identifying characteristic of an element.

Atomic mass

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a super-index on the left hand side of the atomic symbol (e.g.,238U).

The relative atomic mass of an element is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit (u). This number may be a fraction that is not close to a whole number, due to the averaging process. On the other hand, the atomic mass of a pure isotope is quite close to its mass number. Whereas the mass number is a natural (or whole) number, the atomic mass of a single isotope is a real number that is close to a natural number. In general, it differs slightly from the mass number as the mass of the protons and neutrons is not exactly 1 u, the electrons also contribute

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slightly to the atomic mass, and because of the nuclear binding energy. For example, the mass of 19F is 18.9984032 u. The only exception to the atomic mass of an isotope not being a natural number is 12C, which has a mass of exactly 12, because u isdefined as 1/12th of the mass of a free carbon-12 atom.

Isotopes

Isotopes are atoms of the same element (that is, with the same number of protons in their atomic nucleus), but having

different numbers of neutrons. Most (66 of 94) naturally occurring elements have more than one stable isotope. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons in the nucleus, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, the three isotopes of carbon are known as carbon-12, carbon-13, and carbon-14, often abbreviated to 12C, 13C, and 14C. Carbon in everyday life and in chemistry is a mixture of 12C,13C, and14C atoms.

Except in the case of the isotopes of hydrogen (which differ greatly from each other in relative mass —enough to cause chemical effects), the isotopes of the various elements are typically chemically nearly indistinguishable from each other. For example, the three naturally occurring isotopes of carbon have essentially the same chemical properties, but different nuclear properties. In this example, carbon-12 and carbon-13 are stable atoms, but carbon-14 is unstable; it is radioactive, undergoing beta decay into nitrogen-14.

As illustrated by carbon, all of the elements have some isotopes that are radioactive (radioisotopes), which decay into other elements upon radiating an alpha or beta particle. Certain elements onlyhave radioactive isotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one stable isotope, and the mean number of stable isotopes for the 80 stable elements is 3.1 stable isotopes per element. The largest number of stable isotopes that occur for an element is 10 (for tin, element 50).

Allotropes

Atoms of pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple structures (spacial arrangements of atoms), known as allotropes, which differ in their properties. For example, carbon can be found as diamond, which has a tetrahedral structure around each carbon atom; graphite, which has layers of carbon atoms with a hexagonal structure stacked on top of each other; graphene, which is a single layer of graphite that is incredibly strong; fullerenes, which have nearly spherical shapes; and carbon nanotubes, which are tubes with a hexagonal structure (even these may differ from each other in electrical properties). The ability for an element to exist in one of many structural forms is known as 'allotropy'.

Standard state

The standard state, or reference state, of an element is defined as its thermodynamically most stable state at 1 bar at a given temperature (typically at 298.15 K). In thermochemistry, an element is defined to have an enthalpy of formation of zero in its standard state. For example, the reference state for carbon is graphite, because it is more stable than the other allotropes.

Nomenclature

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When it was learned, existing names (e.g., gold, mercury, iron) were kept in most countries, and national differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For example, the Germans use "Wasserstoff" for "hydrogen", "Sauerstoff" for "oxygen" and "Stickstoff" for "nitrogen", while English and some romance languages use "sodium" for "natrium"

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and "potassium" for "kalium", and the French, Italians, Greeks, Portuguese and Poles prefer "azote/azot/azoto" for "nitrogen".

But for international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that "aluminium" and "caesium" take the place of the US spellings "aluminum" and "cesium", while the US "sulfur" takes the place of the British "sulphur". Chemicals that are practical to sell in bulk in many countries, however, still have national names, and those that do not use the Latin alphabet cannot be expected to use the IUPAC name.

According to IUPAC, the full name of an element is not capitalized, even if it is derived from a proper noun such as the elements californium or einsteinium (unless it would be capitalized by some other rule, such as to begin a sentence). Isotopes of chemical elements are also uncapitalized if written out: carbon-12 or uranium-235. Symbols of chemical elements, however, are capitalized: thus the symbols for the elements just discussed are Cf and Es; C-12 and U-235.

In the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have a half life too short for them to remain in any appreciable amounts. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question that delayed naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy).

Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century. For example,lutetium was named in reference to Paris, France. The Germans were reluctant to relinquish naming rights to the French, often calling it cassiopeium. The British discoverer of niobium originally named it columbium, in reference to the New World. It was used extensively as such by American publications prior to international standardization.

Chemical symbols

For the listing of current and not used Chemical symbols, and other symbols that look like chemical symbols, please see List of elements by symbol.

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules.

The current system of chemical notation was invented by Berzelius. In this typographical system chemical symbols are not used as mere abbreviations - though each consists of letters of the Latin alphabet - they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully universal; since Latin was the common language of science at that time, they were abbreviations based on the Latin names of metals - Cu comes from Cuprum, Fe comes from Ferrum, Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Later chemical elements were also assigned unique chemical symbols, based on the name of the element, but not necessarily in English. For example, sodium has the chemical symbol 'Na' after the Latinnatrium. The same applies to "W" (wolfram) for tungsten, "Fe" (ferrum) for Iron, "Hg" (hydrargyrum) for mercury, "Sn" (stannum) for tin, "K" (kalium) for potassium, "Au" (aurum) for gold, "Ag" (argentum) for silver, "Pb" (plumbum) for lead, and "Sb" (stibium) for antimony.

Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not

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be confused with a roman numeral.

The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always lower case (small letters).

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical, meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

Isotope symbols

The three main isotopes of the element hydrogen are often written as H for protium, D for deuterium and T for tritium. This is in order to make it easier to use them in chemical equations, as it replaces the need to write out the mass number for each atom. E.g. the formula for heavy water may be written D₂O instead of ²H₂O.

The periodic table

The periodic table of the chemical elements is a tabular method of displaying the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869. Mendeleev 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.

The periodic table is now ubiquitous within the academic discipline of chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behavior. The table has also found wide application in physics, biology, engineering, and industry. The current standard table contains 118 confirmed elements as of April 10, 2010.

Abundance

During the early phases of the Big Bang, nucleosynthesis of hydrogen nuclei resulted in the production of hydrogen and helium isotopes, as well as very minuscule amounts (on the order of 10−10) of lithium and beryllium. There is argument about whether or not some boron was produced in the Big Bang, as it has been observed in some very young stars,[13]but no heavier elements than boron were produced. As a result, the primordial abundance of atoms consisted of roughly 75%1H, 25%4He, and 0.01% deuterium.[14]Subsequent enrichment of galactic halos occurred due to Stellar nucleosynthesis and Supernova nucleosynthesis.[15] However intergalactic space can still closely resemble the primordial abundance, unless it has been enriched by some means.

The following graph (note log scale) shows abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass.[16] Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. As physical laws and processes appear common throughout the visible universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

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Abundances of the chemical elements in the Solar system.

Element Parts per million by mass Hydrogen 739,000 Helium 240,000 Oxygen 10,400 Carbon 4,600 Neon 1,340 Iron 1,090 Nitrogen 960 Silicon 650 Magnesium 580 Sulfur 440 Potassium 210 Nickel 100

Recently discovered elements

The first transuranium element (element with atomic number greater than 92) discovered was neptunium in 1940. As of February 2010, only the elements up to 112, copernicium, have been confirmed as discovered by IUPAC, while more or less reliable claims have been made for synthesis of elements 113, 114, 115, 116 and 118. The discovery of element 112 was acknowledged in 2009, and the name 'copernicium' and the atomic symbol 'Cn' were suggested for it.[17]The name and symbol were officially endorsed by IUPAC on February 19, 2010.[18]The heaviest element that is believed to have been synthesized to date is element 118, ununoctium, on October 9, 2006, by the Flerov Laboratory of Nuclear Reactions in Dubna, Russia.[19] [20]

Element 117, ununseptium, has been synthesised[21], and its place in the periodic table is preestablished.

On April 24, 2008, Amnon Marinov and six other researchers claimed that element 122 has been detected in purified natural thorium.[22] [23] If confirmed, this would be the first naturally occurring heavy element discovered in more than 50 years. It has yet to be proved as it is still under confirmation by the university but could be a major development as previously all transuranic elements were artificial. The claim of Marinov et al. was criticized by a part of the scientific community, and Marinov says he has submitted the article to the journals Nature and Nature Physicsbut both turned it down without sending it for peer review.[24]

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External links

• Videos for each element[25]by the University of Nottingham

References

[1] International Union of Pure and Applied Chemistry. " (http://goldbook.iupac.org/C01022.html)".Compendium of Chemical Terminology

Internet edition.

[2] A. Earnshaw, Norman Greenwood (1997).Chemistry of the Elements(2 ed.). Butterworth-Heinemann.

[3] http://books.google.com/books?id=xSjvowNydN8C&lpg=PP1&ots=eRla--Y6Ul&dq=Plato%20timaeus&pg=PA45#v=onepage& q=cube&f=false

[4] Hillar, Marian (2004). "The Problem of the Soul in Aristotle's De anima" (http://www.socinian.org/aristotles_de_anima.html). NASA WMAP. . Retrieved 2006-08-10.

[5] Partington, J.R. (1937). A Short History of Chemistry. New York: Dover Publications, Inc.. ISBN 0486659771. [6] Boyle, Robert (1661).The Sceptical Chymist . London. ISBN 0922802904.

[7] Lavoisier, Antoine Laurent (1790). Elements of chemistry translated by Robert Kerr (http://books.google.com/?id=4BzAjCpEK4gC& pg=PA175). Edinburgh. pp. 175 – 176. ISBN 9780415179140. .

[8] Carey, George, W. (1914).The Chemistry of Human Life. Los Angeles. ISBN 0766128407. [9] http://www.nytimes.com/2010/04/07/science/07element.html?hp

[10] Gaitskell, R; Utyonkov, V. K.; Lobanov, Yu. V.; Abdullin, F. Sh.; Polyakov, A. N.; Sagaidak, R. N.; Shirokovsky, I. V.; Tsyganov, Yu. S.et al. (2006). "Evidence for Dark Matter" (http:// gaitskell.brown.edu/physics/talks/0408_SLAC_SummerSchool/

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[15] G. Wallerstein, I. Iben Jr., P. Parker, A. M. Boesgaard, G. M. Hale, A. E. Champagne, C. A. Barnes, F. KM-dppeler, V. V. Smith, R. D. Hoffman, F. X. Timmes, C. Sneden, R.N. Boyd, B.S. Meyer, D.L. Lambert (1999). "Synthesis of the elements in stars: forty years of progress"

Periodic Table

Periodic table

Periodic table

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Articles

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Overview

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Periodic table

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Elemental ideas from

ancient times

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Antoine-Laurent de Lavoisier

Johann Wolfgang Döbereiner

Classifying Elements

Alexandre-Emile Béguyer de Chancourtois

John Newlands

John Newlands

Dmitri Mendeleev

Dmitri Mendeleev

Lothar Meyer

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Henry Moseley

Glenn T. Seaborg

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The periodic table as a cultural icon

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