Periodic table
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Contents Articles Overview
1
Periodic table
1
History
9
Alternative periodic tables
16
Element
19
Isotope
27
Orbital
36
Groups
45
Group
45
Group I
46
Group II
49
Group III
52
Group IV
55
Group V
60
Group VI
61
Group VII
63
Group VIII
64
Group IX
65
Group X
66
Group XI
67
Group XII
69
Group XIII
71
Group XIV
72
Group XV
74
Group XVI
76
Group XVII
78
Group XVIII
83
Periods
85
Period
85
Pediod 1
90
Extensions
92
Blocks
95
Block
95
s-block
95
p-block
96
d-block
97
f-block
99
Other divisions
100
Actinide
100
Lanthanide
104
Metal
109
Metalloid
116
Noble gas
117
Noble metal
128
Nonmetal
131
Platinum group
132
Post-transition metal
135
Transactinide element
137
Transuranium element
138
Transition metal
142
See also
146
Table of nuclides
146
Island of stability
149
References Article Sources and Contributors
154
Image Sources, Licenses and Contributors
159
Article Licenses License
162
1
Overview Periodic table The periodic table of the chemical elements (also periodic table of the elements or just the periodic table) is a tabular display of the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.[1] The periodic table is now ubiquitous within the academic discipline of chemistry, providing a useful framework to classify, systematize, and compare all of the many different forms of chemical behavior. The table has found many applications in chemistry, physics, biology, and engineering, especially chemical engineering. The current standard table contains 118 elements to date. (elements 1 – 118).
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 4 Li Be
5 B
6 C
7 N
8 O
9 F
10 Ne
3
11 12 Na Mg
13 Al
14 Si
15 P
16 S
17 Cl
18 Ar
4
19 20 K Ca
21 Sc
22 23 Ti V
28 29 30 31 Ni Cu Zn Ga
32 Ge
33 As
34 Se
35 Br
36 Kr
5
37 38 Rb Sr
39 Y
40 41 42 43 44 45 Zr Nb Mo Tc Ru Rh
46 47 48 Pd Ag Cd
49 In
50 Sn
51 Sb
52 Te
53 I
54 Xe
6
55 56 * 72 73 Cs Ba Lanthanoids Hf Ta
78 79 80 81 Pt Au Hg Tl
82 Pb
83 Bi
84 Po
85 At
86 Rn
7
87 88 Fr Ra
** Actinoids
24 25 26 Cr Mn Fe
74 W
75 76 Re Os
27 Co
77 Ir
104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo
* Lanthanoids
57 La
58 59 60 61 62 63 Ce Pr Nd Pm Sm Eu
** Actinoids
89 Ac
90 91 Th Pa
64 65 66 67 Gd Tb Dy Ho
92 93 94 95 96 97 98 99 U Np Pu Am Cm Bk Cf Es
68 Er
69 Tm
70 Yb
71 Lu
100 101 102 103 Fm Md No 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.
Periodic table
2 Element categories in the periodic table
Metals Alkali metals
Alkaline earth metals
Inner transition elements
Metalloids Transition elements
Other metals
Nonmetals
Unknown chemical Other Halogens Noble properties nonmetals gases
Lanthanides Actinides
Primordial Solids Liquids Gases Unknown
From decay
Synthetic (Undiscovered)
Other alternative periodic tables exist. Some versions of the table show a dark stair-step line along the metalloids. Metals are to the left of the line and non-metals to the right.[2] The layout of the periodic table demonstrates recurring ("periodic") chemical properties. Elements are listed in order of increasing atomic number (i.e., the number of protons in the atomic nucleus). Rows are arranged so that elements with similar properties fall into the same columns (groups or families). According to quantum mechanical theories of electron configuration within atoms, each row ( period ) in the table corresponded to the filling of a quantum shell of electrons. There are progressively longer periods further down the table, grouping the elements into s-, p-, d- and f-blocks to reflect their electron configuration. In printed tables, each element is usually listed with its element symbol and atomic number; many versions of the table also list the element's atomic mass and other information, such as its abbreviated electron configuration, electronegativity and most common valence numbers. As of 2010, the table contains 118 chemical elements whose discoveries have been confirmed. Ninety-four are found naturally on Earth, and the rest are synthetic elements that have been produced artificially in particle accelerators. Elements 43 (technetium), 61 (promethium) and all elements greater than 83 (bismuth), beginning with 84 (polonium) have no stable isotopes. The atomic mass of each of these element's isotope having the longest half-life is typically reported on periodic tables with parentheses.[3] Isotopes of elements 43, 61, 93 (neptunium) and 94 (plutonium), first discovered synthetically, have since been discovered in trace amounts on Earth as products of natural radioactive decay processes. The primary determinant of an element's chemical properties is its electron configuration, particularly the valence shell electrons. For instance, any atoms with four valence electrons occupying p orbitals will exhibit some similarity. The type of orbital in which the atom's outermost electrons reside determines the "block" to which it belongs. The number of valence shell electrons determines the family, or group, to which the element belongs. 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
Periodic table
3 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.
Periodic table
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 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 This diagram shows the periodic table blocks. 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.
4
Periodic table
Trends of periods Elements in the same period show trends in atomic radius, ionization energy, electron affinity, and electronegativity. Moving left to right across a period, atomic radius usually decreases. This occurs because each successive element has an added proton and electron which causes the electron to be drawn closer to the nucleus. This decrease in atomic radius also causes the ionization energy to Periodic trend for ionization energy. Each period begins at a minimum for the alkali increase when moving from left to metals, and ends at a maximum for the noble gases. 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 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. 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.[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]
5
Periodic table 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]
6
Portrait of Dmitri Mendeleev
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 [13] sub-shells, modern tables have progressively longer periods further down the table. 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]
Periodic table
Gallery
See also • • • • • • • • • • • • •
Alternative periodic tables Abundance of the chemical elements Atomic electron configuration table Discoveries of the chemical elements Extended periodic table History of the periodic table IUPAC's systematic element names Periodic group Chemical elements in East Asian languages Table of chemical elements Table of nuclides Periodic Matrix Sets Photovoltaic effect
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Periodic table
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.
Further reading • Bouma, J. (1989). "An Application-Oriented Periodic Table of the Elements". J. Chem. Ed. 66: 741. doi:10.1021/ed066p741. • Eric Scerri (2007). The periodic table: its story and its significance . Oxford: Oxford University Press. ISBN 0-19-530573-6. • Mazurs, E.G (1974). Graphical Representations of the Periodic System During One Hundred Y ears. Alabama: University of Alabama Press.
External links • • • •
Interactive periodic table [16] WebElements [17] [18] IUPAC periodic table A video for each one of the elements. [19] Made by Brady Haran, featuring Martyn Poliakoff and others, at the University of Nottingham.
pni: ک ا ی ر ی پ ل ب ی
References [1] [2] [3] [4] [5] [6]
IUPAC article on periodic table (http://www.iupac.org/didac/Didac Eng/Didac01/Content/S01.htm) Science Standards of Learning Curriculum Framework (http://www.doe.virginia.gov/VDOE/Instruction/Science/ScienceCF-PS.doc) Dynamic periodic table (http://www.ptable.com/) Ball, p. 100 Ball, p. 101 Newlands, John A. R. (1864-08-20). "On Relations Among the Equivalents" (http://web.lemoyne.edu/~giunta/EA/NEWLANDSann. HTML#newlands3). Chemical News 10: 94 – 95. . [7] Newlands, John A. R. (1865-08-18). "On the Law of Octaves" (http://web.lemoyne.edu/~giunta/EA/NEWLANDSann. HTML#newlands4). Chemical News 12: 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
History The history of the periodic table reflects 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.
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 A modern periodic table with (colored) discovery periods. 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.
9
History
10
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
Antoine Laurent de Lavoisier
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. 2. 3. 4.
chlorine, bromine, and iodine calcium, strontium, and barium sulfur, selenium, and tellurium 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]
History
11
John Newlands John Newlands was an English chemist who in 1865 classified [6] the 56 elements that had been discovered at the time into 11 groups which were based on similar physical properties. Newlands noted that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. However, his law of octaves, likening this periodicity of eights to the musical scale, was ridiculed by his contemporaries. It was not until the J. A. R. Newlands' law of octaves following century, with Gilbert N. Lewis' valence bond theory (1916) and Irving Langmuir's octet theory of chemical bonding[7] [8] (1919) that the importance of the periodicity of eight would be accepted.
Dmitri Mendeleev Dmitri Mendeleev, a Siberian-born Russian chemist, was the first scientist to make a periodic table much like the one we use today. Mendeleev arranged the elements in a table ordered by atomic weight, corresponding to relative molar mass as defined today. It is sometimes said that he played "chemical solitaire" on long train rides using cards with various facts of known elements.[9] On March 6, 1869, a formal presentation was made to the Russian Chemical Society, entitled The
Mendeleev's 1869 periodic table
Dependence Between the Properties of the Atomic Weights of the Elements. His table was published in an obscure Russian journal but quickly republished in a German journal, Zeitschrift für Chemie (Eng., "Chemistry Magazine"), in 1869. It stated: 1. The elements, elements, if arrange arrangedd accordin accordingg to their atomic weights, exhibit an apparent periodicity of properties.
Dmitri Ivanovich Mendeleev
2. Elements which are are similar as regards to their chemical properties properties have atomic weights weights which are either either of nearly the same value (e.g., Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs). 3. The arrangement of the elements, or of groups of elements in the order of their atomic weights, corresponds to 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.
History
12
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. 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. This version of Mendeleev's periodic table from 1891. It is lacking the noble gases
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.
History
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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History
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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History
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See also • • • • •
Prout's hypothesis History of chemistry Discoveries of the chemical elements Periodic table Alternative periodic tables
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] [2] [3] [4] [5] [6]
IUPAC article on periodic table (http://www.iupac.org/didac/Didac Eng/Didac01/Content/S01.htm) Scerri, E. R. (2006). The Periodic Table: Its Story and Its Significance ; New York City, New York; Oxford University Press. "A Brief History of the Development of Periodic Table" (http://www.wou.edu/las/physci/ch412/perhist.htm). . Leicester, Henry M. (1971). The Historical Background of Chemistry ; New York City, New York; Dover Publications. Annales des Mines history page (http://www.annales.org/archives/x/chancourtois.html). 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
Alternative periodic tables
16
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] Alternative
periodic
tables
Aims Theodor Benfey's periodic table
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 [7]
The Janet Periodic Table (with current element symbols).
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.
Alternative periodic tables
17
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. 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.
Mendeleev's 1869 periodic table
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− , which behaves like a multivalent germanium 7 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 on electron configurations [5] • Database of Periodic Tables [17] • Periodic Fractal of the Elements [19]
Alternative periodic tables • 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 atoms Ivan Amato Chemical & Engineering News November 21, 2006 Link (http://pubs.acs.org/cen/news/84/i48/8448notw8.html) [10] Rich, Ronald L. J. Chem. Educ. 2005 82 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
18
Element
19
Element A chemical element is a pure chemical substance consisting of one type of 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.
The periodic table of the chemical elements
Element
20
History 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]
Mendeleev's 1869 periodic table
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]
Element
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
21
Element 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 is defined 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, and 14C 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 only have 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"
22
Element
23
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 Latin natrium. 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
Element 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.
24
Element
25
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 [22] [23] natural thorium. 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 Physics but both turned it down without sending it for peer review.[24]
Element
See also • • • • • • • • • • • • • • •
Abundance of the chemical elements Compound Chemical symbol Chemistry Discovery of the chemical elements Elements song Fictional element Goldschmidt classification Island of stability List of chemical element name etymologies List of elements by atomic number List of elements by name Periodic table Systematic element name Prices of elements and their compounds
Further reading • E.R. Scerri, The Periodic Table, Its Story and Its Significance, Oxford University Press, NY, 2007.
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/ Gaitskell_DMEvidence_v16.pdf) (PDF). Physical Review C 74: timeline on page 10. doi:10.1103/PhysRevC.74.044602. . Retrieved 2008-10-08. [11] Sanderson, Katherine (17 October 2006). "Heaviest element made - again" (http://www.nature.com/news/2006/061016/full/061016-4. html).
[email protected]. Nature. . Retrieved 2006-10-19. [12] Phil Schewe and Ben Stein (17 October 2006). "Elements 116 and 118 Are Discovered" (http://www.aip.org/pnu/2006/797.html). Physics News Update. American Institute of Physics. . Retrieved 2006-10-19. [13] Hubble Observations Bring Some Surprises - New York Times (http://query.nytimes.com/gst/fullpage. html?res=9E0CE5D91F3AF937A25752C0A964958260) [14] Wright, Edward L. (September 12, 2004). "Big Bang Nucleosynthesis" (http://www.astro.ucla.edu/~wright/BBNS.html). UCLA Division of Astronomy. . Retrieved 2007-02-22. [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"
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Element (http://www.cococubed.com/papers/wallerstein97.pdf) (pdf). Reviews of Modern Physics 69 (4): 995 – 1084. doi:10.1103/RevModPhys.69.995. . Retrieved 2006-08-04. [16] Croswell, Ken (February 1996). Alchemy of the Heavens (http://kencroswell.com/alchemy.html). Anchor. ISBN 0-385-47214-5. . [17] "IUPAC Announces Start of the Name Approval Process for the Element of Atomic Number 112" (http://media.iupac.org/news/ 112_Naming_Process_20090720.pdf). 20 July 2009. . Retrieved 2009-08-27. [18] IUPAC (International Union of Pure and Applied Chemistry): Element 112 is Named Copernicium (http://www.iupac.org/web/nt/ 2010-02-20_112_Copernicium) [19] Phil Schewe and Ben Stein (17 October 2006). "Elements 116 and 118 Are Discovered" (http://www.aip.org/pnu/2006/797.html). Physics News Update. American Institute of Physics. . Retrieved 2006-10-19. [20] Oganessian, Yu. Ts. et al.; Utyonkov, V.; Lobanov, Yu.; Abdullin, F.; Polyakov, A.; Sagaidak, R.; Shirokovsky, I.; Tsyganov, Yu. et al. (2006-10-09). "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245Cm+48Ca fusion reactions". Physical Review C 74 (4): 044602. doi:10.1103/PhysRevC.74.044602. [21] http://www. jinr.ru/img_sections/PAC/NP/31/PAK_NP_31_recom_eng.pdf [22] Marinov, A.; Rodushkin, I.; Kolb, D.; Pape, A.; Kashiv, Y.; Brandt, R.; Gentry, R. V.; Miller, H. W. (2008). "Evidence for a long-lived superheavy nucleus with atomic mass number A=292 and atomic number Z=~122 in natural Th" (http://arxiv.org/abs/0804.3869). ArXiv.org 74: 044602. doi:10.1103/PhysRevC.74.044602. . Retrieved 2008-04-28. [23] Battersby, Stephen (2008-05-02). "Has the heaviest element been found?" (http://www.newscientist.com/article/ dn13828-has-the-heaviest-element-been-found.html). NewScientist. . Retrieved 2009-05-01. [24] Van Noorden, Richard (2008-05-02). "Heaviest element claim criticised" (http://www.rsc.org/chemistryworld/News/2008/May/ 02050802. asp). Chemistry World (Royal Society of Chemistry). . Retrieved 2009-05-01. [25] http://periodicvideos.com/
Isotope Isotopes are different types of atoms (nuclides) of the same chemical element, each having a different number of
neutrons. In a corresponding manner, isotopes differ in mass number (or number of nucleons) but never in atomic number.[1] The number of protons (the atomic number) is the same because that is what characterizes a chemical element. For example, carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13 and 14, respectively. The atomic number of carbon is 6, so the neutron numbers in these isotopes of carbon are therefore 12−6 = 6, 13−6 = 7, and 14 – 6 = 8, respectively. A nuclide is an atomic nucleus with a specified composition of protons and neutrons. The nuclide concept emphasizes nuclear properties over chemical properties, while the isotope concept emphasizes chemical over nuclear. The neutron number has drastic effects on nuclear properties, but negligible effects on chemical properties. Since isotope is the older term, it is better known, and is still sometimes used in contexts where nuclide might be more appropriate, such as nuclear technology. An isotope and/or nuclide is specified by the name of the particular element (this indicates the atomic number implicitly) followed by a hyphen and the mass number (e.g. helium-3, carbon-12, carbon-13, iodine-131 and uranium-238). When a chemical symbol is used, e.g., "C" for carbon, standard notation is to indicate the number of nucleons with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left (e.g. 32He, 42He, 126C, 146C, 23592U, and 23992U). Some isotopes are radioactive and are therefore described as radioisotopes or radionuclides, while others have never been observed to undergo radioactive decay and are described as stable isotopes. For example, 14C is a radioactive form of carbon while 12C and 13C are stable isotopes. There are about 339 naturally occurring nuclides on Earth[2] , of which 288 are primordial nuclides. These include 31 nuclides with very long half lives (over 80 million years) and 257 which are formally considered as "stable"[2] . About 30 of these "stable" isotopes have actually been observed to decay, but with half lives too long to be estimated so far. This leaves 227 nuclides that have not been observed to decay at all. Many more apparently "stable" isotopes are predicted by theory to be radioactive, with extremely long half-lives (this does not count the posibility of proton decay, which would make all nuclides unstable). Of the 227 nuclides never observed to decay, only 90 of these (all from the first 40 elements) are stable in theory to all known forms of
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Isotope
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decay. Element 41 (niobium) is theoretically unstable to spontaneous fission, but this has never been detected. Many other stable nuclides are in theory energetically susceptible to other known forms of decay such as alpha decay or double beta decay, but no decay has yet been observed. The half lives for these processes often exceed a million times the estimated age of the universe. Adding in the radioactive nuclides that have been created artificially, there are more than 3100 currently known nuclides.[3] . These include 905 nuclides which are either stable, or have half lives longer than 60 minutes. See list of nuclides for details.
History of the term The term isotope was coined in 1913 by Margaret Todd, a Scottish physician, during a conversation with Frederick Soddy (to whom she [4] was distantly related by marriage). Soddy, a chemist at Glasgow University, explained that it appeared from his investigations as if each position in the periodic table was occupied by multiple entities. Hence Todd made the suggestion, which Soddy adopted, that a suitable name for such an entity would be the Greek term for "at the same place". Soddy's own studies were of radioactive (unstable) atoms. The first observation of different stable isotopes for an element was by J. J. Thomson in 1913. As part of his exploration into the composition of canal rays, Thomson channeled streams of neon ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path. Each stream created a glowing patch on the plate at the point it struck. Thomson observed two separate patches of light on the photographic plate (see image), which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest. F.W. Aston subsequently discovered different stable isotopes for numerous elements using a mass spectrograph.
In the bottom right corner of JJ Thomson's photographic plate are the separate impact marks for the two isotopes of neon: neon-20 and neon-22.
Variation in properties between isotopes Chemical and molecular properties A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and electrons and share a similar electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium ( 1H) and deuterium (2H), because deuterium has twice the mass of protium. The mass effect between deuterium and the relatively light protium also affects the behavior of their respective chemical bonds, by means of changing the center of gravity (reduced mass) of the atomic systems. However, for heavier elements, which have more neutrons than lighter elements, the ratio of the nuclear mass to the collective electronic mass is far greater, and the relative mass difference between isotopes is much less. For these two reasons, the mass-difference effects on chemistry are usually negligible.
Isotope In similar manner, two molecules that differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. As a consequence, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.
Nuclear properties and stability Atomic nuclei consist of protons and neutrons bound together by the residual strong force. Because protons are positively charged, they repel each other. Neutrons, Isotope half lifes. Note that the plot for stable isotopes diverges from the line, protons Z = neutrons N as the element number Z becomes larger which are electrically neutral, stabilize the nucleus in two ways. Their copresence pushes protons slightly apart, reducing the electrostatic repulsion between the protons, and they exert the attractive nuclear force on each other and on protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, so does the ratio of neutrons to protons necessary to ensure a stable nucleus (see graph at right). For example, although the neutron:proton ratio of 32He is 1:2, the neutron:proton ratio of 23892U is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 (Z = N). The nuclide 4020Ca (calcium-40) is the heaviest stable nuclide with the same number of neutrons and protons; all heavier stable nuclides contain more neutrons than protons.
Numbers of isotopes per element Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin). Xenon is the only element that has nine stable isotopes. Cadmium has eight stable isotopes. Five elements have seven stable isotopes, eight have six stable isotopes, ten have five stable isotopes, eight have four stable isotopes, nine have three stable isotopes, 16 have two stable isotopes (counting 180m73Ta as stable), and 26 elements have only a single stable isotope (of these, 19 are so-called mononuclidic elements, having a single primordial stable isotope that dominates and fixes the atomic weight of the natural element to high precision; 3 [5] radioactive mononuclidic elements occur as well). In total, there are 257 nuclides that have not been observed to decay. For the 80 elements that have one or more stable isotopes, the average number of stable isotopes is 257/80 = 3.2 isotopes per element.
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Isotope
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Even and odd nucleon numbers Even/odd N Mass
E
O
All
Stable 145 101 246 Longlived
20
6
26
Primordial 165 107 272
The proton:neutron ratio is not the only factor affecting nuclear stability. Adding neutrons to isotopes can vary their nuclear spins and nuclear shapes, causing differences in neutron capture cross-sections and gamma spectroscopy and nuclear magnetic resonance properties. Even mass number
Beta decay of an even-even nucleus produces an odd-odd nucleus, and vice versa. An even number of protons or of neutrons are more stable (lower binding energy) because of pairing effects, so even-even nuclei are much more stable than odd-odd. One effect is that there are few stable odd-odd nuclei, but another effect is to prevent beta decay of many even-even nuclei into another even-even nucleus of the same mass number but lower energy, because decay proceeding one step at a time would have to pass through an odd-odd nucleus of higher energy. This makes for a larger number of stable even-even nuclei, up to three for some mass numbers, and up to seven for some atomic (proton) numbers. Double beta decay directly from even-even to even-even skipping over an odd-odd nuclide is only occasionally possible, and even then with a half-life greater than a billion times the age of the universe. Even-mass-number nuclides have integer spin and are bosons. Even proton-even neutron
Even/odd Z, N p,n EE OO EO OE
Stable 140
5
53
48
16
4
2
4
Primordial 156
9
55
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Longlived
For example, the extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five or eight nucleons from existing for long enough to serve as platforms for the buildup of heavier elements during fusion formation in stars (see triple alpha process). There are 141 stable even-even isotopes, forming 55% of the 257 stable isotopes. There are also 16 primordial longlived even-even isotopes. As a result, many of the 41 even-numbered elements from 2 to 82 have many primordial isotopes. Half of these even-numbered elements have six or more stable isotopes. All even-even nuclides have spin 0 in their ground state.
Isotope
31
Odd proton-odd neutron
Only five stable nuclides contain both an odd number of protons and an odd number of neutrons: the first four odd-odd nuclides 21H, 63Li, 105B, and 147N (where changing a proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio) and 180m73Ta, which has not yet been observed to decay despite experimental attempts[6] . Also, four long-lived radioactive odd-odd nuclides (4019K, 5023V, 13857La, 17671Lu) occur naturally. Of these 9 primordial odd-odd nuclides, only 147N is the most common isotope of a common element, because it is a part of the CNO cycle; 63Li and 105B are minority isotopes of elements that are rare compared to other light elements, while the other six isotopes make up only a tiny percentage of their elements. Few odd-odd nuclides (and none of the primordial ones) have spin 0 in the ground state. Odd mass number
There is only one beta-stable nuclide per odd mass number because there is no difference in binding energy between even-odd and odd-even comparable to that between even-even and odd-odd, and other nuclides of the same mass are free to beta decay towards the lowest-energy one. For mass numbers 5, 147, 151, and 209 and up, the one beta-stable isobar is able to alpha decay, so that there are no stable isotopes with these mass numbers. This gives a total of 101 stable isotopes with odd mass numbers. Odd-mass-number nuclides have half-integer spin and are fermions. Odd proton-even neutron
These form most of the stable isotopes of the odd-numbered elements, but there is only one stable odd-even isotope for each of the 41 odd-numbered elements from 1 to 81, except for technetium (43Tc) and promethium (61Pm) that have no stable isotopes, and chlorine (17Cl), potassium (19K), copper (29Cu), gallium (31Ga), bromine (35Br), silver (47Ag), antimony (51Sb), iridium (Ir|BL=77), and thallium (81Tl), each of which has two, making a total of 48 stable odd-even isotopes. There are also four primordial long-lived odd-even isotopes, 87 Rb, 115 In, 151 Eu, and 37 49 63 187 Re. 75 Even proton-odd neutron
There are 54 stable isotopes that have an even number of protons and an odd number of neutrons. There are also four primordial long lived even-odd isotopes, 11348Cd (beta decay, half-life is 7.7 × 1015 years); 14762Sm (1.06 × 1011a); and 149 Sm (>2 × 1015a); and the fissile 235 U. 62
92
The only even-odd isotopes that are the most common one for their element are 195 Pt and 9 Be. Beryllium-9 is the 78 4 only stable beryllium isotope because the expected beryllium-8 has higher energy than two alpha particles and therefore decays to them. Odd neutron number
Isotope
32
Even/odd N n
E O
Stable 188 58 Longlived
20
6
Primordial 208 64
The only odd-neutron-number isotopes that are the most common isotope of their element are 19578Pt, 94Be and 14 N. 7 Actinides with odd neutron number are generally fissile, while those with even neutron number are generally not, though they are split when bombarded with fast neutrons.
Occurrence in nature Elements are composed of one or more naturally occurring isotopes. The unstable (radioactive) isotopes are either primordial, in which case they have persisted down to the present because their rate of decay is so slow (e.g., uranium-238 and potassium-40), or they are postprimordial, created by cosmic ray bombardment as cosmogenic nuclides (e.g., tritium, carbon-14) or by the decay of a radioactive primordial isotope to a radioactive radiogenic nuclide daughter (e.g., uranium to radium). As discussed above, only 80 elements have any stable isotopes, and 2 6 of these have only one stable isotope. Thus, about two thirds of stable elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin ( 50Sn). There are about 94 elements found naturally on Earth (up to plutonium inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists estimate that the elements that occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total.[7] Only 257 of these naturally occurring isotopes are stable in the sense of either never having been observed to decay as of the present time (227 nuclides), or having been observed to decay but with a half life too long to estimate (30 nuclides). An additional 31 primordial nuclides (to a total of 288 primordial nuclides), are radioactive with known half lives, but have half lives longer than 80 million years, allowing them to exist from the beginning of the solar system. See list of nuclides for details. All the known stable isotopes occur naturally on Earth; the other naturally occurring-isotopes are radioactive but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. These include the afore-mentioned cosmogenic nuclides and the short-lived radioisotopes formed by decay of a primordial radioactive isotope, such as radon and radium from uranium. An additional ~3000 radioactive isotopes not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale. The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. Before the discovery of isotopes, empirically determined noninteger values of atomic mass confounded scientists. For example, a sample of chlorine contains 75.8% chlorine-35 and 24.2% chlorine-37, giving an average atomic mass of 35.5 atomic mass units. According to generally accepted cosmology theory, only isotopes of hydrogen and helium, and traces of some isotopes of lithium and beryllium were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously produced isotopes. (See nucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar
Isotope system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.
Atomic mass of isotopes The atomic mass (mr) of an isotope is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus (see mass defect), the slight difference in mass between proton and neutron, and the mass of the electrons associated with the atom, the latter because the electron:nucleon ratio differs among isotopes. The mass number is a dimensionless quantity. The atomic mass, on the other hand, is measured using the atomic mass unit based on the mass of the carbon atom. It is denoted with symbols "u" (for unit) or "Da" (for Dalton). The atomic masses of naturally occurring isotopes of an element determine the atomic weight of the element. When the element contains N isotopes, the equation below is applied for the atomic weight M : where m1, m2, ..., m N are the atomic masses of each individual isotope, and x1, ... , x N are the relative abundances of these isotopes.
Applications of isotopes Several applications exist that capitalize on properties of the various isotopes of a given element. Isotope separation is a significant technological challenge, particularly with heavy elements such as uranium or plutonium. Lighter elements such as lithium, carbon, nitrogen, and oxygen are commonly separated by gas diffusion of their compounds such as CO and NO. The separation of hydrogen and deuterium is unusual since it is based on chemical rather than physical properties, for example in the Girdler sulfide process. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectroscopy.
Use of chemical and biological properties • Isotope analysis is the determination of isotopic signature, the relative abundances of isotopes of a given element in a particular sample. For biogenic substances in particular, significant variations of isotopes of C, N and O can occur. Analysis of such variations has a wide range of applications, such as the detection of adulteration of food products.[8] The identification of certain meteorites as having originated on Mars is based in part upon the isotopic signature of trace gases contained in them.[9] • Another common application is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy. For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling). • A technique similar to radioisotopic labeling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials. • Isotopic substitution can be used to determine the mechanism of a reaction via the kinetic isotope effect.
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Isotope
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Use of nuclear properties • Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are 1H, 2D,15N, 13C, and 31P. • Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as 57Fe. • Radionuclides also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes.
See also • • • • • • •
Atom Table of nuclides Radionuclide (or radioisotope) Nuclear medicine (includes medical isotopes) Isotopomer List of particles Isotopes are nuclides having the same number of protons; compare:
• Isotones are nuclides having the same number of neutrons. • Isobars are nuclides having the same mass number, i.e. sum of protons plus neutrons. • Nuclear isomers are different excited states of the same type of nucleus. A transition from one isomer to another is accompanied by emission or absorption of a gamma ray, or the process of internal conversion. (Not to be confused with chemical isomers.) • Bainbridge mass spectrometer
External links • • • • • • • • • • •
Nucleonica Nuclear Science Portal [10] Nucleonica Nuclear Science Wiki [11] International Atomic Energy Agency [12] [13] Atomic weights of all isotopes Atomgewichte, Zerfallsenergien und Halbwertszeiten aller Isotope [14] Chart of the Nuclides [15] produced by the Knolls Atomic Power Laboratory $25 Exploring the Table of the Isotopes [16] at the LBNL Current isotope research and information [17] Radioactive Isotopes [18] by the CDC Interacive Chart of the nuclides, isotopes and Periodic Table [19] The LIVEChart of Nuclides - IAEA [20] with isotope data, in Java [20] or HTML [21]
Isotope
References [1] [2] [3] [4]
IUPAC http://goldbook.iupac.org/I03331.html "Radioactives Missing From The Earth" (http://www.don-lindsay-archive.org/creation/isotope_list.html). . "NuDat 2 Description" (http://www.nndc.bnl.gov/nudat2/help/index. jsp). . Budzikiewicz H, Grigsby RD (2006). "Mass spectrometry and isotopes: a century of research and discussion". Mass spectrometry reviews 25 (1): 146 – 57. doi:10.1002/mas.20061. PMID 16134128. [5] Sonzogni , Alejandro. "Interactive Chart of Nuclides" (http://www.nndc.bnl.gov/chart/). National Nuclear Data Center: Brook haven National Laboratory. . [6] hhttp://bryza.if.uj.edu.pl/zdfk/wp-includes/publications/misiaszek_180mTa_2009.pdf Search for the radioactivity of 180mTa using an underground HPGe sandwich spectrometer, 2009 [7] (http://www.don-lindsay-archive.org/creation/isotope_list.html) [8] E. Jamin et al. (2003). "Improved Detection of Added Water in Orange Juice by Simultaneous Determination of the Oxygen-18/Oxygen-16 Isotope Ratios of Water and Ethanol Derived from Sugars"" (http://pubs.acs.org/cgi-bin/article.cgi/jafcau/2003/51/i18/pdf/jf030167& nbsp;m.pdf). J. Agric. Food Chem. 51: 5202. doi:10.1021/jf030167 m. . [9] A. H. Treiman, J. D. Gleason and D. D. Bogard (2000). ""The SNC meteorites are from Mars"" (http://www.sciencedirect.com/ science?_ob=ArticleURL&_udi=B6V6T-41WBDHD-8&_user=2400262&_coverDate=10/31/2000&_alid=678948366&_rdoc=3& _fmt=summary&_orig=search&_cdi=5823&_sort=r&_docanchor=&view=c&_ct=89&_acct=C000057185&_version=1& _urlVersion=0&_userid=2400262&md5=c5ae2aa8ea60dbd76c2870048730a299). Planet. Space. Sci. 48: 1213. doi:10.1016/S0032-0633(00)00105-7. . [10] http://www.nucleonica.net [11] http://www.nucleonica.net/wiki/index.php/Special:Allpages/Help: [12] http://www.IAEA.org [13] http://physics.nist.gov/cgi-bin/Compositions/stand_alone.pl?ele=&ascii=html&isotype=some [14] http://atom.kaeri.re.kr/ [15] http://www.chartofthenuclides.com/ [16] http://ie.lbl.gov/education/isotopes.htm [17] http://www.isotope.info/ [18] http://www.bt.cdc.gov/radiation/isotopes/ [19] http://www.yoix.org/elements.html [20] http://www-nds.iaea.org/livechart [21] http://www-nds.iaea.org/relnsd/vcharthtml/VChartHTML.html
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Orbital
Orbital An atomic orbital is a mathematical function that describes the wave-like behavior of either one electron or a pair of electrons in an atom.[1] This function can be used to calculate the probability of finding any electron of an atom in any specific region around the atom's nucleus. These functions may serve as three-dimensional graph of an electron’s likely location. The term may thus refer directly to the physical region defined by the function where the electron is likely to be.[2] Specifically, atomic orbitals are the possible quantum states of an individual electron in the collection of electrons around a single atom, as described by the orbital function. Despite the obvious analogy to planets revolving around the Sun, electrons cannot be described as solid particles and so atomic orbitals rarely, if ever, resemble a planet's elliptical path. A more accurate analogy might be that of a large and often oddly-shaped atmosphere (the electron), distributed around a relatively tiny planet (the atomic nucleus). Atomic orbitals exactly describe the shape of this atmosphere only when a single electron is present in an atom. When more electrons are added to a single atom, the additional electrons tend to more evenly fill in a volume of space around the nucleus so that the resulting collection (sometimes termed the atom ’s “electron cloud” [3] ) tends toward a generally spherical zone of probability describing where the atom’s electrons will be found. The idea that electrons might revolve around a compact nucleus with definite angular momentum was convincingly argued in 1913 by Niels Bohr,[4] and the Japanese physicist Hantaro Nagaoka published an orbit-based hypothesis for electronic behavior as early as 1904.[5] However, it was not until 1926 that the solution of the Schrödinger equation for electron-waves in atoms provided the functions for the modern orbitals.[6] Because of the difference from Electron atomic and molecular orbitals. The chart of orbitals (left) is arranged by classical mechanical orbits, the term increasing energy (see Madelung rule). Note that atomic orbits are functions of three "orbit" for electrons in atoms, has been variables (two angles, and the distance from the nucleus, r). These images are faithful to replaced with the term orbital —a term the angular component of the orbital, but not entirely representative of the orbital as a whole. first coined by chemist Robert Mulliken in 1932.[7] Atomic orbitals are typically described as “hydrogen-like” (meaning one-electron) wave functions over space, categorized by n, l, and m quantum numbers, which correspond to the electrons' energy, angular momentum, and an angular momentum direction, respectively. Each orbital is defined by a different set of quantum numbers and contains a maximum of s two electrons. The simple names
36
Orbital orbital, p orbital, d orbital and f orbital refer to orbitals with angular
momentum quantum number l = 0, 1, 2 and 3 respectively. These names indicate the orbital shape and are used to describe the electron configurations as shown on the right. They are derived from the characteristics of their spectroscopic lines: sharp, principal, diffuse, and f undamental, the rest being named in alphabetical order (omitting j).[8] [9] From about 1920, even before the advent of modern quantum mechanics, the aufbau principle (construction principle) that atoms were built up of pairs of electrons, arranged in simple repeating patterns of increasing odd numbers (1,3,5,7..), had been used by Niels Bohr and others to infer the Computed hydrogen atom orbital for n=6, l=0, m=0 presence of something like atomic orbitals within the total electron configuration of complex atoms. In the mathematics of atomic physics, it is also often convenient to reduce the electron functions of complex systems into combinations of the simpler atomic orbitals. Although each electron in a multi-electron atom is not confined to one of the “one-or-two-electron atomic orbitals” in the idealized picture above, still the electron wave-function may be broken down into combinations which still bear the imprint of atomic orbitals; as though, in some sense, the electron cloud of a many-electron atom is still partly “composed” of atomic orbitals, each containing only one or two electrons. The physicality of this view is best illustrated in the repetitive nature of the chemical and physical behavior of elements which results in the natural ordering known from the 19th century as the periodic table of the elements. In this ordering, the repeating periodicity of 2, 6, 10, and 14 elements in the periodic table corresponds with the total number of electrons which occupy a complete set of s, p, d and f atomic orbitals, respectively.
Orbital names Orbitals are given names in the form: where X is the energy level corresponding to the principal quantum number n, type is a lower-case letter denoting the shape or subshell of the orbital and it corresponds to the angular quantum number l, and y is the number of electrons in that orbital. For example, the orbital 1s2 (pronounced "one ess two") has two electrons and is the lowest energy level (n = 1) and has an angular quantum number of l = 0. In X-ray notation, the principal quantum number is given a letter associated with it. For n = 1, 2, 3, 4, 5, ..., the letters associated with those numbers are K , L, M , N , O, ... respectively.
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Orbital
38
Formal quantum mechanical definition In quantum mechanics, the state of an atom, i.e. the eigenstates of the atomic Hamiltonian, is expanded (see configuration interaction expansion and basis (linear algebra)) into linear combinations of anti-symmetrized products (Slater determinants) of one-electron functions. The spatial components of these one-electron functions are called atomic orbitals. (When one considers also their spin component, one speaks of atomic spin orbitals.) In atomic physics, the atomic spectral lines correspond to transitions (quantum leaps) between quantum states of an atom. These states are labelled by a set of quantum numbers summarized in the term symbol and usually associated to particular electron configurations, i.e. by occupations schemes of atomic orbitals (e.g. 1s2 2s2 2 p6 for the ground state of neon -- term symbol: 1S ). 0
This notation means that the corresponding Slater determinants have a clear higher weight in the configuration interaction expansion. The atomic orbital concept is therefore a key concept for visualizing the excitation process associated to a given transition. For example, one can say for a given transition that it corresponds to the excitation of an electron from an occupied orbital to a given unoccupied orbital. Nevertheless one has to keep in mind that electrons are fermions ruled by Pauli exclusion principle and cannot be distinguished from the other electrons in the atom. Moreover, it sometimes happens that the configuration interaction expansion converges very slowly and that one cannot speak about simple one-determinantal wave function at all. This is the case when electron correlation is large. Fundamentally, an atomic orbital is a one-electron wavefunction, even though most electrons do not exist in one-electron atoms, and so the one-electron view is an approximation. When thinking about orbitals, we are often given an orbital vision which (even if it is not spelled out) is heavily influenced by this Hartree – Fock approximation, which is one way to reduce the complexities of molecular orbital theory.
Connection to uncertainty relation Immediately after Heisenberg formulated his uncertainty relation, it was noted by Bohr that the existence of any sort of wave packet implies uncertainty in the wave frequency and wavelength, since a spread of frequencies is needed to create the packet itself. In quantum mechanics, where all particle momenta are associated with waves, it is the formation of such a wave packet which localizes the wave, and thus the particle, in space. In states where a quantum mechanical particle is bound, it must be localized as a wave packet, and the existence of the packet and its minimum size implies a spread and minimal value in particle wavelength, and thus also momentum and energy. In quantum mechanics, as a particle is localized to a smaller region in space, the associated compressed wave packet requires a larger and larger range of momenta, and thus larger kinetic energy. Thus, the binding energy to contain or trap a particle in a smaller region of space, increases without bound, as the region of space grows smaller. Particles cannot be restricted to a geometric point in space, since this would require an infinite particle momentum. In chemistry, Schrödinger, Pauling, Mulliken and others noted that the consequence of Heisenberg's relation was that the electron, as a wave packet, could not be considered to have an exact location in its orbital. Max Born suggested that the electron's position needed to be described by a probability distribution which was connected with finding the electron at some point in the wave-function which described its associated wave packet. The new quantum mechanics did not give exact results, but only the probabilities for the occurrence of a variety of possible such results. Heisenberg held that the path of a moving particle has no meaning if we cannot observe it, as we cannot with electrons in an atom. In the quantum picture of Heisenberg, Schrödinger and others, the Bohr atom number n for each orbital became known as an n-sphere in a three dimensional atom and was pictured as the mean energy of the probability cloud of the electron's wave packet which surrounded the atom. Although Heisenberg used infinite sets of positions for the electron in his matrices, this does not mean that the electron could be anywhere in the universe. Rather there are several laws that show the electron must be in one
Orbital
39
localized probability distribution. An electron is described by its energy in Bohr's atom which was carried over to matrix mechanics. Therefore, an electron in a certain n-sphere had to be within a certain range from the nucleus depending upon its energy. This restricts its location.
Hydrogen-like atoms The simplest atomic orbitals are those that occur in an atom with a single electron, such as the hydrogen atom. In this case the atomic orbitals are the eigenstates of the hydrogen Hamiltonian. They can be obtained analytically (see hydrogen atom). An atom of any other element ionized down to a single electron is very similar to hydrogen, and the orbitals take the same form. For atoms with two or more electrons, the governing equations can only be solved with the use of methods of iterative approximation. Orbitals of multi-electron atoms are qualitatively similar to those of hydrogen, and in the simplest models, they are taken to have the same form. For more rigorous and precise analysis, the numerical approximations must be used. A given (hydrogen-like) atomic orbital is identified by unique values of three quantum numbers: n, l, and ml. The rules restricting the values of the quantum numbers, and their energies (see below), explain the electron configuration of the atoms and the periodic table. The stationary states (quantum states) of the hydrogen-like atoms are its atomic orbital. However, in general, an electron's behavior is not fully described by a single orbital. Electron states are best represented by time-depending "mixtures" (linear combinations) of multiple orbitals. See Linear combination of atomic orbitals molecular orbital method. The quantum number n first appeared in the Bohr model where it determines the radius of each circular electron orbit. In modern quantum mechanics however, n determines the mean distance of the electron from the nucleus; all electrons with the same value of n lie at the same average distance. For this reason, orbitals with the same value of n are said to comprise a "shell". Orbitals with the same value of n and also the same value of l are even more closely related, and are said to comprise a "subshell".
Qualitative characterization Limitations on the quantum numbers An atomic orbital is uniquely identified by the values of the three quantum numbers, and each set of the three quantum numbers corresponds to exactly one orbital, but the quantum numbers only occur in certain combinations of values. The rules governing the possible values of the quantum numbers are as follows: The principal quantum number n is always a positive integer. In fact, it can be any positive integer, but for reasons discussed below, large numbers are seldom encountered. Each atom has, in general, many orbitals associated with each value of n; these orbitals together are sometimes called electron shells. is a non-negative integer. Within a shell where n is some integer n0, ranges across all (integer) values satisfying the relation . For instance, the n = 1 shell has only orbitals The azimuthal quantum number
with , and the n = 2 shell has only orbitals with , and particular value of are sometimes collectively called a subshell. The magnetic quantum number ranges thus:
. The set of orbitals associated with a
is also always an integer. Within a subshell where
is some integer
,
.
The above results may be summarized in the following table. Each cell represents a subshell, and lists the values of available in that subshell. Empty cells represent subshells that do not exist.
Orbital
40
1
2
3
4
2
0
-1, 0, 1
3
0
-1, 0, 1
-2, -1, 0, 1, 2
4
0
-1, 0, 1
-2, -1, 0, 1, 2 -3, -2, -1, 0, 1, 2, 3
5
0
-1, 0, 1
-2, -1, 0, 1, 2 -3, -2, -1, 0, 1, 2, 3 -4, -3, -2 -1, 0, 1, 2, 3, 4
...
...
...
...
...
...
...
...
Subshells are usually identified by their - and -values. is represented by its numerical value, but is represented by a letter as follows: 0 is represented by 's', 1 by 'p', 2 by 'd', 3 by 'f', and 4 by 'g'. For instance, one may speak of the subshell with and as a '2s subshell'.
The shapes of orbitals Any discussion of the shapes of electron orbitals is necessarily imprecise, because a given electron, regardless of which orbital it occupies, can at any moment be found at any distance from the nucleus and in any direction due to the uncertainty principle.
The shapes of the first five atomic orbitals: 1s, 2s, 2px,2py, and 2pz. The colors show the wavefunction phase.
However, the electron is much more likely to be found in certain regions of the atom than in others. Given this, a boundary surface can be drawn so that the electron has a high probability to be found anywhere within the surface, and all regions outside the surface have low values. The precise placement of the surface is arbitrary, but any reasonably compact determination must follow a pattern specified by the behavior of , the square of the wavefunction. This boundary surface is what is meant when the "shape" of an orbital is referred to. Generally speaking, the number determines the size and energy of the orbital for a given nucleus: as increases, the size of the orbital increases. However, in comparing different elements, the higher nuclear charge Z of heavier elements causes their orbitals to contract by comparison to lighter ones, so that the overall size of the whole atom remains very roughly constant, even as the number of electrons in heavier elements (higher Z) increases. Also in general terms, determines an orbital's shape, and its orientation. However, since some orbitals are described by equations in complex numbers, the shape sometimes depends on also. The single -orbitals ( ) are shaped like spheres. For n=1 the sphere is "solid" (it is most dense at the center and fades exponentially outwardly), but for n=2 or more, each single s-orbital is composed of spherically symmetric surfaces which are nested shells (i.e., the "wave-structure" is radial, following a sinusoidal radial component as well). The -orbitals for all n numbers are the only orbitals with an anti-node (a region of high wave function density) at the center of the nucleus. All other orbitals (p, d, f , etc.) have angular momentum, and thus avoid the nucleus (having a wave node at the nucleus). The three -orbitals for n=2 have the form of two ellipsoids with a point of tangency at the nucleus (sometimes referred to as a dumbbell). The three -orbitals in each shell are oriented at right angles to each other, as determined by their respective linear combination of values of .
Orbital
41
Four of the five
-orbitals for n=3 look similar, each with four pear-shaped balls, each ball tangent to two others,
and the centers of all four lying in one plane, between a pair of axes. Three of these planes are the -, -, and -planes, and the fourth has the centres on the and axes. The fifth and final -orbital consists of three regions of high probability density: a torus with two pear-shaped regions placed symmetrically on its axis. There are seven
-orbitals, each with shapes more complex than those of the
-orbitals.
For each s, p, d, f and g set of orbitals, the set of orbitals which composes it forms a spherically symmetrical set of shapes. For non-s orbitals, which have lobes, the lobes point in directions so as to fill space as symmetrically as possible for number of lobes which exist for a set of orientations. For example, the three p orbitals have six lobes which are oriented to each of the six primary directions of 3-D space; for the 5 d orbitals, there are a total of 18 lobes, in which again six point in primary directions, and the 12 additional lobes fill the 12 gaps which exist between each pairs of these 6 primary axes. Additionally, as is the case with the s orbitals, individual p, d, f and g orbitals with n values higher than the lowest possible value, exhibit an additional radial node structure which is reminiscent of harmonic waves of the same type, as compared with the lowest (or fundamental) mode of the wave. As with s orbitals, this phenomenon provides p, d, f, and g orbitals at the next higher possible value of n (for example, 3p orbitals vs. the fundamental 2p), an additional node in each lobe. Still higher values of n further increase the number of radial nodes, for each type of orbital. The shapes of atomic orbitals in one-electron atom are related to 3-dimensional spherical harmonics. These shapes are not unique, and any linear combination is valid, in fact it is possible to generate sets where all the d's are the same shape, just like the p x, p y, and p z are the same shape. [10] [11] Orbitals table
This table shows all orbital configurations for the real hydrogen-like wave functions up to 7 s, and therefore covers the simple electronic configuration for all elements in the periodic table up to radium. It is should also be noted that the p z orbital is the same as the p0 orbital, but the p x and p y are formed by taking linear compbinations of the p+1 and p-1 orbitals (which is why they are listed under the m=±1 label). Also, the p+1 and p-1 are not the same shape as the p0, since they are pure spherical harmonics. s (l=0)
p (l =1)
m=0
m=0
s
p
z
d (l =2)
m=±1 p
x
p
y
m=0 d 2 z
f (l =3)
m=±1 d
xz
d
yz
m=±2 d
xy
d 2 2 x -y
m=0
m=±1
f 3
f 2 f 2
z
m=±2 f
m=±3
xyz
f 2 2
f 2 2
f 2 2
...
...
...
...
...
xz
yz
...
...
z(x -y )
x(x -3y )
y(3x -y )
n=1 n=2 n=3 n=4
n=5 n=6 n=7
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
...
Orbital
42
Orbital energy In atoms with a single electron (hydrogen-like atoms), the energy of an orbital (and, consequently, of any electrons in the orbital) is determined exclusively by . The orbital has the lowest possible energy in the atom. Each successively higher value of has a higher level of energy, but the difference decreases as increases. For high , the level of energy becomes so high that the electron can easily escape from the atom. In atoms with multiple electrons, the energy of an electron depends not only on the intrinsic properties of its orbital, but also on its interactions with the other electrons. These interactions depend on the detail of its spatial probability distribution, and so the energy levels of orbitals depend not only on but also on . Higher values of are associated with higher values of energy; for instance, the 2 p state is higher than the 2s state. When = 2, the increase in energy of the orbital becomes so large as to push the energy of orbital above the energy of the s-orbital in the next higher shell; when = 3 the energy is pushed into the shell two steps higher. The energy sequence of the first 24 subshells is given in the following table. Each cell represents a subshell with and given by its row and column indices, respectively. The number in the cell is the subshell's position in the sequence.
1 1 2 2
3
3 4
5
7
4 6
8
10 13
5 9
11 14 17 21
6 12 15 18 22 26 7 16 19 23 27 31 8 20 24 28 32 36
Note: empty cells indicate non-existent sublevels, while numbers in italics indicate sublevels that could exist, but which do not hold electrons in any element currently known.
Electron placement and the periodic table Several rules govern the placement of electrons in orbitals (electron configuration). The first dictates that no two electrons in an atom may have the same set of values of quantum numbers (this is the Pauli exclusion principle). These quantum numbers include the three that define orbitals, as well as s, or spin quantum number. Thus, two electrons may occupy a single orbital, so long as they have different values of . However, only two electrons, because of their spin, can be associated with each orbital. Additionally, an electron always tends to fall to the lowest possible energy state. It is possible for it to occupy any orbital so long as it does not violate the Pauli exclusion principle, but if lower-energy orbitals are available, this condition is unstable. The electron will eventually lose energy (by releasing a photon) and drop into the lower orbital. Thus, electrons fill orbitals in the order specified by the energy sequence given above. This behavior is responsible for the structure of the periodic table. The table may be divided into several rows (called 'periods'), numbered starting with 1 at the top. The presently known elements occupy seven periods. If a certain period has number , it consists of elements whose outermost electrons fall in the th shell. The periodic table may also be divided into several numbered rectangular 'blocks'. The elements belonging to a given block have this common feature: their highest-energy electrons all belong to the same -state (but the associated with that -state depends upon the period). For instance, the leftmost two columns constitute the 's-block'. The
Orbital
43
outermost electrons of Li and Be respectively belong to the 2s subshell, and those of Na and Mg to the 3s subshell. The number of electrons in a neutral atom increases with the atomic number. The electrons in the outermost shell, or valence electrons, tend to be responsible for an element's chemical behavior. Elements that contain the same number of valence electrons can be grouped together and display similar chemical properties.
Relativistic effects For elements with high atomic number Z, the effects of relativity become more pronounced, and especially so for s electrons, which move at relativistic velocities as they penetrate the screening electrons near the core of high Z atoms. This relativistic increase in momentum for high speed electrons causes a corresponding decrease in wavelength and contraction of 6s orbitals relative to 5d orbitals (by comparison to corresponding s and d electrons in lighter elements in the same column of the periodic table); this results in 6s valence electrons becoming lowered in energy. Examples of significant physical outcomes of this effect include the lowered melting temperature of mercury (which results from 6s electrons not being available for metal bonding) and the golden color of gold and caesium (which results from narrowing of 6s to 5d transition energy to the point that visible light begins to be absorbed). See [12]. In the Bohr Model, an electron has a velocity given by , where Z is the atomic number, is the fine-structure constant, and c is the speed of light. In non-relativistic quantum mechanics, therefore, any atom with an atomic number greater than 137 would require its 1s electrons to be traveling faster than the speed of light. Even in the Dirac equation, which accounts for relativistic effects, the wavefunction of the electron for atoms with Z > 137 is oscillatory and unbound. The significance of element 137, also known as untriseptium, was first pointed out by the physicist Richard Feynman. Element 137 is sometimes informally called feynmanium (symbol Fy). However, Feynman's approximation fails to predict the exact critical value of Z due to the non-point-charge nature of the nucleus and very small orbital radius of inner electrons, resulting in a potential seen by inner electrons which is effectively less than Z. The critical Z value which makes the atom unstable with regard to high-field breakdown of the vacuum and production of electron-positron pairs, does not occur until Z is about 173. These conditions are not seen except transiently in collisions of very heavy nuclei such as lead or uranium in accelerators, where such electron-positron production from these effects has been claimed to be observed. See Extension of the periodic table beyond the seventh period.
See also • • • • • •
Atomic electron configuration table Electron configuration Energy level List of Hund's rules Molecular orbital Quantum chemistry computer programs
Orbital
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Further reading • Tipler, Paul; Ralph Llewellyn (2003). Modern Physics (4 ed.). New York: W. H. Freeman and Company. ISBN 0-7167-4345-0. • Scerri, Eric (2007). The Periodic Table, Its Story and Its Significance. New York: Oxford University Press. ISBN 978-0-19-530573-9.
External links • • • • • • • •
Guide to atomic orbitals [13] Covalent Bonds and Molecular Structure [14] Animation of the time evolution of an hydrogenic orbital [15] The Orbitron [16], a visualization of all common and uncommon atomic orbitals, from 1s to 7g Grand table [17] Still images of many orbitals David Manthey's Orbital Viewer [18] renders orbitals with n ≤ 30 Java orbital viewer applet [19] What does an atom look like? Orbitals in 3D [20]
• Atom Orbitals v.1.5 visualization software [21]
References [1] Milton Orchin,Roger S. Macomber, Allan Pinhas, and R. Marshall Wilson(2005)" Atomic Orbital Theory (http://media.wiley.com/ product_data/excerpt/81/04716802/0471680281.pdf)" [2] Daintith, J. (2004). Oxford Dictionary of Chemistry . New York: Oxford University Press. I SBN 0-19-860918-3. [3] The Feynman Lectures on Physics -The Definitive Edition, Vol 1 lect 6 pg 11. Feynman, Richard; Leighton; Sands. (2006) Addison Wesley ISBN 0-8053-9046-4 [4] Bohr, Niels (1913). "On the Constitution of Atoms and Molecules". Philosophical Magazine 26 (1): 476. [5] Nagaoka, Hantaro (May 1904). "Kinetics of a System of Particles illustrating the Line and the Band Spectrum and the Phenomena of Radioactivity" (http://www.chemteam.info/Chem-History/Nagaoka-1904.html). Philosophical Magazine 7: 445 – 455. . [6] Bryson, Bill (2003). A Short History of Nearly Everything . Broadway Books. pp. 141 – 143. ISBN 0-7679-0818-X. [7] Mulliken, Robert S. (July 1932). "Electronic Structures of Polyatomic Molecules and Valence. II. General Considerations" (http://prola.aps. org/abstract/PR/v41/i1/p49_1). Phys. Rev. 41 (1): 49 – 71. doi:10.1103/PhysRev.41.49. . [8] Griffiths, David (1995). Introduction to Quantum Mechanics. Prentice Hall. pp. 190 – 191. ISBN 0-13-124405-1. [9] Levine, Ira (2000). Quantum Chemistry (5 ed.). Prentice Hall. pp. 144 – 145. ISBN 0-13-685512-1. [10] Powell, Richard E. (1968). "The five equivalent d orbitals". Journal of Chemical Education 45: 45. doi:10.1021/ed045p45. [11] Kimball, George E. (1940). "Directed Valence". The Journal of Chemical Physics 8: 188. doi:10.1063/1.1750628. [12] http://www.chem1.com/acad/webtut/atomic/qprimer/#Q26 [13] http://www.chemguide.co.uk/atoms/properties/atomorbs.html [14] http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_07.html [15] http://strangepaths.com/atomic-orbital/2008/04/20/en/ [16] http://www.shef.ac.uk/chemistry/orbitron/ [17] http://www.orbitals.com/orb/orbtable.htm [18] http://www.orbitals.com/orb/index.html [19] http://www.falstad.com/qmatom/ [20] http://www.hydrogenlab.de/elektronium/HTML/einleitung_hauptseite_uk.html [21] http://taras-zavedy.narod.ru/PROGRAMMS/ATOM_ORBITALS_v_1_5_ENG/Atom_Orbitals_v_1_5_ENG.html
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Groups Group In chemistry, a group (also known as a family) is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table. The modern explanation of the pattern of the table is that the elements in a group have similar configurations of the outermost electron shells of their atoms: as most chemical properties are dominated by the orbital location of the outermost electron. There are three The periodic table of the chemical elements conventional ways of numbering: One using Arabic numerals, and two using Roman numerals. The Roman numeral names are the original traditional names of the groups; the Arabic numeral names are those recommended by the International Union of Pure and Applied Chemistry (IUPAC) to replace the old names in an attempt to reduce the confusion generated by the two older, but mutually confusing, schemes. There is considerable confusion surrounding the two old systems in use (old IUPAC and CAS) that combined the use of Roman numerals with letters. In the old IUPAC system the letters A and B were designated to the left (A) and right (B) part of the table, while in the CAS system the letters A and B were designated to main group elements (A) and transition elements (B). The old IUPAC system was frequently used in Europe while the CAS was most common in America. The new IUPAC scheme was developed to replace both systems as they confusingly used the same names to mean different things. The IUPAC proposal was first circulated in 1985 for public comments, [1] and was later included as part of the 1990 edition of the Nomenclature of Inorganic Chemistry .[2] The periodic table groups are as follows (in the brackets are shown the old systems: European and American): • • • • • • • • • • • • • •
Group 1 (IA,IA): the alkali metals or lithium family Group 2 (IIA,IIA): the alkaline earth metals or beryllium family Group 3 (IIIA,IIIB): the scandium family Group 4 (IVA,IVB): the titanium family Group 5 (VA,VB): the vanadium family Group 6 (VIA,VIB): the chromium family Group 7 (VIIA,VIIB): the manganese family Group 8 (VIII, VIIIB): the iron family Group 9 (VIII, VIIIB): the cobalt family Group 10 (VIII, VIIIB): the nickel family Group 11 (IB,IB): the coinage metals (not an IUPAC-recommended name) or copper family Group 12 (IIB,IIB): the zinc family Group 13 (IIIB,IIIA): the boron family Group 14 (IVB,IVA): the carbon family
Group • • • •
46 Group 15 (VB,VA): the pnictogens or nitrogen family Group 16 (VIB,VIA): the chalcogens or oxygen family Group 17 (VIIB,VIIA): the halogens or fluorine family Group 18 (Group 0, VIIIA): the helium family/neon family; for the first six periods, these are the noble gases
References [1] Fluck, E. New notations in the periodic table. Pure & App. Chem. 1988, 60, 431-436. (http://www.iupac.org/publications/pac/1988/pdf/ 6003x0431. pdf) [2] Leigh, G. J. Nomenclature of Inorganic Chemistry: Recommendations 1990 . Blackwell Science, 1990. ISBN 0632024941.
3. Scerri, E. R. The Periodic Table, Its Story and Its Significance, Oxford University Press, 2007. ISBN 978-0-19-530573-9.
Group I The alkali metals are a series of chemical elements forming Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals). The alkali metals provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behavior down the group.
Properties The alkali metals are all highly reactive and are never found in elemental forms in nature. Because of this, they are usually stored immersed in mineral oil or kerosene (paraffin oil). They also tarnish easily and have low melting points and densities. Physically, the alkali metals are mostly silver-colored, except for metallic caesium, which has a golden tint. These elements are all soft metals of low density. Chemically, all of the alkali metals react aggressively with the halogens to form ionic salts. They all react with water to form strongly alkaline hydroxides. The vigor of reaction increases down the group. All of the atoms of alkali metals have one electron in their outmost electron shells, hence their only way for achieving the equivalent of filled outmost electron shells is to give up one electron to an element with high electronegativity, and hence to become singly charged positive ions, i.e. cations. When it comes to their nuclear physics, the elements potassium and rubidium are naturally weakly radioactive because they each contain a long half-life radioactive isotope. The element hydrogen, with its solitary one electron per atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not counted as an alkali metal. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule. The removal of the single electron of hydrogen requires considerably more energy than removal of the outer electron from the atoms of the alkali metals. As in the halogens, only one additional electron is required to fill in the outermost shell of the hydrogen atom, so hydrogen can in some circumstances behave like a halogen, forming the negative hydride ion. Binary compounds of hydrogen with the alkali metals and some transition metals have been produced in the laboratory, but these are only laboratory curiosities without much practical use. Under extremely high pressures and low temperatures, such as those found at the cores of the planets Jupiter and Saturn, hydrogen does become a metallic element, and it behaves like an alkali metal. (See metallic hydrogen.) The alkali metals have the lowest ionization potentials in their periods of the periodic table, because the removal of their single electrons from their outmost electron shells gives them the stable electron configuration of inert gases. Another way of stating this is that they all have a high electropositivity. The "second ionization potential" of all of
Group I
47
the alkali metals is very high, since removing any electron from an atom having a noble gas configuration is difficult to do. All of the alkali metals are notable for their vigorous reactions with water, and these reactions become increasingly vigorous when going down their column in the periodic table towards the heaviest alkali metals, such as caesium. Their chemical reactions with water are as follows: Alkali metal + water → Alkali metal hydroxide + hydrogen gas For a typical example (M represents an alkali metal): 2 M (s) + 2 H O (l) → 2 MOH (aq) + H (g) 2
2
Series of alkali metals, stored in mineral oil ("natrium" is sodium.)
Trends Like in other columns of the periodic table, the members of the alkali metal family show patterns in their electron configurations, especially their outmost electron shells. This causes similar patterns in their chemical properties: Z
Element
1
Hydrogen 1
3
Lithium
11 Sodium
No. of electrons/shell
2, 1 2, 8, 1
19 Potassium 2, 8, 8, 1 37 Rubidium 2, 8, 18, 8, 1 55 Caesium
2, 8, 18, 18, 8, 1
87 Francium 2, 8, 18, 32, 18, 8, 1
The alkali metals show a number of trends when moving down the group - for instance: decreasing electronegativity, increasing reactivity, and decreasing melting and boiling point. Their densities generally increase, with the notable exception that potassium is less dense than sodium, and the possible exception of francium being less dense than caesium. (The highly radioactive element francium only exists in microscopic quantities.) Alkali metal Standard atomic weight (u) Melting point (K) Boiling point (K)
Density
Electronegativity (Pauling)
−3
(g·cm )
Lithium
6.941
453
1615
0.534
0.98
Sodium
22.990
370
1156
0.968
0.93
Potassium
39.098
336
1032
0.89
0.82
Rubidium
85.468
312
961
1.532
0.82
Caesium
132.905
301
944
1.93
0.79
Francium
(223)
295
950
1.87
0.70
Group I
48
Compounds Alkali metals form a very wide range of amalgams. [1] They tend to form ionically bonded salts with most electronegative elements on the periodic table, like cesium fluoride and sodium chloride.
See also • • • • • • •
alkaline earth metal Lithium Sodium Potassium Rubidium Caesium Francium
References [1] Deiseroth, H (1997). "Alkali metal amalgams, a group of unusual alloys". Progress in Solid State Chemistry 25: 73. doi:10.1016/S0079-6786(97)81004-7.
• Campbell, Linda M., Aaron T. Fisk, Xianowa Wang, Gunter Kock, and Derek C. Muir (2005). "Evidence for Biomagnification of Rubidium in Freshwater and Marine Food Webs". Canadian Journal of Fisheries and Aquatic Sciences 62: 1161 – 1167. doi:10.1139/f05-027. • Chang, Cheng-Hung, and Tian Y. Tsong (2005). "Stochastic Resonance of Na, K-Ion Pumps on the Red Cell Membrane". AIP Conference Proceedings: 18th International Conference on Noise and Fluctuations . 780. American Institute of Physics. pp. 587 – 590. doi:10.1063/1.2036821. ISBN 0-7354-0267-1. • Joffe, Russell T., Stephen T. Sokolov and Anthony J. Levitt (2006). "Lithium and Triiodothyronine Augmentation of Antidepressants". Canadian Journal of Psychiatry 51 (12): 791 – 793. PMID 17168254. • Bauer, Brent A., Robert Houlihan, Michael J. Ackerman, Katya Johnson, and Himeshkumar Vyas (2006). "Acquired Long QT Syndrome Secondary to Cesium Chloride Supplement". Journal of Alternative and Complementary Medicine 12 (10): 1011 – 1014. doi:10.1089/acm.2006.12.1011. PMID 17212573. • Erermis, Serpil, Muge Tamar, Hatice Karasoy, Tezan Bildik, Eyup S. Ercan, and Ahmet Gockay (1997). "Double-Blind Randomised Trial of Modest Salt Restriction in Older People". Lancet 350: 850 – 854. doi:10.1016/S0140-6736(97)02264-2. • Krachler, M, and E Rossipal (1999). "Trace Elements Transfer From Mother to the Newborn - Investigations on Triplets of Colostrum, Maternal and Umbilical Sera". European Journal of Clinical Nutrition 53 (6): 486 – 494. doi:10.1038/sj.ejcn.1600781. PMID 10403586. • Stein, Benjamin P., Stephen G. Benka, Phillip F. Schewe, and Bertram Schwarzhild (1996). "Physics Update". Physics Today 49 (6): 9. doi:10.1063/1.2807642. • "Group 1: The Alkali Metals" (http://www.chemsoc.org/Viselements/pages/data/intro_groupi_data.html). Visual Elements. Royal Society of Chemistry. Retrieved 2009-12-08.
Group I
49
External links • Science aid:Alkali metals (http://www.scienceaid.co.uk/chemistry/periodictable/alkalimetals.html) A simple look at alkali metals • Atomic and Physical Properties of the Group 1 Elements (http://www.chemguide.co.uk/inorganic/group1/ properties.html) An in-depth look at alkali metals
Explanation of above periodic table slice: Alkali metals
Atomic numbers in black are solids
Solid borders indicate primordial elements (older than the Earth)
Dashed borders indicate natural radioactive elements with no isotopes older than the Earth
Group II Group →
2
↓ Period
2
4 Be
3
12 Mg
4
20 Ca
5
38 Sr
6
56 Ba
7
88 Ra
The alkaline earth metals are a series of elements comprising Group 2 (IUPAC style) (Group IIA) of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). This specific group in the periodic table owes its name to their oxides that simply give basic alkaline solutions. These oxides melt at such high temperature that they remain solids ( “earths”) in fires. The alkaline earth metals provide a good example of group trends in properties in the periodic table, with well-characterized homologous behavior down the group. With the exception of Be and Mg, the metals have a distinguishable flame color, brick-red for Ca, magenta-red for Sr, green for Ba and crimson red for Ra. Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:
Group II
50
Z
4
Element
Beryllium
No. of elect rons/shell
2, 2
12 Magnesium 2, 8, 2 20 Calcium
2, 8, 8, 2
38 Strontium
2, 8, 18, 8, 2
56 Barium
2, 8, 18, 18, 8, 2
88 Radium
2, 8, 18, 32, 18, 8, 2
The alkaline earth metals are silver colored, soft metals, which react readily with halogens to form ionic salts, and with water, though not as rapidly as the alkali metals, to form strong alkaline (basic) hydroxides. For example, where sodium and potassium react with water at room temperature, magnesium reacts only with steam and calcium with hot water: Mg + 2 H2O → Mg(OH)2 + H2 Beryllium is an exception: It does not react with water or steam, and its halides are covalent. All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions. The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These oxides are basic (alkaline) when combined with water. "Earth" is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating —properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie ( Elements of Chemistry ) of 1789 he called them salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths.
Biological occurrences • Beryllium's low aqueous solubility means it is rarely available to biological systems; it has no known role in living organisms, and when encountered by them, is generally highly toxic. • Magnesium and calcium are ubiquitous and essential to all known living organisms. They are involved in more than one role, with, for example, Mg/Ca ion pumps playing a role in some cellular processes, magnesium functioning as the active center in some enzymes, and calcium salts taking a structural role (e.g. bones).
The alkaline earth metals.
• Strontium and barium have a lower availability in the biosphere. Strontium plays an important role in marine aquatic life, especially hard corals. They use strontium to build their exoskeleton. These elements have some uses in medicine, for example "barium meals" in radio graphic imaging, whilst strontium compounds are employed in
Group II
51
some toothpastes. • Radium has a low availability and is highly radioactive, making it toxic to life.
Comparison of the physical properties and chemical properties between alkaline earth and alkali metals 2
Just like their names, they do not differ completely. The main difference is the electron configuration, which is ns for alkaline earth metals and ns1 for alkali metals. For the alkaline earth metals, there are two electrons that are available to form a metallic bond, and the nucleus contains an additional positive charge. Also, the elements of group 2A (alkaline earth) have much higher melting points and boiling points compared to those of group 1A (alkali metals). The alkali also have a softer and more lighweight figure whereas the alkaline earth metals are much harder and denser. The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the alkali metals. The second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains much higher ionization energy, they still form an ionic compound with 2+ cations. Beryllium, however, behaves differently. This is because in order to remove two electrons from this particular atom, it requires significantly more energy. It never forms Be2+ and its bonds are polar covalent.
Beryllium As mentioned earlier, Be is “special”; it behaves differently. If the Be2+ ion did exist, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since Be has a high charge density. All compounds that include Be have a covalent bond. Even the compound BeF 2 , which is the most ionic Be compound, has a low melting point and a low electrical conductivity when melted.
Important reactions and compounds Reactions:
Note: E = elements that act as reducing agents 1. The metals reduce halogens to form ionic halides: E(s) + X2 → EX2 (s) where X = F, Cl, Br or I 2. The metals reduce O2 to form the oxides: 2E(s) + O2 → 2EO(s) 3. The larger metals react with water to produce hydrogen gas: E + 2H O → E2+ + 2OH- + H where E = (s) 2 (l) (aq) aq 2 (g) Ca, Sr or Ba Compounds
1. Alkylmagnesium halides (RMgX where R = hydrocarbon group and X = halogen). They are used to synthetise organic compounds. Here’s an example: 3RMgCl + SnCl4 → 3MgCl2 + R3SnCl 2. Magnesium oxide (MgO). It is used as a material to refract furnace brick and wire insulation (melting point of 2852°C). 3. Calcium carbonate (CaCO ). It is mainly used in the construction industry and for making limestone, marble, 3 chalk, and coral.
Group II
52
References Group 2 - Alkaline Earth Metals [1], Royal Chemistry Society. Group 1 Alkali Metals and Group 2 Alkaline Earth Metals [2], Doc Brown's Chemistry Clinic. Science aid: Group 2 Metals [3] Study aid for teens Maguire, Michael E. "Alkaline Earth Metals." Chemistry: Foundations and Applications. Ed. J. J. Lagowski. Vol. 1. New York: Macmillan Reference USA, 2004. 33-34. 4 vols. Gale Virtual Reference Library. Thomson Gale. • Silberberg, M.S., Chemistry: The molecular nature of Matter and Change (3e édition, McGraw-Hill 2009) • Petrucci R.H., Harwood W.S. et Herring F.G., General Chemistry (8e édition, Prentice-Hall 2002) • • • •
Explanation of above periodic table slice: Alkaline earth metals
Atomic numbers in black indicate solids
Solid borders indicate primordial elements (older than the Earth)
References [1] http://www.chemsoc.org/visElements/pages/data/intro_groupii_data.html [2] http://www.wpbschoolhouse.btinternet.co.uk/page07/sblock.htm [3] http://scienceaid.co.uk/chemistry/fundamental/group2.html
Group III Period
Group 3
4
21 Sc 5
39 Y Group 3/ungrouped
* Lanthanides
** Actinides 6
*Lanthanides
Dashed borders indicate natural radioactive elements with no isotopes older than the Earth
Group III
53
57 58 59 60 La Ce Pr Nd 61 62 63 64 Pm Sm Eu Gd 65 66 67 68 Tb Dy Ho Er 69 70 71 Tm Yb Lu 7
**Actinides
89 90 Ac Th
91 Pa
92 U
93 94 95 96 Np Pu Am Cm 97 98 99 100 Bk Cf Es Fm 101 102 103 Md No Lr
The Group 3 elements are chemical elements comprising the third vertical column of the periodic table. IUPAC has not recommended a specific format for the periodic table, so different conventions are permitted and are often used for group 3. The following d-block transition metals are always considered members of group 3: • scandium (Sc) • yttrium (Y)
The metals
Scandium
Yttrium
When defining the remainder of group 3, four different conventions may be encountered: • Some tables [1] include lanthanum (La) and actinium (Ac), (the beginnings of the lanthanide and actinide series of elements, respectively) as the remaining members of group 3. In their most commonly encountered tripositive ion forms, these elements do not possess any partially filled f orbitals, thus resulting in more d-block-like behavior. • Some tables [2] include lutetium (Lu) and lawrencium (Lr) as the remaining members of group 3. These elements terminate the lanthanide and actinide series, respectively. Since the f-shell is nominally full in the ground state electron configuration for both of these metals, they behave most like d-block metals out of all the lanthanides
Group III
54
and actinides, and thus exhibit the most similarities in properties with Sc and Y. For Lr, this behavior is expected, but it has not been observed because sufficient quantities are not available. (See also Periodic table (wide) and Periodic table (extended). ) Some tables [3] refer to all lanthanides and actinides by a marker in group 3. A third and fourth alternative are suggested by this arrangement: • The third alternative is to regard all 30 lanthanide and actinide elements as included in group 3. Lanthanides, as electropositive trivalent metals, all have a closely related chemistry, and all show many similarities to Sc and Y. • The fourth alternative is to include none of the lanthanides and actinides in group 3. The lanthanides possess additional properties characteristic of their partially-filled f orbitals which are not common to Sc and Y. Furthermore, the actinides show a much wider variety of chemistry (for instance, in range of oxidation states) within their series than the lanthanides, and comparisons to Sc and Y are even less useful. The term rare earth elements is often used for group 3 elements including the lanthanides but excluding the actinides.
Occurrence Scandium, yttrium, and the lanthanides (except promethium) tend to occur together in the Earth's crust, and are relatively abundant compared with most d-block metals, but often harder to extract from their ores.
Biological chemistry Group 3 elements are generally hard metals with low aqueous solubility, and have low availability to the biosphere. No group 3 has any documented biological role in living organisms. The radioactivity of the actinides generally makes them highly toxic to living cells.
Notes Explanation of above periodic table slice:
Transition metals
Lanthanide series
Actinide series
Atomic numbers in black indicate solids
Solid borders indicate primordial elements (older than the Earth)
Dashed borders Dotted borders indicate natural indicate synthetic radioactive elements elements
References [1] Periodic table at Lanl.gov (http://periodic.lanl.gov/) [2] "WebElements Periodic Table of the Elements" (http://www.webelements.com). Webelements.com. . Retrieved 2010-04-03. [3] "International Union of Pure and Applied Chemistry > Periodic Table of the Elements" (http://www.iupac.org/reports/periodic_table/ index.html). Iupac.org. . Retrieved 2010-04-03.
Group IV
55
Group IV Group →
4
↓ Period
4
22 Ti 5
40 Zr 6
72 Hf 104 Rf
7
Legend
Transition metal Primordial element Synthetic
The Group 4 elements are a group of chemical elements in the periodic table. In the modern IUPAC nomenclature, Group 4 of the periodic table contains titanium (Ti), zirconium (Zr), hafnium (Hf ) and rutherfordium (Rf ). This group lies in the d-block of the periodic table. The group itself has not acquired a trivial name; it belongs to the broader grouping of the transition metals. The three Group 4 elements that occur naturally are titanium (Ti), zirconium (Zr) and hafnium (Hf). The first three members of the group share similar properties; all three are hard refractory metals under standard conditions. However the fourth element rutherfordium (Rf), has been synthesized in the laboratory, none of them have been found occurring in nature. All isotopes of rutherfordium are radioactive. So far, no experiments in a supercollider were conducted to synthesize the next member of the group Unpentquadium (Upq). As Upq is a late member of period 8 element it is unlikely that this element will be synthesized in the near future.
Characteristics Chemistry Most of the chemistry has been observed only for the first three members of the group, the chemistry of rutherfordium is not very established and therefore the rest of the section deals only with titanium, zirconium, and hafnium. All the elements of the group are reactive metals with a high melting point (1668 °C, 1855 °C, 2233 °C). The reactivity is not always obvious due to the rapid formation of a stable oxide layer, which prevents further reactions. The oxides TiO2, ZrO2 and HfO2 are white solids with high melting points and unreactive against most acids.[1]
Group IV
56
H Li
He Be
B
C
N
O
F
Na Mg
Al
Si
P
S
Cl Ar
As
Se Br Kr
K
Ca Sc
Ti
Rb
Sr
Zr Nb Mo Tc
Ru Rh
Pd Ag Cd In
Sn
Sb
Te
Hf Ta
Pb
Bi
Po At Rn
Y
Cs Ba La Fr
*
V
Cr Mn Fe
Co
Ni
Cu Zn Ga Ge
Ne
W
Re
Os
Ir
Pt
Au Hg Tl
Ra Ac ** Rf Db Sg
Bh
Hs
Mt
Ds
Rg Cn
*
Ce Pr Nd Pm Sm Eu
** Th Pa
U
I
Xe
Gd Tb Dy Ho Er Tm Yb Lu
Np Pu Am Cm Bk Cf
Es Fm Md No Lr
Group 4 in the periodic table
As tetravalent transition metals, all three elements form various inorganic compounds, generally in the oxidation state of +4. For the first three metals, it has been shown that they are resistant to concentrated alkalis, but halogens react with them to form tetrahalides. At higher temperatures, all three metals react with oxygen, nitrogen, carbon, boron, sulfur, and silicon. Because of the lanthanide contraction of the elements in the fifth period, zirconium and hafnium have nearly identical ionic radii. The ionic radius of Zr4+ is 79 picometers and that of Hf 4+ is 78 pm.[1] [2] This similarity results in nearly identical chemical behavior and in the formation of similar chemical compounds.[2] The chemistry of hafnium is so similar to that of zirconium that a separation on chemical reactions was not possible, only the physical properties of the compounds differ. The melting points and boiling points of the compounds and the solubility in solvents are the major differences in the chemistry of these twin elements.[1]
Physical Properties of the Group 4 element Name
Titanium
Zirconium
Hafnium
Rutherfordium
Melting point
1941 K (1668 °C)
2130 K (1857 °C)
2506 K (2233 °C)
?
Boiling point
3560 K (3287 °C)
4682 K (4409 °C)
4876 K (4603 °C)
?
Density
4.507 g·cm−3
6.511 g·cm−3
13.31 g·cm−3
?
Appearance
silver metallic
silver white
silver gray
?
140 pm
155 pm
155 pm
?
Atomic radius
Group IV
57
History While titanium and zirconium, as relatively abundant elements, were discovered in the late 18th century, it took until 1923 for hafnium to be identified. This was only partly due to hafnium's relative scarcity. The chemical similarity between zirconium and hafnium made a separation difficult and, without knowing what to look for, hafnium was left undiscovered, although all samples of zirconium, and all of its compounds, used by chemists for over two centuries contained significant amounts of hafnium.[3] William Gregor, Franz Joseph Muller and Martin Heinrich Klaproth Crystal of the abundant mineral Ilmenite independently discovered titanium between 1791 and 1795. Klaproth named it for the Titans of Greek mythology.[4] Klaproth also discovered zirconium in the mineral zircon in 1789 and named it after the already known Zirkonerde (zirconia). Hafnium had been predicted by Dmitri Mendeleev in 1869 and Henry Moseley measured in 1914 the effective nuclear charge by X-ray spectroscopy to be 72, placing it between the already known elements lanthanum and tantalum. Dirk Coster and Georg von Hevesy were the first to search for the new element in zirconium ores.[5] Hafnium was discovered by the two in 1923 in Copenhagen, Denmark, validating the original 1869 prediction of Mendeleev.[6] Rutherfordium was reportedly first detected in 1966 at the Joint Institute of Nuclear Research at Dubna (then in the Soviet Union). Researchers there bombarded 242Pu with accelerated 22Ne ions and separated the reaction products by gradient thermochromatography after conversion to chlorides by interaction with ZrCl4.[7] 242 Pu 94
− x
− x
+ 2210Ne → 264 104Rf → 264 104RfCl4 The next element after Rutherfordium is expected to be Unpentquadium (Upq). There are no plans to attempt to synthesize the next element in the near future, since it is a late member of the Period 8 elements. Currently none of the period 8 elements have been discovered yet. It is also probable that, due to drip instabilities, only the lower Period 8 elements are physically possible.
Production The production of the metals itself is difficult due to their reactivity. The formation of oxides, nitrides and carbides must be avoided to yield workable metals, this is normally achieved by the Kroll process. The oxides (MO2) are reacted with coal and chlorine to form the chlorides (MCl4).The chlorides of the metals are than reacted with magnesium yielding magnesium chloride and the metals. Further purification is done by a chemical transport reaction developed by Anton Eduard van Arkel and Jan Hendrik de Boer. In a closed vessel, the metal reacts with iodine at temperatures of above 500 °C forming metal(IV) iodide; at a tungsten filament of nearly 2000 °C the reverse reaction happens and the iodine and metal are set free. The metal forms a solid coating at the tungsten filament and the iodine can react with additional metal resulting in a steady turn over.[1] [8] M + 2 I2 (low temp.) → MI4 MI4 (high temp.) → M + 2 I2
Group IV
Occurrence If the abundance of elements in Earth's crust is compared for titanium, zirconium and hafnium, the abundance decreases with increase of atomic mass. Titanium is the seventh most abundant metal in Earth's crust and has an abundance of 6320 ppm, while zirconium has an abundance of 162 ppm and hafnium has only an abundance of 3 ppm.[9] All three stable elements occur in heavy mineral sands ore deposits, which are placer deposits formed, most usually in beach environments, Heavy minerals (dark) in a quartz beach sand by concentration due to the specific gravity of the mineral grains of (Chennai, India). erosion material from mafic and ultramafic rock. The titanium minerals are mostly anatase and rutile, and zirconium occurs in the mineral zircon. Because of the chemical similarity, up to 5% of the zirconium in zircon is replaced by hafnium. The largest producers of the group 4 elements are Australia, South Africa and Canada.[10] [11] [12] [13] [14]
Applications Titanium metal and its alloys have a wide range of applications, where the corrosion resistance, the heat stability and the low density (light weight) are of benefit. The foremost use of corrosion-resistant hafnium and zirconium has been in nuclear reactors. Zirconium has a very low and hafnium has a high thermal neutron-capture cross-section. Therefore, zirconium (mostly as zircaloy) is used as cladding of fuel rods in nuclear reactors,[15] while hafnium is used as control rod for nuclear reactors, because each hafnium atom can absorb multiple neutrons.[16] [17] Smaller amounts of hafnium[18] and zirconium are used in supper alloys to improve the properties of those alloys.[19]
Biological occurrences occurrences The group 4 elements are not known to be involved in the biological chemistry of any living systems. [20] They are hard refractory metals with low aqueous solubility and low availability to the biosphere. Titanium is one of the few first row d-block transition metals with no known biological role. Rutherfordium's radioactivity would make it toxic to living cells.
Precaution Titanium is non-toxic even in large doses and does not play any natural role inside the human body. [20] Zirconium powder can cause irritation, but only contact with the eyes requires medical attention. [21] OSHA recommends for zirconium are 5 mg/m3 time weighted average limit and a 10 mg/m3 short-term exposure limit.[22] Only limited data exists on the toxicology of hafnium.[23]
58
Group IV
References [1] Holleman, Arnold Arnold F.; Wiberg, Wiberg, Egon; Egon; Wiberg, Nils; (1985) (1985) (in German). German). Lehrbuch der Anorganischen Chemie (91-100 ed.). Walter de Gruyter. pp. 1056 – 1057. 1057. ISBN 3110075113. [2] "Los Alamos Alamos National National Laborat Laboratory ory – Hafnium" (http://periodic.lanl.gov/elements/72. (http://periodic.lanl.gov/elements/72.html). html). . Retrieved 2008-09-10. [3] Barksdale, Jelks (1968). (1968). The Encyclopedia Encyclopedia of the Chemical Elements. Elements. Skokie, Illinois: Illinois: Reinhold Book Book Corporation. pp. pp. 732 – 38 38 "Titanium". LCCCN 68-29938. Journal of Chemical Education : 1231 – 1243. [4] Weeks, Mary Elvira (1932). (1932). "III. Some Eighteenth-Century Eighteenth-Century Metals". Journal 1243. [5] Urbain, M. G. (1922). "Sur "Sur les séries L du lutécium et de l'ytterbium et sur l'identification l'identification d'un celtium celtium avec l'élément de nombre nombre atomique 72 (The L series from lutetium to ytterbium and the identification of element 72 celtium" (http://gallica.bnf. (http://gallica.bnf.fr/ark:/12148/bpt6k3127j/f1348. fr/ark:/12148/bpt6k3127j/f1348. table) (in French). Comptes rendus 174: 1347 – 1349. 1349. . Retrieved 2008-10-30. Nature 111: 79 – 79. [6] Coster, D.; Hevesy, Hevesy, G. (1923-01-20). (1923-01-20). "On the Missing Missing Element Element of Atomic Number 72". Nature 79. doi:10.1038/111079a0. [7] Barber, R. C.; Greenwood, N. N.; Hrynkiewicz, Hrynkiewicz, A. Z.; Jeannin, Y. P.; Lefort, Lefort, M.; Sakai, M.; Ulehla, I.; Wapstra, Wapstra, A. P.; Wilkinson, D. H. (1993). "Discovery of the transfermium elements. Part II: Introduction to discovery profiles. Part III : Discovery profiles of the transfermium elements" (http://www.iupac.org/publications/pac/65/8/1757/). (http://www.iupac.org/publications/pac/65/8/1757/). Pure and Applied Chemistry 65 (8): 1757 – 1814. 1814. doi:10.1351/pac199365081757. . [8] van Arkel, A. E.; de Boer, J. H. (1925). (1925). "Darstellung von von reinem Titanium-, Zirkonium-, Zirkonium-, HafniumHafnium- und Thoriummetall Thoriummetall (Production of pure pure titanium, zirconium, hafnium and Thorium metal)" (in German). Zeitschrift für anorganische und allgemeine Chemie 148 (1): 345 – 350. 350. doi:10.1002/zaac.19251480133. [9] "Abundan "Abundance ce in Earth's Earth's Crust" (http://www.webe (http://www.webelemen lements. ts.com/periodicity/abundan com/periodicity/abundance_crust/). ce_crust/). WebElements.com. WebElements.com. . Retrieved 2007-04-14. [10] "Dubbo Zirconia Project Fact Sheet" (http://www.alkane.com.au/ (http://www.alkane.com.au/ projects/nsw/dubbo/DZP Summary June07.pdf) June07.pdf) (PDF). Alkane Resources Limited. June 2007. . Retrieved 2008-09-10. [11] "Zirconiu "Zirconium m and Hafnium" Hafnium" (http://miner (http://minerals.us als.usgs. gs.gov/minerals/pubs/commodity/zirconium/mcs-2008-zirco. gov/minerals/pubs/commodity/zirconium/mcs-2008-zirco.pdf) pdf) (PDF). Mineral Commodity Summaries (US Geological Survey): 192 – 193. 193. January 2008. . Retrieved 2008-02-24. [12] Callaghan, R. (2008-02-21). (2008-02-21). "Zirconium "Zirconium and Hafnium Statistics and Information" Information" (http://minerals.usgs. (http://minerals.usgs.gov/minerals/pubs/commodity/ gov/minerals/pubs/commodity/ zirconium/). US Geological Survey. . Retrieved 2008-02-24. [13] "Minerals Yearbook Yearbook Commodity Commodity Summaries Summaries 2009: 2009: Titanium" Titanium" (http://minerals.usgs.gov/minerals/pubs/commodity/titanium/ (http://minerals.usgs.gov/minerals/pubs/commodity/titanium/ myb1-2007-titan.pdf) myb1-2007-titan.pdf) (PDF). US Geological Survey. May 2009. . Retrieved 2008-02-24. 2008-02-24. [14] Gambogi, Joseph Joseph (January 2009). "Titanium "Titanium and Titanium Titanium dioxide Statistics Statistics and Information" (http://minerals. (http://minerals.usgs. usgs.gov/minerals/pubs/ gov/minerals/pubs/ commodity/titanium/mcs-2009-titan.pdf). commodity/titanium/mcs-2009-titan.pdf). US Geological Survey. . Retrieved 2008-02-24. [15] Schemel, Schemel, J. H. H. (1977) (1977).. ASTM Manual on Zirconium and Hafnium (http://books.google. (http://books. google.com/?id=dI_Lssyd com/?id=dI_LssydVeYC). VeYC). ASTM International. pp. 1 – 5. 5. ISBN 9780803105058. 9780803105058. . [16] Hedrick, Hedrick, James B.. "Hafnium "Hafnium"" (http://minerals (http://minerals..er.usgs.gov/minerals/pubs/commodity/zirconium/731798. er.usgs.gov/minerals/pubs/commodity/zirconium/731798.pdf) pdf) (PDF). United States Geological Survey. . Retrieved 2008-09-10. 2008-09-10. [17] Spink, Donald (1961). "Reactive Metals. Zirconium, Zirconium, Hafnium, and Titanium". Industrial and Engineering Chemistry 53 (2): 97 – 104. 104. doi:10.1021/ie50614a019. [18] Hebda, John (2001). "Niobium "Niobium alloys and and high Temperature Temperature Applications" Applications" (http://www.cbmm. (http://www.cbmm.com. com.br/portug/sources/techlib/ br/portug/sources/techlib/ science_techno/table_content/sub_3/images/pdfs/016.pdf) science_techno/table_con tent/sub_3/images/pdfs/016.pdf) (PDF). CBMM. . Retrieved 2008-09-04. [19] Donachie, Donachie, Matthew Matthew J. (2002) (2002).. Superalloys (http://books.google.com/?id=vjCJ5p (http://books.google.com/?id=vjCJ5pI1QpkC&pg=PA23 I1QpkC&pg=PA235). 5). ASTM International. pp. 235 – 236. 236. ISBN 9780871707499. . [20] Emsley, Emsley, John (2001) (2001).. "Titanium" "Titanium".. Nature's Building Blocks: An A-Z Guide to the Elements . Oxford, England, UK: Oxford University Press. pp. 457 – 456. 456. ISBN 0198503407. [21] "Internati "International onal Chemical Chemical Safety Cards" Cards" (http://www.ilo.org/legacy/english/protection/safework/cis/products/icsc/dtasht/_icsc14/ (http://www.ilo.org/legacy/english/protection/safework/cis/products/icsc/dtasht/_icsc14/ icsc1405.htm). icsc1405.htm). International Labour Organization. October 2004. . Retrieved 2008-03-30. [22] "Zirconiu "Zirconium m Compounds Compounds"" (http://www. (http://www.cdc. cdc.gov/niosh/pel88/7440-67 gov/niosh/pel88/7440-67.html). .html). National Institute f or Occupational Health and Safety. 2007-12-17. . Retrieved 2008-02-17. 2008-02-17. [23] "Occupational Safety & Health Administration: Administration: Hafnium" (http://www.osha.gov/SLTC/healthguidelines/hafnium/index. (http://www.osha.gov/SLTC/healthguidelines/hafnium/index.html). html). U.S. Department of Labor. . Retrieved 2008-09-10.
59
Group V
60
Group V Group →
5
↓ Period
4
23 V 5
41 Nb 6
73 Ta 105 Db
7
A Group 5 element is one in the series of elements in group 5 (IUPAC style) in the periodic table, which consists of vanadium (V), niobium (Nb), tantalum (Ta), and dubnium (Db). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells, though niobium curiously does not follow the trend: Z
Elem lement ent No. No. of of ele elect ctro rons ns/s /she hell ll
23
vanadium 2, 8, 8, 11 11, 2
41
niobium
2, 8, 18, 12, 1
73
tantalum
2, 8, 18 18, 32 32, 11, 2
105 105 dubn dubniu ium m
2, 8, 18, 18, 32, 32, 32, 32, 11, 11, 2
Dubnium can only be produced in the laboratory, and does not exist in nature.
Group V
61
History The discovery of all elements in the group led to controversies. The verification of those discoveries was difficult due to similarity of vanadium and group 6 element chromium, the chemical similarity of niobium and tantalum and the complicated setup which was necessary to produce a few atoms of dubnium.
Biological occurrences Of the group 5 elements, only vanadium has been identified as playing a role in the biological chemistry of living systems: it is involved in some of the enzymes of higher organisms, and also, unusually, in the biochemistry of some marine tunicates.
See also Explanation of periodic table slice on right:
Transition metals
atomic number in black are solids
solid borders are older than the Earth (Primordial elements)
dotted borders are made artificially (Synthetic elements)
• Greenwood, N (2003). "Vanadium to dubnium: from confusion through clarity to complexity". Catalysis Today 78: 5. doi:10.1016/S0920-5861(02)00318-8.
Group VI Group →
6
↓ Period
4
24 Cr 5
42 Mo 6
74 W 106 Sg
7
Legend
Transition metal Primordial element Synthetic
Group VI
62
A Group 6 element is one in the series of elements in group 6 (IUPAC style) in the periodic table, which consists of the transition metals chromium (Cr), molybdenum (Mo), tungsten (W), and seaborgium (Sg). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior: Z
Element
No. of electrons/shell
24
chromium
42
molybdenum 2, 8, 18, 13, 1
74
tungsten
106 seaborgium
2, 8, 13, 1
2, 8, 18, 32, 12, 2 2, 8, 18, 32, 32, 12, 2
"Group 6" is the new IUPAC name for this group; the old style name was "group VIA" in the old European system or "group VIB" in the old US system. Group 6 must not be confused with the group with the old-style group names of either VIB (European system) or VIA (US system); that group is now called group 16.
Biological occurrences Group 6 is notable in that it contains some of the only elements in periods 5 and 6 with a known role in the biological chemistry of living organisms: molybdenum is common in enzymes of many organisms, and tungsten has been identified in an analogous role in enzymes from some archaea, such as Pyrococcus furiosus. In contrast, and unusually for a first-row d-block transition metal, chromium appears to have few biological roles, although it is thought to form part of the glucose metabolism enzyme in some mammals.
See also Explanation of right side periodic table slice:
Transition metals
atomic number in black are solids
solid borders are older than the Earth (Primordial elements)
dotted borders are made artificially (Synthetic elements)
Group VII
63
Group VII Group →
7
↓ Period
4
25 Mn
5
43 Tc
6
75 Re
7
107 Bh
A Group 7 element is one in the series of elements in group 7 (IUPAC style) in the periodic table, which consists of the transition metals manganese ( Mn), technetium (Tc), rhenium (Re), and bohrium (Bh). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior: Z
Element
No. of electrons/shell
25
manganese 2, 8, 13, 2
43
technetium 2, 8, 18, 13, 2
75
rhenium
107 bohrium
2, 8, 18, 32, 13, 2 2, 8, 18, 32, 32, 13, 2
All of these elements are classed in Group 7 because their valence shells hold seven electrons. Technetium has no stable isotopes. Technetium and promethium are the only two such elements before lead, after which (with bismuth having an extremely long-lived isotope) no known element has a stable isotope.
Occurrence Two of the four members of the group 2, technetium and bohrium, are radioactive with short enough half life that they are not present in nature. Furthermore rhenium is a rare element which occurs only in traces in other mineral. These facts make manganese the only abundant element of the group. This is also shown in difference in the annual world production. In 2007 11 million metric tons of manganese were mined while in the same year the world production of rhenium was between 40 and 50 metric tons. They are also very reactive.
See also Explanation of right side periodic table slice:
Transition metals
atomic number in black are solids
solid borders are older than the Earth (Primordial elements)
dotted borders are made artificially (Synthetic elements)
dashed borders have no isotopes older than the earth
Group VIII
64
Group VIII For "Group VIII", the rightmost group on the Periodic Table, see noble gas. Group →
4
↓ Period
4
26 Fe 5
44 Ru 6
76 Os 108 Hs
7
Legend
Transition metal Primordial element Synthetic
A Group 8 element is one in the series of elements in group 8 (IUPAC style) in the periodic table, which consists of the transition metals iron (Fe), ruthenium (Ru), osmium (Os) and hassium (Hs). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells though ruthenium curiously does not follow the trend: Z
Element
No. of electrons/shell
26
iron
44
ruthenium 2, 8, 18, 15, 1
76
osmium
108 hassium
2, 8, 14, 2
2, 8, 18, 32, 14, 2 2, 8, 18, 32, 32, 14, 2
All of these elements are classed in Group 8 because their valence shells hold eight electrons. Hassium can only be produced in the laboratory and has not been found in nature.
Group VIII
65
See also • Platinum group Explanation of right side periodic table slice:
Transition metals
atomic number in black are solids
solid borders are older than the Earth (Primordial elements)
dotted borders are made artificially (Synthetic elements)
Group IX Group →
9
↓ Period
4
27 Co
5
45 Rh
6
77 Ir
7
109 Mt
In modern IUPAC nomenclature, Group 9 of the periodic table contains the elements cobalt (Co), rhodium (Rh), iridium (Ir), and meitnerium (Mt). These are all d-block transition metals. All known isotopes of Mt are radioactive with short half-lives, and it is not known to occur in nature; only minute quantities have been synthesized in laboratories. Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior, though rhodium curiously does not follow the pattern: Z
Element
No. of electrons/shell
27
cobalt
2, 8, 15, 2
45
rhodium
2, 8, 18, 16, 1
77
iridium
2, 8, 18, 32, 15, 2
109 meitnerium 2, 8, 18, 32, 32, 15, 2
Applications • • • •
Alloys with other metals, primarily to add corrosion and wear resistance Industrial Catalysts Superalloys Electrical Components
See also • Platinum group Explanation of right side periodic table slice:
Transition metals
atomic number in black are solids
solid borders are older than the Earth (Primordial elements)
dashed borders have no isotopes older than the earth
Group X
66
Group X Group → 10 ↓ Period
4
28 Ni
5
46 Pd
6
78 Pt
7
110 Ds
A Group 10 element is one in the series of elements in group 10 (IUPAC style) in the periodic table, which consists of the transition metals nickel (Ni), palladium (Pd), platinum (Pt), and darmstadtium (Ds). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells (though for this family it is particularly weak with palladium an exceptional case): Z
Element
No. of electrons/shell
28
nickel
2, 8, 16, 2
46
palladium
2, 8, 18, 18
78
platinum
2, 8, 18, 32, 17, 1
110 darmstadtium 2, 8, 18, 32, 32, 17, 1
Properties Group ten metals are white to light grey in color, and possess a high luster, a resistance to tarnish(oxidation) at STP, are highly ductile, and enter into oxidation states of +2 and +4, with +1 being seen in special conditions. The existence of a +3 state is debated, as the state could be an illusory state created by +2 and +4 states. Theory suggests that group 10 metals may produce a +6 oxidation state under precise conditions, but this remains to be proven conclusively in the laboratory.
Applications The group ten metals share several uses. These include: • • • • •
Decorative purposes, in the form of jewelry and electroplating Catalysts in a variety of chemical reactions Metal Alloys Electrical components, due to their predictable changes in electrical resistivity with regard to temperature. Superconductors, as components in alloys with other metals.
Group X
67
See also • Platinum group Explanation of right side periodic table slice:
Transition metals
atomic number in black are solids
solid borders are older than the Earth (Primordial elements)
dashed borders have no isotopes older than the earth
Group XI Group →
11
↓ Period
4
29 Cu 5
47 Ag 6
79 Au 111 Rg
7
Legend
Transition metal Primordial element Synthetic
A Group 11 element is one in the series of elements in group 11 (IUPAC style) in the periodic table, consisting of transition metals which are the traditional coinage metals of copper ( Cu), silver (Ag), and gold (Au). Roentgenium (Rg) belongs to this group of elements based on its electronic configuration, but cannot be considered coinage metal (short lived transactinide with a 3.6 seconds half-life). The name "coinage metals" should not be used as an alternative name for Group 11 elements, as various cultures have used other metals in coinage including aluminum, lead, nickel, stainless steel, and zinc). The term 'coinage metal' is not recognized by the International Union of Pure and Applied Chemistry (IUPAC) as a designator for group 11 elements.
Group XI
68
History All the elements of the group except Roentgenium have been known since prehistoric times, as all of them occur in metallic form in nature and no extraction metallurgy has to be used to produce them.
Characteristics Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior: Z
Element
No. of electrons/shell
29
copper
2, 8, 18, 1
47
silver
2, 8, 18, 18, 1
79
gold
2, 8, 18, 32, 18, 1
111 roentgenium 2, 8, 18, 32, 32, 18, 1
They are all relatively inert, corrosion-resistant metals. Copper and gold are colored. These elements have low electrical resistivity so they are used for wiring. Copper is the cheapest and most widely used. Bond wires for integrated circuits are usually gold. Silver and silver plated copper wiring are found in some special applications.
Applications These metals, especially silver, have unusual properties that make them essential for industrial applications outside of their monetary or decorative value. They are all excellent conductors of electricity. The most conductive of all metals are silver, copper and gold in that order. Silver is also the most thermally conductive element, and the most light reflecting element. Silver also has the unusual property that the tarnish that forms on silver is still highly electrically conductive. Copper is used extensively in electrical wiring and circuitry. Gold contacts are sometimes found in precision equipment for their ability to remain corrosion-free. Silver is used widely in mission-critical applications as electrical contacts, and is also used in photography (because silver nitrate reverts to metal on exposure to light), agriculture, medicine, audiophile and scientific applications. Gold, silver, and copper are quite soft metals and so are easily damaged in daily use as coins. Precious metal may also be easily abraded and worn away through use. In their numismatic functions these metals must be alloyed with other metals to afford coins greater durability. The alloying with other metals makes the resulting coins harder, less likely to become deformed and more resistant to wear. Gold coins: Gold coins are typically produced as either 90% gold (e.g. with pre-1933 US coins), or 22 carat (92%)
gold (e.g. current collectible coins and Krugerrands), with copper and silver making up the remaining weight in each case. Bullion gold coins are being produced with up to 99.999% gold (in the Canadian Gold Maple Leaf series). Silver coins: Silver coins are typically produced as either 90% silver - in the case of pre 1965 US minted coins
(which were circulated in many countries), or sterling silver (92.5%) coins for pre-1967 British Commonwealth and other silver coinage, with copper making up the remaining weight in each case. Copper coins: Copper coins are often of quite high purity, around 97%, and are usually alloyed with small amounts
of zinc and tin. Inflation has caused the face value of coins to fall below the hard currency value of the historically used metals. This had led to most modern coins being made of base metals - copper nickel (around 80:20, silver in color) is popular as are nickel-brass (copper (75), nickel (5) and zinc (20), gold in color), manganese-brass (copper, zinc, manganese,
Group XI
69
and nickel), bronze, or simple plated steel.
See also Explanation of right side periodic table slice:
Transition metals
atomic numbers in black are solids
solid borders are older than the Earth (Primordial elements)
dashed borders have no isotopes older than the earth
Group XII Group 11
12 13
Period 29
30 31 Cu Zn Ga
4 5
47 48 49 Ag Cd In
6
79 80 81 Au Hg Tl
7
112 Cn
A group 12 element is one of the elements in group 12 (IUPAC style) in the periodic table, consisting of zinc ( Zn), [1] [2] [3] cadmium (Cd) and mercury (Hg). The inclusion of copernicium (Cn) in group 12 is supported by recent [4] experiments on individual Cn atoms.
Some properties of the elements zinc
cadmium
mercury
Electronic configuration
[Ar]3d 4s
[Kr]4d 5s
[Xe]4f 5d 6s
Metallic radius /pm
134
151
151
Ionic radius /pm (M2+)
74
95
102
Electronegativity
1.6
1.7
1.9
Melting point /°C
419.5
320.8
−38.9
Boiling point /°C
907
765
357
10
2
10
2
14
10
2
All elements in this group are metals. The similarity of the metallic radii of cadmium and mercury is a knock-on effect of the lanthanide contraction. So, the trend in this group is unlike the trend in group 2, the alkaline earths, where metallic radius increases smoothly from top to bottom of the group. All three metals have relatively low melting and boiling points, which indicates that the metallic bond is relatively weak, with relatively little overlap between the valence band and the conduction band. Thus, zinc is close to the boundary between metallic and metalloid elements which is usually placed between gallium and germanium, though gallium participates in semi-conductors such as gallium arsenide. Zinc is the most electropositive element in the group and zinc metal is a good reducing agent. The group oxidation state is +2 in which the ions have the rather stable d10 electronic configuration, with a full sub-shell. However, mercury can easily be reduced to the +1 oxidation state; usually, as in the ion Hg22+, two mercury(I) ions come − together to form a metal-metal bond and a diamagnetic species. Cadmium can also form species such as [Cd 2Cl6]4
Group XII
70
in which the metal's oxidation state is +1. Just as with mercury, the formation of a metal-metal bond results in a diamagnetic compound in which there are no unpaired electrons which would otherwise make the species very reactive. Zinc(I) is known only in the gas phase, in such compounds as linear Zn2Cl2, analogous to calomel. −
All three metal ions form many tetrahedral species, such as MCl 42 . When a divalent ion of these elements forms a tetrahedral complex, it obeys the octet rule. Both zinc and cadmium can also form octahedral complexes such as the aqua ions [M(H O) ]2+ which are present in aqueous solutions of salts of these metals. Covalent character is 2 6 achieved by using the 4d or 5d orbitals, respectively, forming sp3d2 hybrid orbitals. Mercury, however, rarely exceeds a coordination number of four; when it does so the 5f orbitals must be involved. Coordination numbers of 2, 3, 5, 7 and 8 are also known. The elements in group 12 are usually considered to be d-block elements, but not transition elements as the d-shell is full. Some authors classify these elements as main-group elements because the valence electrons are in ns2 orbitals. Nevertheless zinc shares many characteristics with the neighbouring transition metal, copper. For instance, zinc complexes merit inclusion in the Irving-Williams series as zinc forms many complexes with the same stoichiometry as complexes of copper(II), albeit with smaller stability constants. There is little similarity between cadmium and silver as compounds of silver(II) are rare and those that do exist are very strong oxidizing agents. Likewise the common oxidation state for gold is +3, which precludes there being much common chemistry between mercury and gold, though there are similarities between mercury(I) and gold(I) such as the formation of linear dicyano − complexes, [M(CN)2] .
References [1] Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419 [2] Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999), Advanced Inorganic Chemistry (6th ed.), New York: Wiley-Interscience, ISBN 0-471-19957-5 [3] Housecroft, C. E.; Sharpe, A. G. (2008). Inorganic Chemistry (3rd ed.). Prentice Hall. ISBN 978-0131755536. [4] Eichler, R; Aksenov, NV; Belozerov, AV; Bozhikov, GA; Chepigin, VI; Dmitriev, SN; Dressler, R; Gäggeler, HWet al. (2007). "Chemical Characterization of Element 112". Nature 447 (7140): 72 – 75. doi:10.1038/nature05761. PMID 17476264.
Group XIII
71
Group XIII Group → 13 ↓ Period
2
5 B
3
13 Al
4
31 Ga
5
49 In
6
81 Tl
The boron group is the series of elements in group 13 (IUPAC style) in the periodic table. The boron group consists of boron (B), aluminium (Al), gallium (Ga), indium (In), thallium (Tl), and ununtrium (Uut). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior: Z
5
Element
boron
No. of electrons/shell
2, 3
13 aluminium 2, 8, 3 31 gallium
2, 8, 18, 3
49 indium
2, 8, 18, 18, 3
81 thallium
2, 8, 18, 32, 18, 3
The group has previously also been referred to as the earth metals and the triels, from the Latin tri, three, stemming from the naming convention of this group as Group IIIB. These elements are characterized by having three electrons in their outer energy levels (valence layers). Boron is considered a metalloid, and the rest are considered metals of the poor metals groups. Boron occurs sparsely probably because of disruption of its nucleus by bombardment with subatomic particles produced from natural radioactivity. Aluminium occurs widely on earth and in fact, it is the third most abundant element in the Earth's crust (7.4%).
Explanation of above periodic table slice: Metalloids
Poor metals
atomic number in black are solid borders are primordial elements (older than dotted borders are radioactive, synthetic solids the Earth) elements
Group XIV
72
Group XIV Group →
14
↓ Period
2
6 C
3
14 Si
4
32 Ge
5
50 Sn
6
82 Pb
7
114 Uuq
The carbon group is a periodic table group consisting of carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and ununquadium (Uuq). In modern IUPAC notation, it is called Group 14. In the old IUPAC and CAS systems, it was called Group IVB and Group IVA, respectively.[1] In the field of semiconductor physics, it is still universally called Group IV. The group was once also known as the tetrels (from Greek tetra, four), stemming from the Roman numeral IV in the group names, or (not coincidentally) from the fact that these elements have four valence electrons (see below). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior: Z
Element
No. of electrons/shell
6
carbon
2, 4
14
silicon
2, 8, 4
32
germanium
2, 8, 18, 4
50
tin
2, 8, 18, 18, 4
82
lead
2, 8, 18, 32, 18, 4
114 ununquadium 2, 8, 18, 32, 32, 18, 4
Each of the elements in this group has 4 electrons in its outer energy level. The last orbital of all these elements is the p2 orbital. In most cases, the elements share their electrons. The tendency to lose electrons increases as the size of the atom increases, as it does with increasing atomic number. Carbon alone forms negative ions, in the form of carbide − (C4 ) ions. Silicon and germanium, both metalloids, each can form +4 ions. Tin and lead both are metals while ununquadium is a synthetic short-lived radioactive metal. Tin and lead are both capable of forming +2 ions. Except for germanium and ununquadium, all of these elements are familiar in daily life either as the pure element or in the form of compounds. However, except for silicon, none of these elements are particularly plentiful in the Earth’s crust. Carbon forms a very large variety of compounds, in both the plant and animal kingdoms. Silicon and silicate minerals are fundamental components of the Earth’s crust; silica (silicon dioxide) is sand. Tin and lead, although with very low abundances in the crust, are nevertheless common in everyday life. They occur in highly concentrated mineral deposits, can be obtained easily in the metallic state from those minerals, and are
Group XIV
73
useful as metals and as alloys in many applications. Germanium, on the other hand, forms few characteristic minerals and is most commonly found only in small concentrations in association with the mineral zinc blende and in coals. Although germanium is indeed one of the rarer elements, it assumed importance upon recognition of its properties as a semiconductor.
History Carbon, tin, and lead, are a few of the elements well known in the ancient world - together with sulfur, iron, copper, mercury, silver, and gold. Carbon as an element was discovered by the first human to handle charcoal from his fire. Modern carbon chemistry dates from the development of coals, petroleum, and natural gas as fuels and from the elucidation of synthetic organic chemistry, both substantially developed since the 1800s. Amorphous elemental silicon was first obtained pure in 1824 by the Swedish chemist Jöns Jacob Berzelius; impure silicon had already been obtained in 1811. Crystalline elemental silicon was not prepared until 1854, when it was obtained as a product of electrolysis. In the form of rock crystal, however, silicon was familiar to the predynastic Egyptians, who used it for beads and small vases; to the early Chinese; and probably to many others of the ancients. The manufacture of glass containing silica was carried out both by the Egyptians — at least as early as 1500 BCE — and by the Phoenicians. Certainly, many of the naturally occurring compounds called silicates were used in various kinds of mortar for construction of dwellings by the earliest people. Germanium is one of three elements the existence of which was predicted in 1871 by the Russian chemist Dmitri Mendeleev when he first devised his periodic table. Not until 1886, however, was germanium identified as one of the elements in a newly found mineral. The origins of tin seem to be lost in history. It appears that bronzes, which are alloys of copper and tin, were used by prehistoric man some time before the pure metal was isolated. Bronzes were common in early Mesopotamia, the Indus Valley, Egypt, Crete, Israel, and Peru. Much of the tin used by the early Mediterranean peoples apparently came from the Scilly Isles and Cornwall in the British Isles, [2] where mining of the metal dates from about 300 – 200 BCE. Tin mines were operating in both the Inca and Aztec areas of South and Central America before the Spanish conquest. Lead is mentioned often in early Biblical accounts. The Babylonians used the metal as plates on which to record inscriptions. The Romans used it for tablets, water pipes, coins, and even cooking utensils; indeed, as a result of the last use, lead poisoning was recognized in the time of Augustus Caesar. The compound known as white lead was apparently prepared as a decorative pigment at least as early as 200 BCE. Modern developments date to the exploitation in the late 1700s of deposits in the Missouri – Kansas – Oklahoma area in the United States.
Explanation of above periodic table slice: Nonmetals Metalloids
Poor metals
atomic number in black are solids
solid borders are primordial elements (older than the Earth)
dotted borders are radioactive, synthetic elements
References [1] Fluck, E. New notations in the periodic table. Pure & App. Chem. 1988, 60, 431-436. (http://www.iupac.org/publications/pac/1988/pdf/ 6003x0431. pdf) [2] Online Encyclopaedia Britannica, Tin (http://www.britannica.com/EBchecked/topic/596431/tin)
Group XV
74
Group XV Group →
15
↓ Period
2
7 N
3
15 P
4
33 As
5
51 Sb
6
83 Bi
7
115 Uup
The nitrogen group is a periodic table group consisting of nitrogen ( N), phosphorus (P), arsenic (As), antimony (Sb), bismuth (Bi) and ununpentium (Uup) (unconfirmed). In modern IUPAC notation, it is called Group 15. In the old IUPAC and CAS systems, it was called Group VB and [1] Group VA, respectively (pronounced "group five B" and "group five A", because "V" is a Roman numeral). In the field of semiconductor physics, it is still universally called Group V.[2] It is also collectively named the [3] pnictogens. The "five" ("V") in the historical names comes from the fact that these elements have five valence electrons (see below). Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior: Z
Element
No. of electrons/shell
7
nitrogen
2, 5
15
phosphorus
2, 8, 5
33
arsenic
2, 8, 18, 5
51
antimony
2, 8, 18, 18, 5
83
bismuth
2, 8, 18, 32, 18, 5
115 ununpentium 2, 8, 18, 32, 32, 18, 5
This group has the defining characteristic that all the component elements have 5 electrons in their outermost shell, that is 2 electrons in the s subshell and 3 unpaired electrons in the p subshell. They are therefore 3 electrons short of filling their outermost electron shell in their non-ionized state. The most important element of this group is nitrogen (chemical symbol N), which in its diatomic form is the principal component of air. Binary compounds of the group can be referred to collectively as pnictides. The spelling derives from the Greek πνίγειν ( pnigein), to choke or stifle, which is a property of nitrogen; they are also mnemonic for the two most common members, P and N. The name pentels (from the Latin penta, five) was also used for this group at one time, stemming from the earlier group naming convention (Group VB).
Group XV
75
These elements are also noted for their stability in compounds due to their tendency for forming double and triple covalent bonds. This is the property of these elements which leads to their potential toxicity, most evident in phosphorus, arsenic and antimony. When these substances react with various chemicals of the body, they create strong free radicals not easily processed by the liver, where they accumulate. Paradoxically it is this strong bonding which causes nitrogen and bismuth's reduced toxicity (when in molecules), as these form strong bonds with other atoms which are difficult to split, creating very unreactive molecules. For example N , the diatomic form of nitrogen, 2 is used for inert atmosphere in situations where argon or another noble gas would be prohibitively expensive.
A collection of nitrogen-group chemical element samples.
The nitrogen group consists of two non-metals, two metalloids, one metal, and one synthetic (presumably metallic) element. All the elements in the group are a solid at room temperature except for Nitrogen which is a gas at room temperature. Nitrogen and bismuth, despite both being part of the nitrogen group, are very different in their physical properties. For example, at STP nitrogen is a transparent nonmetallic gas, while bismuth is a brittle pinkish metallic solid.
See also • oxypnictide includes superconductors discovered in 2008 • ferropnictide includes oxypnictide superconductors.
Notes [1] Fluck, E. New notations in the periodic table. Pure & App. Chem. 1988, 60, 431-436. (http://www.iupac.org/publications/pac/1988/pdf/ 6003x0431. pdf) [2] For example, a 2005 book (http://books.google.com/books?id=J6W5n5Z1EQIC) is titled Properties of group-IV, III-V and II-VI semiconductors. [3] Edited by N G Connelly and T Damhus (with R M Hartshorn and A T Hutton), ed (2005). Nomenclature of Inorganic Chemistry: IUPAC Recommendations 2005 section IR-3.5 (http://www.iupac.org/publications/books/rbook/Red_Book_2005.pdf). ISBN 0-85404-438-8. .
Explanation of above periodic table slice: Nonmetals Met alloids
Poor metals
atomic number in red are gases
atomic number in black are solids
solid borders are primordial elements (older than the Earth)
dotted borders are radioactive, synthetic elements
Group XVI
76
Group XVI Group →
16
↓ Period
2
8 O
3
16 S
4
34 Se
5
52 Te
6
84 Po
7
116 Uuh
The chalcogens (pronounced / ˈkælk ədʒɨ n/) are the chemical elements in group 16 (old-style: VIB or VIA) of the periodic table. This group is also known as the oxygen family. It consists of the elements oxygen (O), sulfur (S), selenium (Se), tellurium (Te), the radioactive element polonium (Po), and the synthetic element ununhexium (Uuh). Although all group 16 elements of the periodic table, including oxygen are defined as chalcogens, oxygen and oxides are usually distinguished from chalcogens and chalcogenides. The term chalcogenide is more commonly reserved for sulfides, selenides, and tellurides, rather than for oxides. Oxides are usually not indicated as chalcogenides.[1] [2] [3] [4] Binary compounds of the chalcogens are called chalcogenides (rather than chalcides, which breaks the pattern of halogen / halide and pnictogen / pnictide).
Properties Members of this group show similar patterns in their electron configuration, especially the outermost shells, resulting in similar trends in chemical behavior: Z
Element
No. of electrons/shell
8
oxygen
2, 6
16
sulfur
2, 8, 6
34
selenium
2, 8, 18, 6
52
tellurium
2, 8, 18, 18, 6
84
polonium
2, 8, 18, 32, 18, 6
116 ununhexium 2, 8, 18, 32, 32, 18, 6
Oxygen and sulfur are nonmetals, and selenium, tellurium, and polonium are metalloid semiconductors (that means, their electrical properties are between those of a metal and an insulator). Nevertheless, tellurium, as well as selenium, is often referred to as a metal when in elemental form. Metal chalcogenides are common as minerals. For example, pyrite (FeS2) is an iron ore. The rare mineral calaverite is the ditelluride AuTe2.
Group XVI
77
The formal oxidation number of the most common chalcogen copounds is −2. Other values, such as −1 in pyrite, can be attained. The highest formal oxidation number +6 is found in sulfates, selenates and tellurates, such as in sulfuric acid or sodium selenate (Na2SeO4).
Explanation of above periodic table slice: Nonmetals Met alloids
Poor metals
Atomic Atomic numbers in red numbers in are gases black are solids
Solid borders indicate primordial elements (older than the Earth)
Dashed borders indicate radioactive natural elements
Dotted borders indicate radioactive synthetic elements
Etymology The name chalcogen comes from the Greek words χαλκος (chalkos, literally "copper"), and γενεσ (genes, born)[5] . Thus the chalcogens give birth to, produce copper. It was first used around 1930 by Wilhelm Biltz's group at the University of Hanover, where it was proposed by a man named Werner Fischer.[6] Although the literal meanings of the Greek words imply that chalcogen means "copper-former", this is misleading because the chalcogens have nothing to do with copper in particular. "Ore-former" has been suggested as a better translation,[7] both because the vast majority of metal ores are chalcogenides, and because the word χαλκος in ancient Greek was associated with metals and metal-bearing rock in general (because copper, and its alloy bronze, was one of the first metals to be used by humans).
See also • Gold chalcogenides
References [1] A Second Note on the Term "Chalcogen" (http://jchemed.chem.wisc.edu/Journal/Issues/2001/Oct/abs1333_1.html) [2] Francesco Devillanova (Editor) Handbook of Chalcogen Chemistry - New Perspectives in Sulfur, Selenium and Tellurium (http://books. google.com/books?id=IvGnUAaSqOsC&pg=PT24) Royal Society of Chemistry, 2007; ISBN 0854043667, 9780854043668 [3] IUPAC goldbook amides (http://goldbook.iupac.org/A00266.html). Chalcogen replacement analogues (of amides) are called thio-, selenoand telluro-amides. [4] Ohno Takahisa Passivation of GaAs(001) surfaces by chalcogen atoms (S, Se and Te) (http://www.sciencedirect.com/ science?_ob=ArticleURL&_udi=B6TVX-46T3HTC-JF&_user=10&_rdoc=1&_fmt=&_orig=search&_sort=d&_docanchor=&view=c& _acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=aae920d4ee8b01faf1150483b04710a8) Surface Science; Volume 255, Issue 3, 2 September 1991, Pages 229-236; doi:10.1016/0039-6028(91)90679-M [5] Online Etymology Dictionary -gen (http://www.etymonline.com/index.php?term=-gen) [6] Werner Fischer (2001). "A Second Note on the Term "Chalcogen"". Journal of Chemical Education 78 (10): 1333. doi:10.1021/ed078p1333.1. [7] William B. Jensen (1997). Journal of Chemical Education 74 (9): 1063. doi:10.1021/ed074p1063.
Group XVII
78
Group XVII Group →
17
↓ Period
2
9 F
3
17 Cl
4
35 Br
5
53 I
6
85 At
7
117 Uus
Legend
Halogen Gas Liquid Primordial element From decay Synthetic
The halogens or halogen elements are a series of nonmetal elements from Group 17 IUPAC Style (formerly: VII, VIIA) of the periodic table, comprising fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117, provisionally referred to by the systematic name ununseptium, may also be a halogen. The group of halogens is the only periodic table group which contains elements in all three familiar states of matter at standard temperature and pressure.
Abundance Owing to their high reactivity, the halogens are found in the environment only in compounds or as ions. Halide ions − and oxoanions such as iodate (IO ) can be found in many minerals and in seawater. Halogenated organic 3 compounds can also be found as natural products in living organisms. In their elemental forms, the halogens exist as diatomic molecules, but these only have a fleeting existence in nature and are much more common in the laboratory and in industry. At room temperature and pressure, fluorine and chlorine are gases, bromine is a liquid and iodine and astatine are solids; Group 17 is therefore the only periodic table group exhibiting all three states of matter at room temperature.
Group XVII
79
Etymology The Swedish chemist Baron Jöns Jakob Berzelius coined the term "halogen" – ἅλς (háls), "salt" or "sea", and γεν(gen-), from γίγνομαι (gí gnomai), "come to be" – for an element that produces a salt when it forms a compound with a metal.[1]
Properties Like other groups, the candidates of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:
Fluorine, (F); chlorine, (Cl); bromine, (Br); iodine, (I) at room temperature. The first two are gaseous, the third is liquid and the fourth is solid.
Z
Element
No. of electrons/shell
9
fluorine
2, 7
17
chlorine
2, 8, 7
35
bromine
2, 8, 18, 7
53
iodine
2, 8, 18, 18, 7
85
astatine
2, 8, 18, 32, 18, 7
117 ununseptium 2, 8, 18, 32, 32, 18, 7
The halogens show a series of trends when moving down the group —for instance, decreasing electronegativity and reactivity, and increasing melting and boiling point.
Group XVII
80
Halogen Standard Atomic Weight (u) Melting Point (K) Boiling Point (K) Electronegativity (Pauling)
Fluorine
18.998
53.53
85.03
3.98
Chlorine
35.453
171.60
239.11
3.16
Bromine
79.904
265.80
332.00
2.96
Iodine
126.904
386.85
457.40
2.66
(210)
575
610 (?)
2.20
Astatine
Diatomic halogen molecules halogen molecule
structure
model
d (X−X) /
d (X−X) / pm
pm
(solid phase)
(gas phase)
fluorine
F2
143
149
chlorine
Cl2
199
198
bromine
Br2
228
227
iodine
I2
266
272
The elements become less reactive and have higher melting points as the atomic number increases.
Chemistry Reactivity Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in sufficient quantities. This high reactivity is due to the atoms being highly electronegative due to their high effective nuclear charge. They can gain an electron by reacting with atoms of other elements. Fluorine is one of the most reactive elements in existence, attacking otherwise inert materials such as glass, and forming compounds with the heavier noble gases. It is a corrosive and highly toxic gas. The reactivity of fluorine is such that if used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4). Thus fluorine must be handled with substances such as Teflon (which is itself an organofluorine compound), extremely dry glass, or metals such as copper or steel which form a protective layer of fluoride on their surface. The high reactivity of fluorine means that once it does react with something, it bonds with it so strongly that the resulting molecule is very inert and non-reactive to anything else. For example, Teflon is fluorine bonded with carbon. Both chlorine and bromine are used as disinfectants for drinking water, swimming pools, fresh wounds, spas, dishes, and surfaces. They kill bacteria and other potentially harmful microorganisms through a process known as sterilization. Their reactivity is also put to use in bleaching. Sodium hypochlorite, which is produced from chlorine, is the active ingredient of most fabric bleaches and chlorine-derived bleaches are used in the production of some paper products.
Group XVII
Hydrogen halides The halogens all form binary compounds with hydrogen known as the hydrogen halides (HF, HCl, HBr, HI, and HAt), a series of particularly strong acids. When in aqueous solution, the hydrogen halides are known as hydrohalic acids. HAt, or "hydroastatic acid", should also qualify, but it is not typically included in discussions of hydrohalic acid due to astatine's extreme instability toward alpha decay.
Interhalogen compounds The halogens react with each other to form interhalogen compounds. Diatomic interhalogen compounds such as BrF, ICl, and ClF bear resemblance to the pure halogens in some respects. The properties and behaviour of a diatomic interhalogen compound tend to be intermediate between those of its parent halogens. Some properties, however, are found in neither parent halogen. For example, Cl2 and I2 are soluble in CCl4, but ICl is not since it is a polar molecule due to the relatively large electronegativity difference between I and Cl.
Organohalogen compounds Many synthetic organic compounds such as plastic polymers, and a few natural ones, contain halogen atoms; these are known as halogenated compounds or organic halides. Chlorine is by far the most abundant of the halogens, and the only one needed in relatively large amounts (as chloride ions) by humans. For example, chloride ions play a key role in brain function by mediating the action of the inhibitory transmitter GABA and are also used by the body to produce stomach acid. Iodine is needed in trace amounts for the production of thyroid hormones such as thyroxine. On the other hand, neither fluorine nor bromine are believed to be essential for humans.
Polyhalogenated compounds Polyhalogenated compounds are industrially created compounds substituted with multiple halogens. Many of them are very toxic and bioaccumulate in humans, and have a very wide application range. They include the much maligned PCB's, PBDE's, and PFC's as well as numerous other compounds.
Drug discovery In drug discovery, the incorporation of halogen atoms into a lead drug candidate results in analogues that are usually more lipophilic and less water soluble.[2] Consequently, halogen atoms are used to improve penetration through lipid membranes and tissues. Consequently, there is a tendency for some halogenated drugs to accumulate in adipose tissue. The chemical reactivity of halogen atoms depends on both their point of attachment to the lead and the nature of the halogen. Aromatic halogen groups are far less reactive than aliphatic halogen groups, which can exhibit considerable chemical reactivity. For aliphatic carbon-halogen bonds the C-F bond is the strongest and usually less chemically reactive than aliphatic C-H bonds. The other aliphatic-halogen bonds are weaker, their reactivity increasing down the periodic table. They are usually more chemically reactive than aliphatic C-H bonds. Consequently, the most common halogen substitutions are the less reactive aromatic fluorine and chlorine groups.
81
Group XVII
82
Reactivity with water Fluorine reacts vigorously with water to produce oxygen (O ) and hydrogen fluoride (HF):[3] 2
2 F2(g) + 2 H2O(l) → O2(g) + 4 HF(aq)
Chlorine has minimal solubility of 0.7g Cl2 per kg of water at ambient temperature (21 oC).[4] Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid, a solution that can be used as a disinfectant or bleach: Cl2(g) + H2O(l) → HCl(aq) + HClO(aq)
Bromine has a solubility of 3.41 g per 100 g of water,[5] but it slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO): Br (g) + H O(l) → HBr(aq) + HBrO(aq) 2
2
[6]
Iodine, however, is minimally soluble in water (0.03 g/100 g water @ 20 °C) and does not react with it. However, iodine will form an aqueous solution in the presence of iodide ion, such as by addition of potassium iodide (KI), because the triiodide ion is formed.
See also • Pseudohalogen • Halogen bond • Halogen lamp
Further reading • N. N. Greenwood, A. Earnshaw, Chemistry of the Elements [7], 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
Explanation of above periodic table slice: Halogens
Atomic numbers in red are gases
Atomic numbers in green are liquids
Atomic numbers in black are solids
Solid borders indicate primordial elements (older than the Earth)
Dashed borders indicate radioactive natural elements
Dotted borders indicate radioactive synthetic elements
No borders indicates undiscovered elements
References [1] [2] [3] [4] [5] [6] [7]
Online Etymology Dictionary halogen (http://www.etymonline.com/index.php?search=halogen). G. Thomas, Medicinal Chemistry an Introduction , John Wiley & Sons, West Sussex, UK, 2000. The Oxidising Ability of the Group 7 Elements (http://www.chemguide.co.uk/inorganic/group7/halogensasoas.html) Solubility of chlorine in water (http://www.resistoflex.com/chlorine_graphs.htm#9) Properties of bromine (http://www.bromaid.org/hand_chap1.htm) Iodine MSDS (http://www. jtbaker.com/msds/englishhtml/I2680.htm) http://www.knovel.com/knovel2/Toc. jsp?BookID=402
Group XVIII
83
Group XVIII Group →
18
↓ Period
1
2 He
2
10 Ne
3
18 Ar
4
36 Kr
5
54 Xe
6
86 Rn
7
118 Uuo
Legend
Noble gas Gas Primordial element From decay Synthetic
A group 18 element is any chemical element from the last column of the standard periodic table. For the first six periods, the group 18 elements are exactly the noble gases. However, the seventh member of group 18 (the synthetic element ununoctium) is probably not a noble gas. Group 18 was previously called 'group 8A' or 'group 0'.
Properties According to the classical shell model for electrons, the group 18 elements have a fully filled outer shell, rendering them inert to most chemical reactions. This holds true for the first six elements of this group (though they tend to become slightly less inert with increasing periods). For the seventh period group 18 element (ununoctium), this "nobility" is predicted to break down due to relativistic effects.[1]
See also Noble gas
Group XVIII
References [1] Clinton S. Nash (2005). "Atomic and Molecular Properties of Elements 112, 114, and 118". J. Phys. Chem. A 109 (15): 3493 – 3500. doi:10.1021/jp050736o. PMID 16833687.
84
85
Periods Period In the periodic table of the elements, elements are arranged in a series of rows (or periods) so that those with similar properties appear in vertical columns. Elements of the same period have the same number of electron shells; with each group across a period, the elements have one more proton and electron and become less metallic. This arrangement reflects the periodic recurrence of similar properties as the atomic number increases. For example, the alkaline metals lie in one group (group 1) and share similar properties, such as high reactivity and the tendency to lose one electron to arrive at a noble-gas electronic configuration. Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. As atomic number increases, shells fill with electrons in approximately the order shown below. The filling of each shell corresponds to a row in the table. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 8s In the s-block and p-block of the periodic table, elements within the same period generally do not exhibit trends and similarities in properties (vertical trends down groups are more significant). However in the d-block, trends across periods become significant, and in the f-block elements show a high degree of similarity across periods (particularly the lanthanides).
Periods Seven periods of elements occur naturally on Earth. For period 8, which includes elements which may be synthesized after 2010, see the extended periodic table. A group in chemistry means a family of objects with similarities like different families.
Chemical elements in the first period Group 1/17 2/18 # Name
1 H
2 He
The first period contains fewer elements than any other, with only two, hydrogen and helium. They therefore do not follow the octet rule. Chemically, helium behaves as a noble gas, and thus is taken to be part of the group 18 elements. However, in terms of its nuclear structure it belongs to the s block, and is therefore sometimes classified as a group 2 element, or simultaneously both 2 and 18. Hydrogen readily loses and gains an electron, and so behaves chemically as both a group 1 and a group 17 element.
Period
86
• Hydrogen (H) is the most abundant of the chemical elements, constituting roughly 75% of the universe's elemental mass.[1] Ionized hydrogen is just a proton. Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from hydrocarbons such as methane. Hydrogen can form compounds with most elements and is present in water and most organic compounds.[2] [3] • Helium (He) exists only as a gas except in extreme conditions. It is the second lightest element and is the second most abundant in the universe.[4] Most helium was formed during the Big Bang, but new helium is created through nuclear fusion of hydrogen in stars.[5] On Earth, helium is relatively rare, only occurring as a byproduct of the natural decay of some radioactive elements.[6] Such 'radiogenic' helium is trapped within natural gas in concentrations of up to seven percent by volume.[7]
Chemical elements in the second period Group 1
2
13 14 15 16 17 18
3
4
5
#
Name Li Be B
6 C
7 8 N O
9 10 F Ne
Period 2 elements involve the 2s and 2p orbitals. They include the biologically most essential elements besides hydrogen: carbon, nitrogen, and oxygen. • Lithium is the lightest metal and the least dense solid element.[8] In its non-ionized state it is one of the most reactive elements, and so is only ever found naturally in compounds. It is the heaviest primordial element forged in large quantities during the Big Bang. • Beryllium has one of the highest melting points of all the light metals. Small amounts of beryllium were synthesised during the Big Bang, although most of it decayed or reacted further within stars to create larger nucleii, like carbon, nitrogen or oxygen. Beryllium is classified by the International Agency for Research on Cancer as Group 1 carcinogens. [9] Between 1% and 15% of people are sensitive to beryllium and may develop an inflammatory reaction in their respiratory system and skin, called chronic beryllium disease.[10] • Boron (B) does not occur naturally as a free element, but in compounds such as borates. It is an essential plant micronutrient, required for cell wall strength and development, cell division, seed and fruit development, sugar transport and hormone development,[11] [12] though high levels are toxic. • Carbon (C) is the fourth most abundant element in the universe by mass after hydrogen, helium and oxygen[13] and is the second most abundant element in the human body by mass after oxygen,[14] the third most abundant by number of atoms.[15] There are an almost infinite number of compounds that contain carbon due to carbon's ability to form long stable chains of C —C bonds.[16] [17] All organic compounds, those essential for life, contain at least one atom of carbon;[16] [17] combined with hydrogen, oxygen, nitrogen, sulfur, and phosphorus, carbon is the basis of every important biological compound.[17] • Nitrogen (N) is found mainly as mostly inert diatomic gas, N , which makes up 78% of the earth's atmosphere. It 2 is an essential component of proteins and therefore of life. • Oxygen (O) comprising 21% of the atmosphere and is required for respiration by all (or nearly all) animals, as well as being the principal component of water. Oxygen is the third most abundant element in the universe, and oxygen compounds dominate the earth's crust. • Fluorine (F) is the most reactive element in its non-ionized state, and so is never found that way in nature. • Neon is a noble gas used in neon lights.
Period
87
Chemical elements in the third period Group
1
2
13 14 15 16 17 18
11 12 13 14 15 16 17 18 Name Na Mg Al Si P S Cl Ar #
All period three elements occur in nature and have at least one stable isotope. All but the noble gas argon are all essential to basic geology and biology. • Sodium (symbol Na) is an alkali metal. It is present in Earth's oceans in large quantities in the form of sodium chloride (table salt). • Magnesium (symbol Mg) is an alkaline earth metal. Magnesium ions are found in chlorophyll. • Aluminium (symbol Al) is a poor metal. It is the most abundant metal in the Earth's crust. • Silicon (symbol Si) is a metalloid. It is a semiconductor, making it the principal component in many integrated circuits. Silicon dioxide is the principal constituent of sand. • Phosphorus (symbol P) is a nonmetal essential to DNA. It is highly reactive, and as such is never found in nature as a free element. • Sulfur (symbol S) is a nonmetal. It is found in two amino acids: cysteine and methionine. • Chlorine (symbol Cl) is a halogen. It is used as a disinfectant, especially in swimming pools. • Argon (symbol Ar) is a noble gas, making it almost entirely nonreactive. Incandescent lamps are often filled with noble gasses such as argon in order to preserve the filaments at high temperatures.
Chemical elements in the fourth period Group Atomic number Name electron configuration all begin with [Ar]
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19 20 21 22 23 24 25 26 27 28 K Ca Sc Ti V Cr Mn Fe Co Ni
29 Cu
30 Zn
31 Ga
32 Ge
33 As
34 Se
35 Br
36 Kr
4s1 4s2 3d1 3d2 3d3 3d5 3d5 3d6 3d7 3d8 3d10 3d10 3d10 3d10 3d10 3d10 3d10 3d10 2 2 2 1 2 2 2 2 1 2 2 2 2 2 2 2 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4s 4p1 4p2 4p3 4p4 4p5 4p6
Period 4 includes the biologically essential elements potassium and calcium, and is the first period in the d-block with the lighter transition metals. These include iron, the heaviest element forged in main-sequence stars and a principal component of the earth, as well as other important metals such as cobalt, nickle, copper, and zinc. Almost all have biological roles.
From left to right, aqueous solutions of: Co(NO3)2 (red); K2Cr2O7 (orange); K2CrO4 (yellow); NiCl2 (green); CuSO4 (blue); KMnO4 (purple).
Period
88
Chemical elements in the fifth period Group
1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Name Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe #
Period 5 includes the important metals silver and tin and the biologically important element iodine. Also in period 5 is the lightest purely radioactive element, technetium, the first element to be artificially synthesized.
Chemical elements in the sixth period Group
1
2
55 56 57 58 59 60 61
#
3 (Lanthanides)
62 63 64 65 66 67 68 69
4
5
6
7
8
9 10 11 12 13 14 15 16 17 18
70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Name 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
Period 6 is the first period to include the F block, with the lanthanides aka rare earth elements, and includes the heaviest stable elements. Many of these heavy metals are toxic and some are radioactive, but platinum and gold are largely inert.
Chemical elements in the seventh period Group 1
2
3 (Actinides)
4
5
6
7
8
9 10 11 12 13
14
15
16
17
18
87 88 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 Name 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 #
All elements of period 7 are radioactive. This period contains the heaviest element which occurs naturally on earth, uranium. Most of the subsequent elements in the period have been synthesized artificially. Whilst some of these (e.g. plutonium) are now available in tonne quantities, most are extremely rare, having only been prepared in microgram amounts or less. Some of the later elements have only ever been identified in laboratories in quantities of a few atoms at a time. Although the rarity of many of these elements means that experimental results are not very extensive, periodic and group trends in behaviour appear to be less well defined for period 7 than for other periods. Whilst francium and radium do show typical properties of Groups 1 and 2 respectively, the actinides display a much greater variety of behaviour and oxidation states that than the lanthanides. Initial studies suggest Group 14 element ununquadium appears to be a noble gas instead of a poor metal, and group 18 element ununoctium probably is not a noble gas.[18] These peculiarities of period 7 may be due to a variety of factors, including a large degree of spin-orbit coupling and relativistic effects, ultimately caused by the very high positive electrical charge from their massive atomic nuclei.
Chemical elements in the eighth period No element of the eighth period has yet been synthesized. A G block is predicted. It is not clear if all elements predicted for the eighth period are in fact physically possible. There may therefore be no ninth period. Element categories in the periodic table
Period
89
Metals Alkali metals
Alkaline earth metals
Inner transition elements
Metalloids Transition elements
Other metals
Nonmetals Other nonmetals
Halogens
Unknown chemical Noble properties gases
Lanthanides Actinides
References [1] Palmer, David (November 13, 1997). "Hydrogen in the Universe" (http://imagine.gsfc.nasa.gov/docs/ask_astro/answers/971113i.html). NASA. . Retrieved 2008-02-05. [2] "hydrogen". Encyclopædia Britannica. 2008. [3] "Helium: physical properties" (http://www.webelements.com/helium/physics.html). WebElements. . Retrieved 2008-07-15. [4] "Helium: geological information" (http://www.webelements.com/helium/geology.html). WebElements. . Retrieved 2008-07-15. [5] Cox, Tony (1990-02-03). "Origin of the chemical elements" (http://www.newscientist.com/article/mg12517027. 000-origin-of-the-chemical-elements.html). New Scientist . . Retrieved 2008-07-15. [6] "Helium supply deflated: production shortages mean some industries and partygoers must squeak by.". Houston Chronicle. 2006-11-05. [7] Brown, David (2008-02-02). "Helium a New Target in New Mexico" (http://www.aapg.org/explorer/2008/02feb/helium.cfm). American Association of Petroleum Geologists. . Retrieved 2008-07-15. [8] Lithium (http://www.webelements.com/lithium/) at WebElements. [9] "IARC Monograph, Volume 58" (http://www.inchem.org/documents/iarc/vol58/mono58-1.html). International Agency for Research on Cancer. 1993. . Retrieved 2008-09-18. [10] Information (http://www.chronicberylliumdisease.com/medical/med_bediseases.htm#cbd) about chronic beryllium disease. [11] "Functions of Boron in Plant Nutrition" (http://www.borax.com/agriculture/files/an203.pdf) (PDF). U.S. Borax Inc.. . [12] Blevins, Dale G.; Lukaszewski, Krystyna M. (1998). "Functions of Boron in Plant Nutrition". Annual Review of Plant Physiology and Plant Molecular Biology 49: 481 – 500. doi:10.1146/annurev.arplant.49.1.481. PMID 15012243. [13] Ten most abundant elements in the universe, taken from The Top 10 of Everything, 2006, Russell Ash, page 10. Retrieved October 15, 2008. (http://plymouthlibrary.org/faqelements.htm) [14] Chang, Raymond (2007). Chemistry, Ninth Edition. McGraw-Hill. pp. 52. ISBN 0-07-110595-6. [15] Freitas Jr., Robert A. (1999). Nanomedicine (http://www.foresight.org/Nanomedicine/Ch03_1.html) ,. Landes Bioscience. Tables 3-1 & 3-2. ISBN 1570596808. [16] "Structure and Nomenclature of Hydrocarbons" (http://chemed.chem.purdue.edu/genchem/topicreview/bp/1organic/organic.html). Purdue University. . Retrieved 2008-03-23. [17] Alberts, Bruce; Alexander Johnson, Julian Lewis, Martin Raff, Keith Roberts, Peter Walter. Molecular Biology of the Cell (http://www. ncbi.nlm.nih.gov/books/bv.fcgi?highlight=carbon&rid=mboc4.section.165). Garland Science. . [18] See references in the articles Ununquadium, Ununoctium
Pediod 1
Pediod 1 A period 1 element is one of the chemical elements in the first row (or period) of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The first period contains fewer elements than any other row in the table, with only two: hydrogen and helium. This situation can be explained by modern theories of atomic structure.
Elements Hydrogen Hydrogen (H) is the chemical element with atomic number 1. At standard temperature and pressure, hydrogen is a colorless, odorless, nonmetallic, tasteless, highly flammable diatomic gas with the molecular formula H . With an 2 atomic mass of 1.00794 amu, hydrogen is the lightest element.[1] Hydrogen is the most abundant of the chemical elements, constituting roughly 75% of the universe's elemental mass.[2] Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from hydrocarbons such as methane, after which most elemental hydrogen is used "captively" (meaning locally at the production site), with the largest markets almost equally divided between fossil fuel upgrading, such as hydrocracking, and ammonia production, mostly for the fertilizer market. Hydrogen may be produced from water using the process of electrolysis, but this process is significantly more expensive commercially than hydrogen production from natural gas.[3] The most common naturally occurring isotope of hydrogen, known as protium, has a single proton and no neutrons.[4] In ionic compounds, it can take on either a positive charge, becoming a cation composed of a bare proton, or a negative charge, becoming an anion known as a hydride. Hydrogen can form compounds with most elements and is present in water and most organic compounds. [5] It plays a particularly important role in acid-base chemistry, in which many reactions involve the exchange of protons between soluble molecules.[6] As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and spectrum of [7] the hydrogen atom has played a key role in the development of quantum mechanics. The interactions of hydrogen with various metals are very important in metallurgy, as many metals can suffer hydrogen embrittlement,[8] and in developing safe ways to store it for use as a fuel. [9] Hydrogen is highly soluble in many compounds composed of rare earth metals and transition metals[10] and can be dissolved in both crystalline and amorphous metals.[11] Hydrogen solubility in metals is influenced by local distortions or impurities in the metal crystal lattice.[12]
Helium Helium (He) is a colorless, odorless, tasteless, non-toxic, inert monatomic chemical element that heads the noble gas series in the periodic table and whose atomic number is 2.[13] Its boiling and melting points are the lowest among the elements and it exists only as a gas except in extreme conditions.[14] Helium was discovered in 1868 by French astronomer Pierre Janssen, who first detected the substance as an unknown yellow spectral line signature in light from a solar eclipse.[15] In 1903, large reserves of helium were found in the natural gas fields of the United States, which is by far the largest supplier of the gas.[16] The substance is used in cryogenics,[17] in deep-sea breathing systems,[18] to cool superconducting magnets, in helium dating, [19] for inflating balloons,[20] for providing lift in airships,[21] and as a protective gas for industrial uses such as arc welding [22] and growing silicon wafers. Inhaling a small volume of the gas temporarily changes the timbre and quality of the
90
Pediod 1 human voice.[23] The behavior of liquid helium-4's two fluid phases, helium I and helium II, is important to researchers studying quantum mechanics and the phenomenon of superfluidity in particular, [24] and to those looking at the effects that temperatures near absolute zero have on matter, such as with superconductivity.[25] Helium is the second lightest element and is the second most abundant in the observable universe.[26] Most helium was formed during the Big Bang, but new helium is being created as a result of the nuclear fusion of hydrogen in stars.[27] On Earth, helium is relatively rare and is created by the natural decay of some radioactive elements [28] because the alpha particles that are emitted consist of helium nuclei. This radiogenic helium is trapped with natural [29] gas in concentrations of up to seven percent by volume, from which it is extracted commercially by a [30] low-temperature separation process called fractional distillation.
References • Bloch, D. R. (2006). Organic Chemistry Demystified [31]. McGraw-Hill Professional. ISBN 0-07-145920-0.
References [1] "Hydrogen – Energy" (http://www.eia.doe.gov/kids/energyfacts/sources/IntermediateHydrogen.html). Energy Information Administration. . Retrieved 2008-07-15. [2] Palmer, David (November 13, 1997). "Hydrogen in the Universe" (http://imagine.gsfc.nasa.gov/docs/ask_astro/answers/971113i.html). NASA. . Retrieved 2008-02-05. [3] Staff (2007). "Hydrogen Basics — Production" (http://www.fsec.ucf.edu/en/consumer/hydrogen/basics/production.htm). Florida Solar Energy Center. . Retrieved 2008-02-05. [4] Sullivan, Walter (1971-03-11). "Fusion Power Is Still Facing Formidable Difficulties". The New York Times. [5] "hydrogen". Encyclopædia Britannica. 2008. [6] Eustis, S. N.; Radisic, D; Bowen, KH; Bachorz, RA; Haranczyk, M; Schenter, GK; Gutowski, M (2008-02-15). "Electron-Driven Acid-Base Chemistry: Proton Transfer from Hydrogen Chloride to Ammonia". Science 319 (5865): 936 – 939. doi:10.1126/science.1151614. PMID 18276886. [7] "Time-dependent Schrödinger equation". Encyclopædia Britannica. 2008. [8] Rogers, H. C. (1999). "Hydrogen Embrittlement of Metals". Science 159 (3819): 1057 – 1064. doi:10.1126/science.159.3819.1057. PMID 17775040. [9] Christensen, C. H.; Nørskov, J. K.; Johannessen, T. (July 9, 2005). "Making society independent of fossil fuels — Danish researchers reveal new technology" (http://www.dtu.dk/English/About_DTU/News.aspx?guid={E6FF7D39-1EDD-41A4-BC9A-20455C2CF1A7}). Technical University of Denmark. . Retrieved 2008-03-28. [10] Takeshita, T.; Wallace, W.E.; Craig, R.S. (1974). "Hydrogen solubility in 1:5 compounds between yttrium or thorium and nickel or cobalt". Inorganic Chemistry 13 (9): 2282 – 2283. doi:10.1021/ic50139a050. [11] Kirchheim, R.; Mutschele, T.; Kieninger, W (1988). "Hydrogen in amorphous and nanocrystalline metals". Materials Science and Engineering 99: 457 – 462. doi:10.1016/0025-5416(88)90377-1. [12] Kirchheim, R. (1988). "Hydrogen solubility and diffusivity in defective and amorphous metals". Progress in Materials Science 32 (4): 262 – 325. doi:10.1016/0079-6425(88)90010-2. [13] "Helium: the essentials" (http://www.webelements.com/helium/). WebElements. . Retrieved 2008-07-15. [14] "Helium: physical properties" (http://www.webelements.com/helium/physics.html). WebElements. . Retrieved 2008-07-15. [15] "Pierre Janssen" (http://encarta.msn.com/encyclopedia_762508746/pierre_janssen.html). MSN Encarta. . Retrieved 2008-07-15. [16] Theiss, Leslie (2007-01-18). "Where Has All the Helium Gone?" (http://www.blm.gov/wo/st/en/info/newsroom/2007/january/ NR0701_2. html). Bureau of Land Management. . Retrieved 2008-07-15. [17] Timmerhaus, Klaus D. (2006-10-06). Cryogenic Engineering: Fifty Years of Progress . Springer. ISBN 0-387-33324-X. [18] Copel, M. (September 1966). "Helium voice unscrambling". Audio and Electroacoustics 14 (3): 122 – 126. doi:10.1109/TAU.1966.1161862. [19] "helium dating". Encyclopædia Britannica. 2008. [20] Brain, Marshall. "How Helium Balloons Work" (http://www.howstuffworks.com/helium.htm). How Stuff Works. . Retrieved 2008-07-15. [21] Jiwatram, Jaya (2008-07-10). "The Return of the Blimp" (http://www.popsci.com/military-aviation-space/article/2008-07/ return-blimp). Popular Science. . Retrieved 2008-07-15. [22] "When good GTAW arcs drift; drafty conditions are bad for welders and their GTAW arcs.".Welding Design & Fabrication. 2005-02-01. [23] Montgomery, Craig (2006-09-04). "Why does inhaling helium make one's voice sound strange?" (http://www.sciam.com/article. cfm?id=why-does-inhaling-helium). Scientific American. . Retrieved 2008-07-15. [24] "Probable Discovery Of A New, Supersolid, Phase Of Matter" (http://www.sciencedaily.com/releases/2004/09/040903085531.htm). Science Daily. 2004-09-03. . Retrieved 2008-07-15.
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92
[25] Browne, Malcolm W. (1979-08-21). "Scientists See Peril In Wasting Helium; Scientists See Peril in Waste of Helium".The New York Times. [26] "Helium: geological information" (http://www.webelements.com/helium/geology.html). WebElements. . Retrieved 2008-07-15. [27] Cox, Tony (1990-02-03). "Origin of the chemical elements" (http://www.newscientist.com/article/mg12517027. 000-origin-of-the-chemical-elements.html). New Scientist . . Retrieved 2008-07-15. [28] "Helium supply deflated: production shortages mean some industries and partygoers must squeak by.". Houston Chronicle. 2006-11-05. [29] Brown, David (2008-02-02). "Helium a New Target in New Mexico" (http://www.aapg.org/explorer/2008/02feb/helium.cfm). American Association of Petroleum Geologists. . Retrieved 2008-07-15. [30] Voth, Greg (2006-12-01). "Where Do We Get the Helium We Use?". The Science Teacher. [31] http://books.google.com/?id=yVPcSIn5xjAC
Extensions There are currently seven periods in the periodic table of chemical elements, culminating with atomic number 118. If further elements with higher atomic numbers than this are discovered, they will be placed in additional periods, laid out (as with the existing periods) to illustrate periodically recurring trends in the properties of the elements concerned. Any additional periods are expected to contain a larger number of elements than the seventh period, as they are calculated to have an additional so-called g-block, containing 18 elements with partially filled g-orbitals in each period. An eight-period table containing this block was suggested by Glenn T. Seaborg in 1969.[1] No elements in this region have been synthesized or discovered in nature. (Element 122 was claimed to exist naturally in April 2008, but this claim was widely believed to be erroneous.)[2] The f irst element of the g-block may have atomic number 121, and thus would have the systematic name unbiunium. Elements in this region are likely to be highly unstable with respect to radioactive decay, and have extremely short half lives, although element 126 is hypothesized to be within an island of stability that is resistant to fission but not to alpha decay. It is not clear how many elements beyond the expected island of stability are physically possible, if period 8 is complete, or if there is a period 9. If period 9 does exist, it is likely to be the last. According to the orbital approximation in quantum mechanical descriptions of atomic structure, the g-block would correspond to elements with partially-filled g-orbitals. However, spin-orbit coupling effects reduce the validity of the orbital approximation substantially for elements of high atomic number.[3]
Extended periodic table, including the g-block Extended Periodic Table 1
2
3
4
5
6
7
[4]
1
2
H
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
58
59
60
61
62
63
64
65
66
67
68
69
70
71
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
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
87
88
89
90
91
92
93
94
95
96
97
98
99
100
101
102
103
104
105
106
107
108
109
110
111
112
113
114
115
116
117
118
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
Extensions 8
119
120
121
122
123
93 124
125
126
127
128
129
130 131 132 133 134 135 136 137 138 139
140
141
142
143
144
145
146
147
148
149
150
151
152
153
154
155
156
157
158
159
160
161
162
163
164
165
166
167
168
Uue Ubn Ubu Ubb Ubt Ubq Ubp Ubh Ubs Ubo Ube Utn Utu Utb Utt Utq Utp Uth Uts Uto Ute Uqn Uqu Uqb Uqt Uqq Uqp Uqh Uqs Uqo Uqe Upn Upu Upb Upt Upq Upp Uph Ups Upo Upe Uhn Uhu Uhb Uht Uhq Uhp Uhh Uhs Uho
9
169
170
171
172
173
Uhe
Usn
Usu
Usb
Ust
Blocks of the periodic table s-block p-block d-block f-block g-block (Undiscovered (theorized) elements are coloured in a lighter shade)
All of these hypothetical undiscovered elements are named by the International Union of Pure and Applied Chemistry (IUPAC) systematic element name standard which creates a generic name for use until the element has been discovered, confirmed, and an official name approved. The positioning of the g-block in the table (to the left of the f-block, to the right, or in between) is speculative. The positions shown in the table above corresponds to the assumption that the Madelung rule will continue to hold at higher atomic number; this assumption may or may not be true. At element 118, the orbitals 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, 5s, 5p, 5d, 5f, 6s, 6p, 6d, 7s and 7p are assumed to be filled, with the remaining orbitals unfilled. The orbitals of the eighth period are predicted to be filled in the order 8s, 5g, 6f, 7d, 8p. However, after approximately element 120, the proximity of the electron shells makes placement in a simple table problematic; for example, calculations suggest that it may be elements 165 and 166 which occupy the 9s block (leaving the 8p orbital incomplete) assuming they are physically possible.[5]
End of the periodic table The number of physically possible elements is unknown. The light-speed limit on electrons orbiting in ever-bigger electron shells theoretically limits neutral atoms to a Z of approximately 173,[6] after which it would be nonsensical to assign the elements to blocks on the basis of electron configuration. However, it is likely that the periodic table actually ends much earlier, possibly soon after the island of stability,[7] which is expected to center around Z = 126.[8] Additionally the extension of the periodic and nuclides tables is restricted by the proton drip line and the neutron drip line.
Bohr model breakdown The Bohr model exhibits difficulty for atoms with atomic number greater than 137, for the speed of an electron in a 1s electron orbital, v, is given by
where Z is the atomic number, and α is the fine structure constant, a measure of the strength of electromagnetic interactions.[9] Under this approximation, any element with an atomic number of greater than 137 would require 1s electrons to be traveling swifter than c, the speed of light. Hence a non-relativistic model such as the Bohr model is inadequate for such calculations.
Extensions
The Dirac equation The semi-relativistic Dirac equation also has pr oblems for Z > 137, for the ground state energy is where m0 is the rest mass of the electron. For Z > 137, the wave function of the Dirac ground state is oscillatory, rather than bound, and there is no gap between the positive and negative energy spectra, as in the Klein paradox.[10] Richard Feynman pointed out this effect, so the last element expected under this model, 137 (untriseptium), is sometimes called feynmanium. However, a realistic calculation has to take into account the finite extension of the nuclear-charge distribution. This results in a critical Z of ≈ 173 (unseptrium), such that non-ionized atoms may be limited to elements equal to or lower than this.[6]
See also • Electron configuration • Nuclear shell model • Table of nuclides (combined)
References [1] (http://acs.lbl.gov/Seaborg.talks/65th-anniv/29.html) [2] "Heaviest element claim criticised" (http://www.rsc.org/chemistryworld/News/2008/May/02050802.asp). Rsc.org. 2008-05-02. . Retrieved 2010-03-16. [3] For example, an element in the column labeled g1 may indeed have exactly one valence-shell g-electron (as the name suggests), but it is also possible that it would have more, or none at all. [4] The labels "g1", etc. are inspired by the Madelung rule, but this is merely an empirical rule, with well-known exceptions such as copper. [5] Pekka Pyykkö, Peter Schwerdtfeger (2004), Relativistic electronic structure theory, p 23. [6] Walter Greiner and Stefan Schramm (2008). American Journal of Physics 76: 509. doi:10.1119/1.2820395., and references therein. [7] Encyclopædia Britannica. "transuranium element (chemical element) - Britannica Online Encyclopedia" (http://www.britannica.com/ EBchecked/topic/603220/transuranium-element). Britannica.com. . Retrieved 2010-03-16. [8] S. Cwiok, P.-H. Heenen and W. Nazarewicz (2005). "Shape coexistence and triaxiality in the superheavy nuclei". Nature 433: 705. [9] See for example R. Eisberg and R. Resnick, Quantum Physics of Atoms, Molecules, Solids, Nuclei and Particles , Wiley (New York: 1985). [10] James D. Bjorken and Sidney D. Drell, Relativistic Quantum Mechanics, McGraw-Hill (New York:1964).
• http://www.springerlink.com/content/j303171428652143/
External links • Images of g-orbitals (http://www.uky.edu/~holler/html/g.html) from the University of Kentucky • jeries.rihani.com (http://jeries.rihani.com) - The extended periodic table of the elements. • Eric Scerri, The Periodic Table, Its Story and Its Significance, Oxford University Press, 2007.
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95
Blocks Block A block of the periodic table of elements is a set of adjacent groups. The term appears to have been first used (in French) by Charles Janet. [1] The respective highest-energy electrons in each element in a block belong to the same atomic orbital type. Each block is named after its characteristic orbital; thus, the blocks are: • • • • •
s-block p-block d-block f-block g-block (hypothetical)
The block names (s, p, d, f. and g) are derived from the quality of the spectroscopic lines of the associated atomic orbitals: sharp, principal, diffuse and f undamental, the rest being named in alphabetical order. Blocks are sometimes called families. [1] Charles Janet, La classification helicoidal des elements chimiques , Beauvais, 1928
s-block Chemical elements in s-block Group
1
2
18
Period 1
1 H
2
3 4 Li Be
3
11 12 Na Mg
4
19 20 K Ca
5
37 38 Rb Sr
6
55 56 Cs Ba
7
87 88 Fr Ra
2 He
The s-block of the periodic table of elements consists of the first two groups: the alkali metals and alkaline earth metals, plus hydrogen and helium. These elements are distinguished by the property that in the atomic ground state, the highest-energy electron is in an s-orbital. Except in hydrogen and helium, these electrons are very easily lost to form positive ions. The helium
s-block
96
configuration is chemically exceedingly stable and thus helium has no known stable compounds; thus it is generally grouped with the noble gases. The other elements of the s-block are all extremely powerful reducing agents, so much so that they never occur naturally in the free state. The metallic metallic forms forms of these elements elements can only be extracted extracted by electrolysis of a molten molten salt, since water is much more easily reduced to hydrogen than the ions of these metals. Sir Humphry Davy, in 1807 and 1808, was the first to isolate all of these metals except lithium, beryllium, rubidium and caesium. Beryllium was isolated independently by F. Wooler and A.A. Bussy in 1828, while lithium was isolated by Robert Bunsen in 1854, who isolated rubidium nine years later after having observed it and caesium spectroscopically. Caesium was not isolated until 1881 when Carl Setterberg electrolysed the molten cyanide. The s-block metals vary from extremely soft (all the alkali metals) to quite hard (beryllium). With the exception of beryllium and magnesium, the metals are too reactive for any structural use except as very minor components (5%
>1% >.1%
Pd
>7%
Sn
I
fission product yield
Only thorium and uranium occur naturally in the Earth's crust in anything more than trace quantities. Protactinium and actinium, which are both decay products of uranium, are the only remaining actinides that were discovered in nature before they were synthesized. Neptunium and plutonium have also been known to show up naturally in trace amounts in uranium ores as a result of decay or bombardment, but this was only discovered after they were synthesized. The remaining actinides were synthesized in particle colliders or nuclear reactors, and none of them has been found to occur naturally on earth. Actinides beyond californium possess exceedingly short half-lives.
Actinide Isotopes of all of the transuranium elements up to and including fermium can be produced by rapid neutron bombardment of lighter nuclides. The nuclei created have an excess of neutrons. β-decay occurs with a neutron decaying to a proton and an electron, increasing the atomic number in the process. Conditions suitable for the synthesis of transuranium elements occur in supernovae. These elements may also be produced in specialized nuclear reactors. They may be created in a nuclear explosion and come to earth as nuclear fallout from an atmospheric test explosion. The heavier elements may be synthesized by bombardment with heavier particles, such as α particles or heavier nuclei. In 1961, Antoni Przybylski discovered a star, HD 101065, commonly called Przybylski's star, that contains unusually high amounts of actinides.
See also • Actinides in the environment
Further reading • Tamer Andrea and Moris S. Eisen (2008). "Recent advances in organothorium and organouranium catalysis". Chem. Soc. Rev. 37: 550 - 567. doi:10.1039/b614969n. • Morss, Lester R.; Edelstein, Norman M.; Fuger, Jean, eds (2006). The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands: Springer. ISBN 13978-1-4020-3555-5.
External links • The Columbia Encyclopedia, Sixth Edition. [8] • Chemical Elements website [9] • Lawrence Berkeley Laboratory image of historic periodic table by Seaborg showing actinide series for the first time [10] • Lawrence Livermore National Laboratory, Uncovering the Secrets of the Actinides [11] • Los Alamos National Laboratory, Actinide Research Quarterly [12]
References [1] [2] [3] [4]
IUPAC Periodic Table (http://www.iupac.org/reports/periodic_table) IUPAC Periodic Table 2007 .pdf (http://www.iupac.org/reports/periodic_table/IUPAC_Periodic_Table-22Jun07b.pdf) Connelly, Neil G.; et al. (2005). "Elements". Nomenclature of Inorganic Chemistry . London: Royal Society of Chemistry. pp. 52. Seaborg, Glenn T. (1946). "The Transuranium Elements" (http://www. jstor.org/stable/1675046). Science 104 (2704): 379 – 386. doi:10.1126/science.104.2704.379. . [5] Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419, p 1263 [6] Arnold F. Holleman, Nils Wiberg: Lehrbuch der Anorganischen Chemie, 102. Auflage, de Gruyter, Berlin 2007, S. 1956; ISBN 978-3-11-017770-1. [7] dtv-Atlas zur Chemie 1981, Teil 1, S. 224. [8] http://www.bartleby.com/65/ac/actinide.html [9] http://www.chemicalelements.com/groups/rareearth.html [10] http://imglib.lbl.gov/ImgLib/COLLECTIONS/BERKELEY-LAB/SEABORG-ARCHIVE/index/96B05654.html [11] http://www.llnl.gov/str/pdfs/06_00.2.pdf#search=%22actinide%20series%22 [12] http://arq.lanl.gov/
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Lanthanide The lanthanide or lanthanoid (IUPAC nomenclature)[1] series comprises the fourteen elements with atomic numbers 58 through 71, from cerium to lutetium.[2] All lanthanides are f-block elements, corresponding to the filling of the 4f electron shell. Lanthanum, which is a d-block element, may also be considered to be a lanthanide. All lanthanide elements form trivalent cations, Ln3+, whose chemistry is largely determined by the ionic radius, which decreases steadily from lanthanum to lutetium.
Classification Atomic No.
57 58 59 60 61 62 63 64 65 66 67 68 69
Name
La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
M3+ f electrons 0 1
2
3
4
5
6
7
8
9
10 11 12
70 71
13 14
The lanthanide elements are the group of elements with atomic number increasing from 58 (cerium) to 71 (lutetium). They are termed lanthanide because the lighter elements in the series are chemically similar to lanthanum. Strictly speaking lanthanum is a group 3 element element and the ion La 3+ has no f electrons. However this element is often included in any general discussion of the chemistry of the lanthanide elements.
Chemistry The electronic structure of the lanthanide elements, with minor exceptions is [Xe]6s24f n. In their compounds, the 6s electrons are lost and the ions have the configuration [Xe]4f m.[3] The chemistry of the lanthanides differs from main group elements and transition metals because of the nature of the 4f orbitals. These orbitals are "buried" inside the atom and are shielded from the atom's environment by the 4d and 5p electrons. As a consequence of this the chemistry of the elements is largely determined by their size, which decreases gradually from 102 pm (La 3+) with increasing atomic number to 86 pm (Lu3+), the so-called lanthanide contraction. All the lanthanide elements exhibit the oxidation state +3. In addition Ce3+ can lose its single f electron to form Ce4+ with the stable electronic configuration of xenon. Also, Eu3+ can gain an electron to form Eu 2+ with the f 7 configuration which has the extra stability of a half-filled shell. Promethium is effectively a man-made element as all its isotopes are radioactive with half-lives of less than 20 y. The similarity in ionic radius between adjacent lanthanide elements makes it difficult to separate them from each other in naturally occurring ores and other mixtures. Historically the very laborious processes of cascading and fractional crystallization was used. Because the lanthanide ions have slightly different radii, the lattice energy of their salts and hydration energies of the ions will be slightly different, leading to a small difference in solubility. Salts of the formula Ln(NO3)3.2NH4NO3.4H2O can be used. Industrially, the elements are separated from each other by solvent extraction. Typically an aqueous solution of nitrates is extracted into kerosene containing tri-n-butylphosphate, (BunO)3PO. The strength of the complexes formed increases as the ionic radius decreases, so solubility in the organic phase increases. Complete separation can be achieved continuously by use of countercurrent exchange methods. The elements can also be separated by ion-exchange chromatography, making use of the fact that the stability constant for formation of EDTA complexes increases for log K ≈ 15.5 for [La(EDTA)]- to log K ≈ 19.8 - [4] [5] for [Lu(EDTA)] . The process, involving two columns, is described in detail in Greenwood & Earnshaw Ce(IV) is a useful oxidising agent, and Eu(II) is a useful reducing agent. The trivalent lanthanides mostly form ionic salts. The trivalent ions are hard acceptors and form more stable complexes with oxygen-donor ligands than with nitrogen-donor ligands. The larger ions are 9-coordinate in aqueous solution, [Ln(H 2O)9]3+ but the smaller ions are 8-coordinate, [Ln(H2O)8]3+. There is some evidence that the later lanthanides have more water molecules in the
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second coordination sphere.[6] Complexation with monodentate ligands is generally weak because it is difficult to displace water molecules from the first coordination sphere. Stronger complexes are formed with chelating ligands because of the chelate effect.
Magnetic and spectroscopic spectroscopic properties All the trivalent lanthanide ions, except lutetium, have unpaired f electrons. However the magnetic moments deviate considerably from the spin-only values because of strong spin-orbit coupling. The maximum number of unpaired electrons is 7, in Gd3+, with a magnetic moment of 7.94 B.M., but the largest magnetic moments, at 10.4-10.7 B.M., are exhibited by Dy3+ and Ho3+. However, in Gd3+ all the electrons have parallel spin and this property is important for the use of gadolinium complexes as contrast reagent in MRI scans. Crystal field splitting is rather small for the lanthanide ions and is less important than spin-orbit coupling in regard to energy levels.[7] Transitions of electrons between f orbitals are forbidden by the Laporte rule. Furthermore, because of the "buried" nature of the f orbitals, coupling with molecular vibrations is weak. Consequently the spectra of lanthanide ions are rather weak and the absorption bands are similarly narrow. Glass containing holmium oxide and holmium oxide solutions (usually in perchloric acid) have sharp optical absorption peaks in the spectral range 200 – 900 900 nm and can be used as a wavelength calibration standard for optical spectrophotometers[8] , and are available commercially.[9]
A solution of 4% holmium oxide in 10% perchloric acid, permanently fused into a quartz cuvette as a wavelength calibration standard
As f-f transitions are Laporte-forbidden, once an electron has been excited, decay to the ground state will be slow. This makes them suitable for use in lasers as it makes the population inversion easy to achieve. The Nd:YAG laser is one that is widely used. Lanthanide ions are also fluorescent as a result of the forbidden nature of f-f transitions. Europium-doped yttrium vanadate was the first red phosphor to enable the development of colour television screens.[10]
Lanthanide
Organometallic chemistry Metal-carbon σ bonds are found in alkyls of the lanthanide elements such as [LnMe ]3- and Ln[CH(SiMe ) ].[11] The 6 3 3 cyclopentadiene complexes, of formula [Ln(C5H5)3] and [Ln(C5H5)2Cl] may have η-1, η-2, and η-5 rings. Analogues to uranocene are formed with the cyclo-octadienide ion, C8H82- which is a Hückel's rule aromatic ring.
Geochemistry The trivial name "rare earths" is sometimes used to describe all the lanthanides together with scandium and yttrium. This name arises from the minerals from which they were isolated, which were uncommon oxide-type minerals. However, the use of the name is deprecated by IUPAC, as the elements are neither rare in abundance nor "earths" (an obsolete term for water-insoluble strongly basic oxides of electropositive metals incapable of being smelted into metal using late 18th century technology) . Cerium is the 26th most abundant element in the Earth's crust, Abundance of elements in the Earth crust per million of Si atoms neodymium is more abundant than gold and even thulium (the least common naturally occurring lanthanide) is more abundant than iodine.[12] Despite their abundance, even the technical term "lanthanides" could be interpreted to reflect a sense of elusiveness on the part of these elements, as it comes from the Greek λανθανειν (lanthanein), "to lie hidden". However, if not referring to their natural abundance, but rather to their property of "hiding" behind each other in minerals, this interpretation is in fact appropriate. The etymology of the term must be sought in the first discovery of lanthanum, at that time a so-called new rare earth element "lying hidden" in a cerium mineral, but we might call it a fortunate twist of irony that exactly lanthanum was later identified as the first in an entire series of chemically similar elements and could give name to the whole series. The lanthanide contraction is responsible for the great geochemical divide that splits the lanthanides into light and heavy-lanthanide enriched minerals, the latter being almost inevitably associated with and dominated by yttrium. This divide is reflected in the first two "rare earths" that were discovered: yttria (1794) and ceria (1803). The geochemical divide has put more of the light lanthanides in the Earth's crust, but more of the heavy members in the Earth's mantle. The result is that although large rich ore-bodies are found that are enriched in the light lanthanides, correspondingly large ore-bodies for the heavy members are few. The principal ores are monazite and bastnaesite. Monazite sands usually contain all the lanthanide elements, but the heavier elements are lacking in bastnaesite. The lanthanides obey the Oddo-Harkins rule - odd-numbered elements are less abundant than their even-numbered neighbours. Three of the lanthanide elements have radioactive isotopes with long half-lives ( 138La, 147Sm and 176Lu) that can be used to date minerals and rocks from Earth, the Moon and meteorites.[13]
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Biological effects Lanthanides entering the human body due to exposure to various industrial processes can affect metabolic processes. Trivalent lanthanide ions, especially La3+ and Gd3+, can interfere with calcium channels in human and animal cells. Lanthanides can also alter or even inhibit the action of various enzymes. Lanthanide ions found in neurons can regulate synaptic transmission, as well as block some receptors (for example, glutamate receptors).[14]
Applications Most lanthanides are widely used in lasers. These elements deflect ultraviolet and infrared radiation and are commonly used in the production of sunglass lenses. Other applications are summarized in the following table:[12] Application
Percentage
Catalytic converters
45
Petrol Petroleum eum refini refining ng cataly catalysts sts
25
Permanent magnets
12
Glass polishing and ceramics
7
Metallurgical
7
Phosphors
3
Other
1
See also • • • •
Actinoid Grou Groupp 3 eleme lement nt Lant Lantha hani nide de con contr trac acti tion on Rare Rare eart earthh ele eleme ment nt
External links • lant lantha hani nide de Spa Spark rkle le Mod Model el [15], used in the computational chemistry of lanthanoid complexes • USGS USGS Rare Eart Earths hs Statis Statistic ticss and Inform Informati ation on [16] • Ana de Bettencour Bettencourt-Dias t-Dias:: Chemistry Chemistry of the lanthanides lanthanides and and lanthanide-co lanthanide-contain ntaining ing materials materials [17]
References [1] the current IUPAC recommendation recommendation is that the name name lanthanoid be used rather than lanthanide, as the suffix "-ide" is preferred for negative ions whereas the suffix "-oid" indicates similarity to one of the members of the containing family of elements. However, lanthanide is still favored in most (~90%) scientific articles and is currently adopted on wikipedia. In the older literature, the name "lanthanon" was often used. [2] Holden, Norman E.; E.; Coplen, Coplen, Tyler (January-February 2004). The Periodic Table of the Elements (IUPAC) 26 (1): 8. http://www.iupac. http://www.iupac.org/ org/ publications/ci/2004/2601/2_holden.html. publications/ci/2004/2601/2_ho lden.html. Retrieved March 23, 2010. Lanthanum ionisation energies (http://www.webelements.com/lanthanum/atoms. [3] [3] Mark Mark Wint Winter er.. (http://www.webelements.com/lanthanum/atoms.html). html). WebElements Ltd, UK. . Retrieved 02-09-2010. [4] L. Pettit Pettit and K. Powell, Powell, SC-database SC-database (http://www (http://www..acadsoft.co. acadsoft.co.uk/scdbase/scdbase. uk/scdbase/scdbase.htm) htm) [5] Greenwood Greenwood,, Norman N.; Earnsha Earnshaw, w, A. (1997), (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419 p 1231 [6] Burge Burgess, ss,,, J. (1978). (1978). 'Metal ions in solution' . , New York: Ellis Horwood. ISBN 0853120277. [7] Greenwood Greenwood,, Norman N.; Earnsha Earnshaw, w, A. (1997), (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419 p 1242
Lanthanide [8] R. P. MacDonald (1964). "Uses for a Holmium Oxide Filter in Spectrophotometry" (http://www.clinchem.org/cgi/reprint/10/12/1117. pdf). Clinical Chemistry 10: 1117. . [9] "Holmium Glass Filter for Spectrophotometer Calibration" (http://www.labshoponline.com/ holmium-glass-filter-spectrophotometer-calibration-p-88.html). . Retrieved 2009-06-06. [10] Levine, Albert K.; Palilla, Frank C. (1964). "A new, highly efficient red-emitting cathodoluminescent phosphor (YVO4:Eu) for color television". Applied Physics Letters 5: 118. doi:10.1063/1.1723611. [11] Cotton, S.A. (1997). "Aspects of the lanthanide-carbon σ-bond". Coord. Chem. Revs. 160: 93 – 127. doi:10.1016/S0010-8545(96)01340-9. [12] Helen C. Aspinall (2001). Chemistry of the f-block elements (http://books.google.com/?id=bLI2maI1_xAC). CRC Press. p. 8. ISBN 905699333X. . [13] There exist other naturally occurred radioactive isotopes of lanthanides with long half-lives (144Nd, 150Nd, 148Sm, 151Eu, 152Gd) but they are not used as chronometers. [14] Pałasz, A; Czekaj, P (2000). "Toxicological and cytophysiological aspects of lanthanides action." (http://www.actabp.pl/pdf/4_2000/ 1107-1114s.pdf). Acta biochimica Polonica 47 (4): 1107 – 14. PMID 11996100. . [15] http://www.sparkle.pro.br [16] http://minerals.usgs.gov/minerals/pubs/commodity/rare_earths/ [17] http://www.chem.unr.edu/faculty/abd/
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Metal Alkali metals
Lithium, Sodium, Potassium Rubidium, Caesium, Francium Alkaline earth metals
Beryllium, Magnesium, Calcium Strontium, Barium, Radium Transition metals
Zinc, Molybdenum, Cadmium Scandium, Titanium, Vanadium Chromium, Manganese, Iron Cobalt, Nickel, Copper Yttrium, Zirconium, Niobium Technetium, Ruthenium, Rhodium Palladium, Silver, Hafnium Tantalum, Tungsten, Rhenium Osmium, Iridium, Platinum Gold, Mercury, Rutherfordium, Dubnium, Seaborgium, Bohrium, Hassium, Meitnerium, Darmstadtium, Roentgenium, Copernicium Post-transition metals
Aluminium, Gallium, Indium Tin, Thallium, Lead, Bismuth Ununtrium, Ununquadium Ununpentium, Ununhexium Lanthanoids
Lanthanum, Cerium, Praseodymium Neodymium, Promethium, Samarium Europium, Gadolinium, Terbium Dysprosium, Holmium, Erbium Thulium, Ytterbium, Lutetium Actinoids
Actinium, Thorium, Protactinium Uranium, Neptunium, Plutonium Americium, Curium, Berkelium Californium, Einsteinium, Fermium Mendelevium, Nobelium, Lawrencium
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A metal is a chemical element that is a good conductor of both electricity and heat and forms cations and ionic bonds with non-metals. In chemistry, a metal (from Greek "μέταλλον" - métallon, "mine"[1] ) is an element, compound, or alloy characterized by high electrical conductivity. In a metal, atoms readily lose electrons to form positive ions (cations). Those ions are surrounded by delocalized electrons, which are responsible for the conductivity. The solid thus produced is held by electrostatic interactions between the ions and the electron cloud, which are called metallic bonds.[2] Usage in astronomy is quite different.
Definition Metals are sometimes described as an arrangement of positive ions surrounded by a sea of delocalized electrons. They are one of the three groups of elements as distinguished by their ionization and bonding properties, along with the metalloids and non-metals. Metals occupy the bulk of the periodic table, while non-metallic elements can only be found on the right-hand-side of the Periodic Table of the Elements. A diagonal line drawn from boron (B) to polonium (Po) separates the metals from the nonmetals. Most elements on this line are metalloids, sometimes called semiconductors. This is because these elements exhibit electrical properties common to both conductors and insulators. Elements to the lower left of this division line are called metals, while elements to the upper right of the division line are called non-metals. An alternative definition of metal refers to the band theory. If one fills the energy bands of a material with available electrons and ends up with a top band partly filled then the material is a metal. This definition opens up the category for metallic polymers and other organic metals, which have been made by researchers and employed in high-tech devices. These synthetic materials often have the characteristic silvery gray reflectiveness (luster) of elemental metals.
Astronomy In the specialized usage of astronomy and astrophysics, the term "metal" is often used to refer collectively to all elements other than hydrogen or helium, including substances as chemically non-metallic as neon, fluorine, and oxygen. Nearly all the hydrogen and helium in the Universe was created in Big Bang nucleosynthesis, whereas all the "metals" were produced by nucleosynthesis in stars or supernovae. The Sun and the Milky Way Galaxy are composed of roughly 74% hydrogen, 24% helium, and 2% "metals" (the rest of the elements; atomic numbers 3-118) [3] by mass. The concept of a metal in the usual chemical sense is irrelevant in stars, as the chemical bonds that give elements their properties cannot exist at stellar temperatures.
Properties Chemical Metals are usually inclined to form cations through electron loss,[2] reacting with oxygen in the air to form oxides over changing timescales (iron rusts over years, while potassium burns in seconds). Examples: 4 Na + O2 → 2 Na2O (sodium oxide) 2 Ca + O2 → 2 CaO (calcium oxide) 4 Al + 3 O 2 → 2 Al2O3 (aluminium oxide) The transition metals (such as iron, copper, zinc, and nickel) take much longer to oxidize. Others, like palladium, platinum and gold, do not react with the atmosphere at all. Some metals form a barrier layer of oxide on their surface which cannot be penetrated by further oxygen molecules and thus retain their shiny appearance and good
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conductivity for many decades (like aluminium, magnesium, some steels, and titanium). The oxides of metals are generally basic, as opposed to those of nonmetals, which are acidic. Painting, anodizing or plating metals are good ways to prevent their corrosion. However, a more reactive metal in the electrochemical series must be chosen for coating, especially when chipping of the coating is expected. Water and the two metals form an electrochemical cell, and if the coating is less reactive than the coatee, the coating actually promotes corrosion.
Physical Metals in general have high electrical conductivity, thermal conductivity, luster and density, and the ability to be deformed under [2] stress without cleaving. While there are several metals that have low density, hardness, and melting points, these (the alkali and alkaline earth metals) are extremely reactive, and are rarely encountered in their elemental, metallic form. Optically speaking, metals are opaque, shiny and lustrous. This is because visible lightwaves are not readily transmitted through the bulk of their microstructure. The large number of free electrons in any typical metallic solid (element or alloy) is responsible for the fact that they can never be categorized as transparent materials.
Gallium crystals
The majority of metals have higher densities than the majority of nonmetals. [2] Nonetheless, there is wide variation in the densities of metals; lithium is the least dense solid element and osmium is the densest. The metals of groups I A and II A are referred to as the light metals because they are exceptions to this generalization [2] . The high density of most metals is due to the tightly packed crystal lattice of the metallic structure. The strength of metallic bonds for different metals reaches a maximum around the center of the transition series, as those elements have large amounts of delocalized electrons in a metallic bond. However, other factors (such as atomic radius, nuclear charge, number of bonding orbitals, overlap of orbital energies, and crystal form) are involved as well.[2]
Electrical The electrical and thermal conductivity of metals originate from the fact that in the metallic bond, the outer electrons of the metal atoms form a gas of nearly free electrons, moving as an electron gas in a background of positive charge formed by the ion cores. Good mathematical predictions for electrical conductivity, as well as the electrons' contribution to the heat capacity and heat conductivity of metals can be calculated from the free electron model, which does not take the detailed structure of the ion lattice into account. When considering the exact band structure and binding energy of a metal, it is necessary to take into account the positive potential caused by the specific arrangement of the ion cores - which is periodic in crystals. The most important consequence of the periodic potential is the formation of a small band gap at the boundary of the Brillouin zone. Mathematically, the potential of the ion cores can be treated by various models, the simplest being the nearly free electron model.
Mechanical Mechanical properties of metals include ductility, which is largely due to their inherent capacity for plastic deformation. Reversible elasticity in metals can be described by Hooke's Law for restoring forces, where the stress is linearly proportional to the strain. Forces larger than the elastic limit, or heat, may cause a permanent (irreversible) deformation of the object, known as plastic deformation or plasticity. This irreversible change in atomic arrangement may occur as a result of:
Metal • The action of an applied force (or work). An applied force may be tensile (pulling) force, compressive (pushing) force, shear, bending or torsion (twisting) forces. • A change in temperature (or heat). A temperature change may affect the mobility of the structural defects such as grain boundaries, point vacancies, line and screw dislocations, stacking faults and twins in both crystalline and non-crystalline solids. The movement or displacement of such mobile defects is thermally activated, and thus limited by the rate of atomic diffusion. Viscous flow near grain boundaries, for example, can give rise to internal slip, creep and fatigue in metals. It can also contribute to significant changes in the microstructure like grain growth and localized densification due to the elimination of intergranular porosity. Screw dislocations may slip in the direction of any lattice plane containing the dislocation, while the principal driving force for "dislocation climb" is the movement or diffusion of vacancies through a crystal lattice. In addition, the nondirectional nature of metallic bonding is also thought to Hot metal work from a blacksmith. contribute significantly to the ductility of most metallic solids. When the planes of an ionic bond slide past one another, the resultant change in location shifts ions of the same charge into close proximity, resulting in the cleavage of the crystal; such shift is not observed in covalently bonded crystals where fracture and crystal fragmentation occurs.[4]
Alloys An alloy is a mixture of two or more elements in solid solution in which the major component is a metal. Most pure metals are either too soft, brittle or chemically reactive for practical use. Combining different ratios of metals as alloys modifies the properties of pure metals to produce desirable characteristics. The aim of making alloys is generally to make them less brittle, harder, resistant to corrosion, or have a more desirable color and luster. Of all the metallic alloys in use today, the alloys of iron (steel, stainless steel, cast iron, tool steel, alloy steel) make up the largest proportion both by quantity and commercial value. Iron alloyed with various proportions of carbon gives low, mid and high carbon steels, with increasing carbon levels reducing ductility and toughness. The addition of silicon will produce cast irons, while the addition of chromium, nickel and molybdenum to carbon steels (more than 10%) results in stainless steels. Other significant metallic alloys are those of aluminium, titanium, copper and magnesium. Copper alloys have been known since prehistory —bronze gave the Bronze Age its name —and have many applications today, most importantly in electrical wiring. The alloys of the other three metals have been developed relatively recently; due to their chemical reactivity they require electrolytic extraction processes. The alloys of aluminium, titanium and magnesium are valued for their high strength-to-weight ratios; magnesium can also provide electromagnetic shielding. These materials are ideal for situations where high strength-to-weight ratio is more important than material cost, such as in aerospace and some automotive applications. Alloys specially designed for highly demanding applications, such as jet engines, may contain more than ten elements.
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Categories Base metal In chemistry, the term base metal is used informally to refer to a metal that oxidizes or corrodes relatively easily, and reacts variably with dilute hydrochloric acid (HCl) to form hydrogen. Examples include iron, nickel, lead and zinc. Copper is considered a base metal as it oxidizes relatively easily, although it does not react with HCl. It is commonly used in opposition to noble metal. In alchemy, a base metal was a common and inexpensive metal, as opposed to precious metals, mainly gold and silver. A longtime goal of the alchemists was the transmutation of base metals into precious metals. In numismatics, coins used to derive their value primarily from the precious metal content. Most modern currencies are fiat currency, allowing the coins to be made of base metal.
Ferrous metal The term "ferrous" is derived from the Latin word meaning "containing iron". This can include pure iron, such as wrought iron, or an alloy such as steel. Ferrous metals are often magnetic, but not exclusively.
Noble metal Noble metals are metals that are resistant to corrosion or oxidation, unlike most base metals. They tend to be precious metals, often due to perceived rarity. Examples include tantalum, gold, platinum, silver and rhodium.
Precious metal A precious metal is a rare metallic chemical element of high economic value. Chemically, the precious metals are less reactive than most elements, have high luster and high electrical conductivity. Historically, precious metals were important as currency, but are now regarded mainly as investment and industrial commodities. Gold, silver, platinum and palladium each have an ISO 4217 currency code. The best-known precious metals are gold and silver. While both have industrial uses, A gold nugget they are better known for their uses in art, jewelry, and coinage. Other precious metals include the platinum group metals: ruthenium, rhodium, palladium, osmium, iridium, and platinum, of which platinum is the most widely traded. Plutonium and uranium could also be considered precious metals. The demand for precious metals is driven not only by their practical use, but also by their role as investments and a store of value. Palladium was, as of summer 2006, valued at a little under half the price of gold, and platinum at around twice that of gold. Silver is substantially less expensive than these metals, but is often traditionally considered a precious metal for its role in coinage and jewelry.
Metal
Extraction Metals are often extracted from the Earth by means of mining, resulting in ores that are relatively rich sources of the requisite elements. Ore is located by prospecting techniques, followed by the exploration and examination of deposits. Mineral sources are generally divided into surface mines, which are mined by excavation using heavy equipment, and subsurface mines. Once the ore is mined, the metals must be extracted, usually by chemical or electrolytic reduction. Pyrometallurgy uses high temperatures to convert ore into raw metals, while hydrometallurgy employs aqueous chemistry for the same purpose. The methods used depend on the metal and their contaminants. When a metal ore is an ionic compound of that metal and a non-metal, the ore must usually be smelted — heated with a reducing agent — to extract the pure metal. Many common metals, such as iron, are smelted using carbon as a reducing agent. Some metals, such as aluminium and sodium, have no commercially practical reducing agent, and are extracted using electrolysis instead.[5] Sulfide ores are not reduced directly to the metal but are roasted in air to convert them to oxides.
Metallurgy Metallurgy is a domain of materials science that studies the physical and chemical behavior of metallic elements, their intermetallic compounds, and their mixtures, which are called alloys.
Applications Some metals and metal alloys possess high structural strength per unit mass, making them useful materials for carrying large loads or resisting impact damage. Metal alloys can be engineered to have high resistance to shear, torque and deformation. However the same metal can also be vulnerable to fatigue damage through repeated use or from sudden stress failure when a load capacity is exceeded. The strength and resilience of metals has led to their frequent use in high-rise building and bridge construction, as well as most vehicles, many appliances, tools, pipes, non-illuminated signs and railroad tracks. The two most commonly used structural metals, iron and aluminium, are also the most abundant metals in the Earth's crust.[6] Metals are good conductors, making them valuable in electrical appliances and for carrying an electric current over a distance with little energy lost. Electrical power grids rely on metal cables to distribute electricity. Home electrical systems, for the most part, are wired with copper wire for its good conducting properties. The thermal conductivity of metal is useful for containers to heat materials over a flame. Metal is also used for heat sinks to protect sensitive equipment from overheating. The high reflectivity of some metals is important in the construction of mirrors, including precision astronomical instruments. This last property can also make metallic jewelry aesthetically appealing. Some metals have specialized uses; radioactive metals such as uranium and plutonium are used in nuclear power plants to produce energy via nuclear fission. Mercury is a liquid at room temperature and is used in switches to complete a circuit when it flows over the switch contacts. Shape memory alloy is used for applications such as pipes, fasteners and vascular stents.
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Trade The World Bank reports that China was the top importer of ores and metals in 2005 followed by the U.S.A. and Japan.[7]
See also • • • • • • • • • • • •
Amorphous metal ASM International (society) Ductility Electric field screening Metal theft Metalworking Periodic table (metals and non-metals) Properties and uses of metals Solid Steel Structural steel Transition metal
Metal and ore imports in 2005
External links • Martindale's 'The Reference Desk' - International Art, Business, Science & Technology [8]
References [1] μέταλλον (http://www.perseus.tufts.edu/hopper/text?doc=Perseus:text:1999.04.0057:entry=me/tallon), Henry George Liddell, Robert Scott, A Greek-English Lexicon , on Perseus Digital Library [2] Mortimer, Charles E. (1975). Chemistry: A Conceptual Approach (3rd ed.). New York:: D. Van Nostrad Company. [3] Sparke, Linda S.; Gallagher, John S. (2000). Galaxies in the Universe (1 ed.). Cambridge University Press. p. 8. ISBN 0521592410. [4] Ductility - strength of materials (http://www.engineersedge.com/material_science/ductility.htm) [5] "Los Alamos National Laboratory – Sodium" (http://periodic.lanl.gov/elements/11.html). . Retrieved 2007-06-08. [6] Frank Kreith and Yogi Goswami, eds. (2004). The CRC Handbook of Mechanical Engineering, 2nd edition. Boca Raton. p. 12-2. [7] Structure of merchandise imports (http://siteresources.worldbank.org/DATASTATISTICS/Resources/table4_5.pdf) [8] http://www.martindalecenter.com/GradMaterial_4_MA.html
Metalloid
116
Metalloid 2
3
4
5
6
13
14
15
16
17
B
C
N
O
F
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Al
Si
P
S
Cl
Aluminium
Silicon
Phosphorus
Sulfur
Chlorine
Ga
Ge
As
Se
Br
Gallium
Germanium
Arsenic
In
Sn
Sb
Te
I
Indium
Tin
Antimony
Tellurium
Iodine
Tl
Pb
Bi
Po
At
Thallium
Lead
Bismuth
Selenium Bromine
Polonium Astatine
Metalloid, or semi metal, is a term used in chemistry when classifying the chemical elements. On the basis of their
general physical and chemical properties, nearly every element in the periodic table can be termed either a metal or a nonmetal. However, a few elements with intermediate properties are referred to as metalloids (from the Greek metallon = "metal" and eidos = "sort"). The line that separates metalloids from nonmetals in the periodic table is referred to as the "amphoteric line". There is no rigorous definition of the term, but the following properties are usually considered characteristic of metalloids: • metalloids often form amphoteric oxides. • metalloids often behave as semiconductors (B, Si, Ge). The concepts of metalloid and semiconductor should not be confused. Metalloid refers to the properties of certain elements in relation to the periodic table. Semiconductor refers to the physical properties of materials (including alloys, compounds) and there is only partial overlap between the two. The following elements are generally considered metalloids:[1] [2] • • • • • • •
Boron (B) Silicon (Si) Germanium (Ge)[3] [4] Arsenic (As)[5] Antimony (Sb)[5] Tellurium (Te)[5] [6] Polonium (Po)[7] [8]
Some allotropes of elements exhibit more pronounced metal, metalloid or non-metal behavior than others. For example, for the element carbon, its diamond allotrope is clearly non-metallic, but the graphite allotrope displays limited electric conductivity more characteristic of a metalloid. Phosphorus, tin, and bismuth also have allotropes that display borderline behavior. In the standard layout of the periodic table, metalloids occur along the diagonal line through the p block from boron to astatine. Elements to the upper right of this line display increasing nonmetallic behaviour; elements to the lower left display increasing metallic behaviour. This line is called the "stair-step" or "staircase." The poor metals are to the left and down and the nonmetals are to the right and up.
Metalloid
117
References [1] E. Sherman and G.J. Weston (1966). Chemistry of the non-metallic elements . Pergamon Press, New York. p. 64. [2] Boylan, P.J. (1962). Elements of Chemistry . Allyn and Bacon, Boston. p. 493. [3] Liu, E (1978). "Fluorination of dimethylmercury, tetramethylsilane and tetramethylgermanium. Synthesis and characterization of polyfluorotetramethylsilanes, polyfluorotetramethylgermanes,bis(trifluoromethyl)mercury and tetrakis(trifluoromethyl)germanium". Journal of Organometallic Chemistry 145: 167. doi:10.1016/S0022-328X(00)91121-5. [4] Schnepf, Andreas (2008). "Metalloid Cluster Compounds of Germanium: Synthesis – Properties – Subsequent Reactions". European Journal of Inorganic Chemistry 2008: 1007. doi:10.1002/ejic.200700969. [5] Casiot, C (2002). "Optimization of the hyphenation between capillary zone electrophoresis and inductively coupled plasma mass spectrometry for the measurement of As-, Sb-, Se- and Te-species, applicable to soil extracts". Spectrochimica Acta Part B Atomic Spectroscopy 57: 173. doi:10.1016/S0584-8547(01)00365-2. [6] Chasteen, Thomas G.; Bentley, R (2003). "Biomethylation of Selenium and Tellurium: Microorganisms and Plants". Chemical Reviews 103 (1): 1. doi:10.1021/cr010210. PMID 12517179. [7] Polonium-210 Information Sheet (http://www.hps.org/documents/po210_information_sheet.pdf) [8] Rubin, K (1997). "Degassing of metals and metalloids from erupting seamount and mid-ocean ridge volcanoes: Observations and predictions". Geochimica et Cosmochimica Acta 61: 3525. doi:10.1016/S0016-7037(97)00179-8.
Noble gas Group →
18
↓ Period
1
2 He
2
10 Ne
3
18 Ar
4
36 Kr
5
54 Xe
6
86 Rn
7
118 Uuo
Legend
Noble gas Gas Primordial element From decay Synthetic
The noble gases are a group of chemical elements with very similar properties: under standard conditions, they are all odorless, colorless, monatomic gases, with very low chemical reactivity. The six noble gases that occur naturally are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn).
Noble gas For the first six periods of the periodic table, the noble gases are exactly the members of group 18 of the periodic table. However, this no longer holds in the seventh period (due to relativistic effects): the next member of group 18, ununoctium, is probably not a noble gas.[1] Instead, group 14 member ununquadium exhibits noble-gas-like properties.[2] The properties of the noble gases can be well explained by modern theories of atomic structure: their outer shell of valence electrons is considered to be "full", giving them little tendency to participate in chemical reactions, and it has only been possible to prepare a few hundred noble gas compounds. The melting and boiling points for each noble gas are close together, differing by less than 10 °C (18 °F); consequently, they are liquids over only a small temperature range. Neon, argon, krypton, and xenon are obtained from air using the methods of liquefaction of gases and fractional distillation. Helium is typically separated from natural gas, and radon is usually isolated from the radioactive decay of dissolved radium compounds. Noble gases have several important applications in industries such as lighting, welding, and space exploration. A helium-oxygen breathing gas is often used by deep-sea divers at depths of seawater over 180 feet (55 m) to keep the diver from experiencing oxygen toxemia, the lethal effect of high-pressure oxygen, and nitrogen narcosis, the distracting narcotic effect of the nitrogen in air beyond this partial-pressure threshold. After the risks caused by the flammability of hydrogen became apparent, it was replaced with helium in blimps and balloons.
History Noble gas is translated from the German noun Edelgas, first used in 1898 by Hugo Erdmann[3] to indicate their extremely low level of reactivity. The name makes an analogy to the term "noble metals", which also have low reactivity. The noble gases have also been referred to as inert gases, but this label is now deprecated as many noble gas compounds are now known.[4] Rare gases is another term that was used, [5] but this is also inaccurate because argon forms a fairly considerable part (0.94% by volume, 1.3% by mass) of the Earth's atmosphere.[6] Pierre Janssen and Joseph Norman Lockyer discovered a new element on August 18, 1868 while looking at the chromosphere of the Sun, and named it helium after the Greek word for the Sun, ήλιος (ílios or helios).[7] No chemical analysis was Helium was first detected in the Sun due to its characteristic spectral lines. possible at the time, but helium was later found to be a noble gas. Before them, in 1784, the English chemist and physicist Henry Cavendish had discovered that air contains a small proportion of a substance less reactive than nitrogen. [8] A century later, in 1895, Lord Rayleigh discovered that samples of nitrogen from the air were of a different density than nitrogen resulting from chemical reactions. Along with scientist William Ramsay at University College, London, Lord Rayleigh theorized that the nitrogen extracted from air was mixed with another gas, leading to an experiment that successfully isolated a new element, argon, from the Greek word αργός (argós, "inactive").[8] With this discovery, they realized an entire class of gases was missing from the periodic table. During his search for argon, Ramsay also managed to isolate helium for the first time while heating cleveite, a mineral. In 1902, having accepted the evidence for the elements helium and argon, Dmitri Mendeleev included these noble gases as group 0 in his arrangement of the elements, which would later become the periodic table.[9] Ramsay continued to search for these gases using the method of fractional distillation to separate liquid air into several components. In 1898, he discovered the elements krypton, neon, and xenon, and named them after the Greek words κρυπτός (kryptós, "hidden"), νέος (néos, "new"), and ξένος ( xénos, "stranger"), respectively. Radon was first [10] identified in 1898 by Friedrich Ernst Dorn, and was named radium emanation, but was not considered a noble gas until 1904 when its characteristics were found to be similar to those of other noble gases. [11] Rayleigh and Ramsay
118
Noble gas
119
received the 1904 Nobel Prizes in Physics and in Chemistry, respectively, for their discovery of the noble gases; [12] [13] in the words of J. E. Cederblom, then president of the Royal Swedish Academy of Sciences, "the discovery of an entirely new group of elements, of which no single representative had been known with any certainty, is something utterly unique in the history of chemistry, being intrinsically an advance in science of peculiar significance".[13] The discovery of the noble gases aided in the development of a general understanding of atomic structure. In 1895, French chemist Henri Moissan attempted to form a reaction between fluorine, the most electronegative element, and argon, one of the noble gases, but failed. Scientists were unable to prepare compounds of argon until the end of the 20th century, but these attempts helped to develop new theories of atomic structure. Learning from these experiments, Danish physicist Niels Bohr proposed in 1913 that the electrons in atoms are arranged in shells surrounding the nucleus, and that for all noble gases except helium the outermost shell always contains eight electrons.[11] In 1916, Gilbert N. Lewis formulated the octet rule, which concluded an octet of electrons in the outer shell was the most stable arrangement for any atom; this arrangement caused them to be unreactive with other elements since they did not require any more electrons to complete their outer shell.[14] In 1962 Neil Bartlett discovered the first chemical compound of a noble gas, xenon hexafluoroplatinate.[15] Compounds of other noble gases were discovered soon after: in 1962 for radon, radon fluoride, [16] and in 1963 for [17] krypton, krypton difluoride (KrF2). The first stable compound of argon was reported in 2000 when argon fluorohydride (HArF) was formed at a temperature of 40 K (−233.2 °C; −387.7 °F).[18] In December 1998, scientists at the Joint Institute for Nuclear Research working in Dubna, Russia bombarded plutonium (Pu) with calcium (Ca) to produce a single atom of element 114,[19] , ununquadium (Uuq).[20] Preliminary chemistry experiments have indicated this element may be the first superheavy element to show abnormal noble-gas-like properties, even though it is a member of group 14 on the periodic table. [21] In October 2006, scientists from the Joint Institute for Nuclear Research and Lawrence Livermore National Laboratory successfully created synthetically ununoctium (Uuo), the seventh element in group 18,[22] by bombarding californium (Cf) with calcium (Ca).[23]
Chemical properties The noble gases are colorless, odorless, tasteless, and nonflammable under standard conditions. They were once labeled group 0 in the periodic table because it was believed they had a valence of zero, meaning their atoms cannot combine with those of other elements to form compounds. However, it was later discovered some do indeed form compounds, causing this label to fall into disuse.[11] Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:
Neon, like all noble gases, has a full valence shell. Noble gases have eight electrons in the outermost shell, except in the case of helium, which has two.
Noble gas
120
Z Element No. of electrons/shell
2
helium
2
10 neon
2, 8
18 argon
2, 8, 8
36 krypton
2, 8, 18, 8
54 xenon
2, 8, 18, 18, 8
86 radon
2, 8, 18, 32, 18, 8
The noble gases have full valence electron shells. Valence electrons are the outermost electrons of an atom and are normally the only electrons that participate in chemical bonding. Atoms with full valence electron shells are extremely stable and therefore do not tend to form chemical bonds and have little tendency to gain or lose electrons.[24] However, heavier noble gases such as radon are held less firmly together by electromagnetic force than lighter noble gases such as helium, making it easier to remove outer electrons from heavy noble gases. As a result of a full shell, the noble gases can be used in conjunction with the electron configuration notation to form the noble gas notation. To do this, the nearest noble gas that precedes the element in question is written first, and then the electron configuration is continued from that point forward. For example, the electron notation of carbon is 1s²2s²2p², and the noble gas notation is [He]2s²2p². This notation makes it easier to identify elements, and is shorter [25] than writing out the full notation of atomic orbitals.
Compounds The noble gases show extremely low chemical reactivity; consequently, only a few hundred noble gas compounds have been formed. Neutral compounds in which helium and neon are involved in chemical bonds have not been formed (although there is some theoretical evidence for a few helium compounds), while xenon, krypton, and argon have shown only minor reactivity.[26] The reactivity follows the order Ne < He < Ar < Kr < Xe < Rn. In 1933, Linus Pauling predicted that the heavier noble gases could form compounds with fluorine and oxygen. He predicted the existence of krypton hexafluoride (KrF6) and xenon hexafluoride (XeF6), Structure of XeF4, one of the first noble gas speculated XeF8 might exist as an unstable compound, and suggested compounds to be discovered. xenic acid could form perxenate salts. [27] [28] These predictions were shown to be generally accurate, except XeF 8 is now thought to be both thermodynamically and kinetically unstable.[29] Xenon compounds are the most numerous of the noble gas compounds that have been formed.[30] Most of them have the xenon atom in the oxidation state of +2, +4, +6, or +8 bonded to highly electronegative atoms such as fluorine or oxygen, as in xenon difluoride (XeF ), xenon tetrafluoride (XeF ), xenon hexafluoride (XeF ), xenon tetroxide 2 4 6 (XeO4), and sodium perxenate (Na4XeO6). Some of these compounds have found use in chemical synthesis as oxidizing agents; XeF2, in particular, is commercially available and can be used as a fluorinating agent. [31] As of 2007, about five hundred compounds of xenon bonded to other elements have been identified, including organoxenon compounds (those bonded to carbon), and xenon bonded to nitrogen, chlorine, gold, mercury, and xenon itself.[26] [32] Compounds of xenon bound to boron, hydrogen, bromine, iodine, beryllium, sulphur, titanium, copper, and silver have also been observed but only at low temperatures in noble gas matrices, or in supersonic noble gas jets.[26]
Noble gas
121
In theory, radon is more reactive than xenon, and therefore should form chemical bonds more easily than xenon does. However, due to the high radioactivity and short half-life of radon isotopes, only a few fluorides and oxides of radon have been formed in practice.[33] Krypton is less reactive than xenon, but several compounds have been reported with krypton in the oxidation state of +2.[26] Krypton difluoride is the most notable and easily characterized. Compounds in which krypton forms a single bond to nitrogen and oxygen have also been characterized,[34] but are only stable below −60 °C (−76 °F) and −90 °C (−130 °F) respectively).[26] Krypton atoms chemically bound to other nonmetals (hydrogen, chlorine, carbon) as well as some late transition metals (copper, silver, gold) have also been observed, but only either at low temperatures in noble gas matrices, or in supersonic noble gas jets.[26] Similar conditions were used to obtain the first few compounds of argon in 2000, such as argon fluorohydride (HArF), and some bound to the late transition metals copper, silver, and gold. [26] As of 2007, no stable neutral molecules involving covalently bound helium or neon are known.[26] The noble gases —including helium —can form stable molecular ions in the gas phase. The simplest is the helium hydride molecular ion, HeH+, discovered in 1925.[35] Because it is composed of the two most abundant elements in the universe, hydrogen and helium, it is believed to occur naturally in the interstellar medium, although it has not been detected yet.[36] In addition to these ions, there are many known neutral excimers of the noble gases. These are compounds such as ArF and KrF that are stable only when in an excited electronic state; some of them find application in excimer lasers. In addition to the compounds where a noble gas atom is involved in a covalent bond, noble gases also form [37] non-covalent compounds. The clathrates, first described in 1949, consist of a noble gas atom trapped within cavities of crystal lattices of certain organic and inorganic substances. The essential condition for their formation is that the guest (noble gas) atoms must be of appropriate size to fit in the cavities of the host crystal lattice. For instance, argon, krypton, and xenon form clathrates with hydroquinone, but helium and neon do not because they are too small or insufficiently polarizable to be retained.[38] Neon, argon, krypton, and xenon also form clathrate hydrates, where the noble gas is trapped in ice.[39] Noble gases can form endohedral fullerene compounds, in which the noble gas atom is trapped inside a fullerene molecule. In 1993, it was discovered that when C , a spherical molecule consisting of 60 60 carbon atoms, is exposed to noble gases at high pressure, complexes such as He@C60 can be formed (the @ notation indicates He is contained inside C but not covalently bound to it).[40] As of 2008, 60 endohedral complexes with helium, neon, argon, krypton, and xenon have been obtained.[41] These compounds have found use in the study of the structure and reactivity of fullerenes by means of the nuclear magnetic resonance of the noble gas atom.[42] An endohedral fullerene compound containing a noble gas atom
Noble gas Noble gas compounds such as xenon difluoride (XeF2) are considered to be hypervalent because they violate the octet rule. Bonding in such compounds can be explained using a 3-center-4-electron bond model.[43] [44] This model, first proposed in 1951, considers bonding of three collinear atoms. For example, bonding in XeF2 is described by a set of three molecular orbitals (MOs) derived from p-orbitals on Bonding in XeF2 according to the 3-center-4-electron bond model each atom. Bonding results from the combination of a filled p-orbital from Xe with one half-filled p-orbital from each F atom, resulting in a filled bonding orbital, a filled non-bonding orbital, and an empty antibonding orbital. The highest occupied molecular orbital is localized on the two terminal atoms. This represents a localization of charge which is facilitated by the high electronegativity of fluorine.[45] The chemistry of heavier noble gases, krypton and xenon, are well established. The chemistry of the lighter ones, argon and helium, is still at an early stage, while a neon compound is still yet to be identified.
Occurrence and production The abundances of the noble gases in the universe decrease as their atomic numbers increase. Helium is the most common element in the universe after hydrogen, with a mass fraction of about 24%. Most of the helium in the universe was formed during Big Bang nucleosynthesis, but the amount of helium is steadily increasing due to the fusion of hydrogen in stellar nucleosynthesis.[46] [47] Abundances on Earth follow different trends; for example, helium is only the third most abundant noble gas in the atmosphere. The reason is that there is no primordial helium [48] in the atmosphere; due to the small mass of the atom, helium cannot be retained by the Earth's gravitational field. Helium on Earth comes from the alpha decay of heavy elements such as uranium and thorium found in the Earth's crust, and tends to accumulate in natural gas deposits.[48] The abundance of argon, on the other hand, is increased as a result of the beta decay of potassium-40, also found in the Earth's crust, to form argon-40, which is the most abundant isotope of argon on Earth despite being relatively rare in the Solar System. This process is the base for the potassium-argon dating method.[49] Xenon has an unexpectedly low abundance in the atmosphere, in what has been called the missing xenon problem ; one theory is that the missing xenon may be trapped in minerals inside the Earth's crust.[50] Radon is formed in the lithosphere as from the alpha decay of radium. It can seep into buildings through cracks in their foundation and accumulate in areas that are not well ventilated. Due to its high radioactivity, radon presents a significant health hazard; it is implicated in an estimated 21,000 lung cancer deaths per year in the United States alone.[51]
122
Noble gas
123
Abundance
Helium
Neon
Argon
[52] Solar System (for each atom of silicon)
2343
2.148
0.1025
Earth's atmosphere (volume fraction in [53] ppm)
5.20
18.20
9340.00
[55] Igneous rock (mass fraction in ppm)
Krypton
5.515 × 10−5 5.391 × 10 −6
–
0.09
[54] (0.06 – 18) × 10−19
–
–
1.7 × 10−10
2004 price 3 [56] (USD/m )
Helium (industrial grade)
4.20 – 4.90
Helium (laboratory grade)
22.30 – 44.90
Argon
2.70 – 8.50
Neon
60 – 120
Krypton
400 – 500
Xenon
Radon
1.10
3 × 10−3 7 × 10−5 4 × 10−2
Gas
Xenon
4000 – 5000
Neon, argon, krypton, and xenon are obtained from air using the methods of liquefaction of gases, to convert elements to a liquid state, and fractional distillation, to separate mixtures into component parts. Helium is typically produced by separating it from natural gas, and radon is isolated from the radioactive decay of radium compounds.[11] The prices of the noble gases are influenced by their natural abundance, with argon being the cheapest and xenon the most expensive. As an example, the table to the right lists the 2004 prices in the United States for laboratory quantities of each gas.
Applications Noble gases have very low boiling and melting points, which makes them useful as cryogenic refrigerants.[57] In particular, liquid helium, which boils at 4.2 K (−268.95 °C; −452.11 °F), is used for superconducting magnets, such as those needed in nuclear magnetic resonance imaging and nuclear magnetic resonance.[58] Liquid neon, although it does not reach temperatures as low as liquid helium, also finds use in cryogenics because it has over 40 times more refrigerating capacity than liquid helium and over three times more than liquid hydrogen.[54] Liquid helium is used to cool the superconducting magnets in modern MRI scanners.
Helium is used as a component of breathing gases to replace nitrogen, due its low solubility in fluids, especially in lipids. Gases are absorbed by the blood and body tissues when under pressure like in scuba diving, which causes an anesthetic effect known as [59] nitrogen narcosis. Due to its reduced solubility, little helium is taken into cell membranes, and when helium is used to replace part of the breathing mixtures, such as in trimix or heliox, a decrease in the narcotic effect of the gas at depth is obtained.[60] Helium's reduced solubility offers further advantages for the condition known as decompression sickness, or the bends.[11] [61] The reduced amount of dissolved gas in the body means that fewer gas bubbles form during the decrease in pressure of the ascent. Another noble gas, argon, is considered the best option for use as a drysuit inflation gas for scuba diving.[62]
Noble gas
124
Since the Hindenburg disaster in 1937,[63] helium has replaced hydrogen as a lifting gas in blimps and balloons due to its lightness and incombustibility, despite an 8.6%[64] decrease in buoyancy.[11] In many applications, the noble gases are used to provide an inert atmosphere. Argon is used in the synthesis of air-sensitive compounds The Spirit of Goodyear , one of the iconic that are sensitive to nitrogen. Solid argon is also used for the study of Goodyear Blimps very unstable compounds, such as reactive intermediates, by trapping them in an inert matrix at very low temperatures.[65] Helium is used as the carrier medium in gas chromatography, as a filler gas for thermometers, and in devices for measuring radiation, such as the Geiger counter and the bubble chamber.[56] Helium and argon are both commonly used to shield welding arcs and the surrounding base metal from the atmosphere during welding and cutting, as well as in other metallurgical processes and in the production of silicon for the semiconductor industry.[54]
15,000-watt xenon short-arc lamp used in IMAX projectors
Noble gases are commonly used in lighting because of their lack of chemical reactivity. Argon, mixed with nitrogen, is used as a filler gas for incandescent light bulbs.[54] Krypton is used in high-performance light bulbs, which have higher color temperatures and greater efficiency, because it reduces the rate of evaporation of the filament more than argon; halogen lamps, in particular, use krypton mixed with small amounts of compounds of iodine or bromine.[54] The noble gases glow in distinctive colors when used inside gas-discharge lamps, such as neon lights, which produce an orange-red color. Xenon is commonly used in xenon arc lamps which, due to their nearly continuous spectrum that resembles daylight, find application in film projectors and as automobile headlamps.[54]
The noble gases are used in excimer lasers, which are based on short-lived electronically excited molecules known as excimers. The excimers used for lasers may be noble gas dimers such as Ar 2, Kr2 or Xe2, or more commonly, the noble gas is combined with a halogen in excimers such as ArF, KrF, XeF, or XeCl. These lasers produce ultraviolet light which, due to its short wavelength (193 nm for ArF and 248 nm for KrF), allows for high-precision imaging. Excimer lasers have many industrial, medical, and scientific applications. They are used for microlithography and microfabrication, which are essential for integrated circuit manufacture, and for laser surgery, including laser angioplasty and eye surgery.[66] Some noble gases have direct application in medicine. Helium is sometimes used to improve the ease of breathing of asthma sufferers.[54] Xenon is used as an anesthetic because of its high solubility in lipids, which makes it more potent than the usual nitrous oxide, and because it is readily eliminated from the body, resulting in faster recovery.[67] Xenon finds application in medical imaging of the lungs through hyperpolarized MRI.[68] Radon, which is highly radioactive and is only available in minute amounts, is used in radiotherapy.[11]
Discharge color
Noble gas
125
Colors and spectra (bottom row) of electric discharge in pure noble gases
Helium
Neon
Argon (with some Hg in the "Ar" image)
Krypton
Xenon
The color of gas discharge emission depends on several factors, including the following:[69] • discharge parameters (local value of current density and electric field, temperature, etc. – note the color variation along the discharge in the top row); • gas purity (even small fraction of certain gases can affect color); • color balance and saturation level of the image recording medium; • material of the discarge tube envelope – note suppression of the UV and blue components in the bottom-row tubes made of thick household glass.
See also • • • •
Noble gas (data page), for extended tables of physical properties. Noble metal, for metals that are resistant to corrosion or oxidation. Inert gas, for any gas that is not reactive under normal circumstances. Industrial gas
References • Bennett, Peter B.; Elliott, David H. (1998). The Physiology and Medicine of Diving. SPCK Publishing. ISBN 0702024104. • Bobrow Test Preparation Services (2007-12-05). CliffsAP Chemistry. CliffsNotes. ISBN 047013500X. • Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. • Harding, Charlie J.; Janes, Rob (2002). Elements of the P Block . Royal Society of Chemistry. ISBN 0854046909. • Holloway, John H. (1968). Noble-Gas Chemistry . London: Methuen Publishing. ISBN 0412211009. • Mendeleev, D. (1902 – 1903) (in Russian). Osnovy Khimii (The Principles of Chemistry) [70] (7th ed.). [71] • Ozima, Minoru; Podosek, Frank A. (2002). Noble Gas Geochemistry . Cambridge University Press. ISBN 0521803667. • Weinhold, F.; Landis, C. (2005). Valency and bonding. Cambridge University Press. ISBN 0521831288.
Noble gas
References [1] Clinton S. Nash (2005). "Atomic and Molecular Properties of Elements 112, 114, and 118". J. Phys. Chem. A 109 (15): 3493 – 3500. doi:10.1021/jp050736o. PMID 16833687. [2] "Flerov laboratory of nuclear reactions" (http://www1.jinr.ru/Reports/2008/english/06_flnr_e.pdf). JINR. . Retrieved 2009-08-08. [3] Renouf, Edward (1901). "Noble gases". Science 13: 268 – 270. doi:10.1126/science.13.320.268. [4] Ozima 2002, p. 30 [5] Ozima 2002, p. 4 [6] "argon" (http://www.britannica.com/eb/article-9009382/argon). Encyclopædia Britannica. 2008. . [7] Oxford English Dictionary (1989), s.v. "helium". Retrieved December 16, 2006, from Oxford English Dictionary Online. Also, from quotation there: Thomson, W. (1872). Rep. Brit. Assoc. xcix: "Frankland and Lockyer find the yellow prominences to give a very decided bright line not far from D, but hitherto not identified with any terrestrial flame. It seems to indicate a new substance, which they propose to call Helium." [8] Ozima 2002, p. 1 [9] Mendeleev 1903, p. 497 [10] Partington, J. R. (1957). "Discovery of Radon". Nature 179 (4566): 912. doi:10.1038/179912a0. [11] "Noble Gas" (http://www.britannica.com/eb/article-9110613/noble-gas). Encyclopædia Britannica. 2008. . [12] Cederblom, J. E. (1904). "The Nobel Prize in Physics 1904 Presentation Speech" (http://nobelprize.org/nobel_prizes/physics/laureates/ 1904/press.html). . [13] Cederblom, J. E. (1904). "The Nobel Prize in Chemistry 1904 Presentation Speech" (http://nobelprize.org/nobel_prizes/chemistry/ laureates/1904/press.html). . [14] Gillespie, R. J.; Robinson, E. A. (2007). "Gilbert N. Lewis and the chemical bond: the electron pair and the octet rule from 1916 to the present day". J Comput Chem 28 (1): 87 – 97. doi:10.1002/jcc.20545. PMID 17109437. [15] Bartlett, N. (1962). "Xenon hexafluoroplatinate Xe+[PtF6] – ". Proceedings of the Chemical Society (6): 218. doi:10.1039/PS9620000197. [16] Fields, Paul R.; Stein, Lawrence; Zirin, Moshe H. (1962). "Radon Fluoride". Journal of the American Chemical Society 84 (21): 4164 – 4165. doi:10.1021/ja00880a048. [17] Grosse, A. V.; Kirschenbaum, A. D.; Streng, A. G.; Streng, L. V. (1963). "Krypton Tetrafluoride: Preparation and Some Properties".Science 139 (3559): 1047 – 1048. doi:10.1126/science.139.3559.1047. PMID 17812982. [18] Khriachtchev, Leonid; Pettersson, Mika; Runeberg, Nino; Lundell, Jan; Räsänen, Markku (2000). "A stable argon compound". Nature 406 (406): 874 – 876. doi:10.1038/35022551. PMID 10972285. [19] Oganessian, Yu. Ts.; Utyonkov, V.; Lobanov, Yu.; Abdullin, F.; Polyakov, A.; Shirokovsky, I.; Tsyganov, Yu.; Gulbekian, G. et al. (1999). "Synthesis of Superheavy Nuclei in the 48Ca + 244Pu Reaction". Physical Review Letters (American Physical Society) 83: 3154. doi:10.1103/PhysRevLett.83.3154. [20] Woods, Michael (2003-05-06). "Chemical element No. 110 finally gets a name —darmstadtium" (http://www.post-gazette.com/ healthscience/20030506element0506p4.asp). Pittsburgh Post-Gazette . . Retrieved 2008-06-26. [21] "Gas Phase Chemistry of Superheavy Elements" (http://lch.web.psi.ch/files/lectures/TexasA&M/TexasA&M.pdf) (PDF). Texas A&M University. . Retrieved 2008-05-31. [22] Wilson, Elaine (2005). "Making Meaning in Chemistry Lessons" (http://ejlts.ucdavis.edu/sites/ejlts.ucdavis.edu/files/articles/Wilson. pdf). Electronic Journal of Literacy through Science 4 (2). . Retrieved 2009-08-01. [23] Oganessian, Yu. Ts.; Utyonkov, V.; Lobanov, Yu.; Abdullin, F.; Polyakov, A.; Sagaidak, R.; Shirokovsky, I.; Tsyganov, Yu. et al. (2006). "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245Cm + 48Ca fusion reactions". Physical Review C 74 (4): 44602. doi:10.1103/PhysRevC.74.044602. [24] Ozima 2002, p. 35 [25] CliffsNotes 2007, p. 15 [26] Grochala, Wojciech (2007). "Atypical compounds of gases, which have been called noble". Chemical Society Reviews 36 (36): 1632 – 1655. doi:10.1039/b702109g. PMID 17721587. [27] Pauling, Linus (1933). "The Formulas of Antimonic Acid and the Antimonates". Journal of the American Chemical Society 55 (5): 1895 – 1900. doi:10.1021/ja01332a016. [28] Holloway 1968 [29] Seppelt, Konrad (1979). "Recent developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research 12: 211 – 216. doi:10.1021/ar50138a004. [30] Moody, G. J. (1974). "A Decade of Xenon Chemistry" (http://www.eric.ed. gov/ERICWebPortal/custom/portlets/recordDetails/ detailmini. jsp?_nfpb=true&_&ERICExtSearch_SearchValue_0=EJ111480&ERICExtSearch_SearchType_0=no&accno=EJ111480). Journal of Chemical Education 51 (10): 628 – 630. doi:10.1021/ed051p628. . Retrieved 2007-10-16. [31] Zupan, Marko; Iskra, Jernej; Stavber, Stojan (1998). "Fluorination with XeF2. 44. Effect of Geometry and Heteroatom on the Regioselectivity of Fluorine Introduction into an Aromatic Ring". J. Org. Chem 63 (3): 878 – 880. doi:10.1021/jo971496e. PMID 11672087. [32] Harding 2002, pp. 90 – 99 [33] .Avrorin, V. V.; Krasikova, R. N.; Nefedov, V. D.; Toropova, M. A. (1982). "The Chemistry of Radon". Russian Chemical Review 51 (1): 12 – 20. doi:10.1070/RC1982v051n01ABEH002787.
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"Stable compounds of helium and neon. He@C60 and Ne@C60". Science 259 (5100): 1428 – 1430. doi:10.1126/science.259.5100.1428. PMID 17801275. [41] Saunders, Martin; Jimenez-Vazquez, Hugo A.; Cross, R. James; Mroczkowski, Stanley; Gross, Michael L.; Giblin, Daryl E.; Poreda, Robert J. (1994). "Incorporation of helium, neon, argon, krypton, and xenon into fullerenes using high pressure". J. Am. Chem. Soc. 116 (5): 2193 – 2194. doi:10.1021/ja00084a089. [42] Frunzi, Michael; Cross, R. James; Saunders, Martin (2007). "Effect of Xenon on Fullerene Reactions". Journal of the American Chemical Society 129 (43): 13343. doi:10.1021/ja075568n. PMID 17924634. [43] Greenwood 1997, p. 897 [44] Weinhold 2005, pp. 275 – 306 [45] Pimentel, G. C. (1951). "The Bonding of Trihalide and Bifluoride Ions by the Molecular Orbital Method". The Journal of Chemical Physics 19 (4): 446 – 448. doi:10.1063/1.1748245. [46] Weiss, Achim. "Elements of the past: Big Bang Nucleosynthesis and observation" (http://www.einstein-online.info/en/spotlights/ BBN_obs/index.html). Max Planck Institute for Gravitational Physics. . Retrieved 2008-06-23. [47] Coc, A.; et al. (2004). "Updated Big Bang Nucleosynthesis confronted to WMAP observations and to the Abundance of Light Elements". Astrophysical Journal 600: 544. doi:10.1086/380121. [48] Morrison, P.; Pine, J. (1955). "Radiogenic Origin of the Helium Isotopes in Rock". Annals of the New York Academy of Sciences 62 (3): 71 – 92. doi:10.1111/j.1749-6632.1955.tb35366.x. 40 39 [49] Scherer, Alexandra (2007-01-16). " Ar/ Ar dating and errors" (http://www.geoberg.de/text/geology/07011601.php). Technische Universität Bergakademie Freiberg. . Retrieved 2008-06-26. [50] Sanloup, Chrystèle; et al. (2005). "Retention of Xenon in Quartz and Earth's Missing Xenon". Science 310 (5751): 1174 – 1177. doi:10.1126/science.1119070. PMID 16293758. [51] "A Citizen's Guide to Radon" (http://www.epa.gov/radon/pubs/citguide.html). U.S. Environmental Protection Agency. 2007-11-26. . Retrieved 2008-06-26. [52] Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal 591: 1220 – 1247. doi:10.1086/375492. [53] "The Atmosphere" (http://www.srh.noaa.gov/jetstream//atmos/atmos_intro.htm). National Weather Service. . Retrieved 2008-06-01. [54] Häussinger, Peter; Glatthaar, Reinhard; Rhode, Wilhelm; Kick, Helmut; Benkmann, Christian; Weber, Josef; Wunschel, Hans-Jörg; Stenke, Viktor; Leicht, Edith; Stenger, Hermann ( 2002). "Noble gases". Ullmann's Encyclopedia of Industrial Chemistry . Wiley. doi:10.1002/14356007.a17_485. [55] Greenwood 1997, p. 891 [56] Hwang, Shuen-Chen; Lein, Robert D.; Morgan, Daniel A. (2005). "Noble Gases". Kirk Othmer Encyclopedia of Chemical Technology . Wiley. pp. 343 – 383. doi:10.1002/0471238961.0701190508230114.a01. [57] "Neon". Encarta. 2008. [58] Zhang, C. J.; Zhou, X. T.; Yang, L. (1992). "Demountable coaxial gas-cooled current leads for MRI superconducting magnets". Magnetics, IEEE Transactions on (IEEE) 28 (1): 957 – 959. doi:10.1109/20.120038. —a critical review" (http://archive. [59] Fowler, B; Ackles, K. N.; Porlier, G. (1985). "Effects of inert gas narcosis on behavior rubicon-foundation.org/3019). Undersea Biomed. Res. 12 (4): 369 – 402. ISSN 0093-5387. OCLC 2068005. PMID 4082343. . Retrieved 2008-04-08. [60] Bennett 1998, p. 176 [61] Vann, R. D. (ed) (1989). "The Physiological Basis of Decompression" (http://archive.rubicon-foundation.org/6853). 38th Undersea and Hyperbaric Medical Society Workshop 75(Phys)6-1-89: 437. . Retrieved 2008-05-31. [62] Maiken, Eric (2004-08-01). "Why Argon?" (http://www.decompression.org/maiken/Why_Argon.htm). Decompression. . Retrieved 2008-06-26. [63] "Disaster Ascribed to Gas by Experts". The New York Times : p. 1. 1937-05-07. [64] Freudenrich, Craig (2008). "How Blimps Work" (http://science.howstuffworks.com/blimp2.htm). HowStuffWorks. . Retrieved 2008-07-03. [65] Dunkin, I. R. (1980). "The matrix isolation technique and its application to organic chemistry". Chem. Soc. Rev. 9: 1 – 23. doi:10.1039/CS9800900001. [66] Basting, Dirk; Marowsky, Gerd (2005). Excimer Laser Technology . Springer. ISBN 3540200568.
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[67] Sanders, Robert D.; Ma, Daqing; Maze, Mervyn (2005). "Xenon: elemental anaesthesia in clinical practice". British Medical Bulletin 71 (1): 115 – 135. doi:10.1093/bmb/ldh034. PMID 15728132. [68] Albert, M. S.; Balamore, D. (1998). "Development of hyperpolarized noble gas MRI". Nuclear Instruments and Methods in Physics Research A 402: 441 – 453. doi:10.1016/S0168-9002(97)00888-7. PMID 11543065. [69] Ray, Sidney F. (1999). Scientific photography and applied imaging (http://books.google.com/?id=AEFPNfghI3QC&pg=PA383). Focal Press. pp. 383 – 384. ISBN 0240513231. . [70] http://www.archive.org/details/principlesofchem00menduoft [71] http://books.google.com/?id=CBM2LJDvRtgC
Noble metal Noble metals are metals that are resistant to corrosion and oxidation in
moist air, unlike most base metals. They tend to be precious, often due to their rarity in the Earth's crust. The noble metals are considered to be (in order of increasing atomic number)[1] ruthenium, rhodium, palladium, silver, osmium, iridium, platinum, gold . Other sources include mercury[2] [3] [4] or even rhenium[5] as a noble metal. On the other hand, neither titanium nor niobium nor tantalum are called noble metals despite the fact that they are very resistant to corrosion. Noble metals should not be confused with precious metals (although many noble metals are precious).
The noble metals including mercury and r henium together with the non-noble metal copper ordered according their position in the periodic table of the elements
Introduction Palladium, osmium, platinum, gold and mercury can be dissolved in aqua regia, a highly concentrated mixture of hydrochloric acid and nitric acid, but iridium and silver cannot. (Silver can dissolve in nitric acid though.) Ruthenium can be dissolved in aqua regia only when in the presence of oxygen, while rhodium must be in a fine pulverized form. Niobium and tantalum are resistant to acids, including aqua regia. [6] This term can also be used in a relative sense, considering "noble" as an adjective for the word "metal". A "galvanic series" is a hierarchy of metals (or other electrically conductive materials, including composites and semimetals) that runs from noble to active, and allows designers to see at a glance how materials will interact in the environment used to generate the series. In this sense of the word, graphite is more noble than silver and the relative nobility of many materials is highly dependent upon context, as for aluminium and stainless steel in conditions of varying pH.[7] In physics, the definition of a noble metal is even more strict. It is required that the d-bands of the electronic structure are filled. Taking this into account, only copper, silver and gold are noble metals, as all d-like band are filled and don't cross the Fermi level. [8] For platinum two d-bands cross the Fermi level, changing its chemical behaviour; it is used as a catalyst. The different reactivity can easily be seen while preparing clean metal surfaces in ultra high vacuum; surfaces of "physical defined" noble metals (e.g., gold) are easy to clean and stay clean for a long time, while those of platinum or palladium, for example, are covered by carbon monoxide very quickly.[9]
Noble metal
129
Electrochemistry Metallic elements, including several non-noble metals, sorted by their chemical "nobility" (noble metals bolded): [10] element
group
reaction
potential
Gold
Ib/6
Platinum
VIIIb/6 Pt → Pt2+ + 2 e−
1.18 V
Iridium
VIIIb/6 Ir → Ir3+ + 3 e−
1.156 V
Palladium
VIIIb/5 Pd → Pd2+ + 2 e−
0.987 V
Osmium
VIIIb/6 Os + 4 H O → OsO + 8 H+ + 8 e− 2 4
0.838 V
Silver
Ib/5
Ag → Ag + e− +
0.7996 V
Mercury
IIb/6
2 Hg → Hg22+ + 2 e−
0.7973 V
Polonium
VIa/6
Po → Po2+ + 2 e−
0.65 V [11]
Rhodium
VIIIb/5 Rh → Rh2+ + 2 e−
Au → Au3+ + 3 e−
1.498 V
0.600 V
Ruthenium VIIIb/5
Ru → Ru2+ + 2 e−
0.455 V
Copper
Ib/4
Cu → Cu2+ + 2 e−
0.337 V
Bismuth
Va/6
Bi → Bi + 3 e−
0.308 V
3+
Technetium VIIb/5 Tc + 2 H O → TcO + 4 H+ + 4 e− 2 2
0.272 V
Rhenium
VIIb/6
Re + 2 H2O → ReO2 + 4 H+ + 4 e−
0.259 V
Antimony
Va/5
2 Sb + 3 H2O → Sb2O3 + 6 H+ + 6 e− 0.152 V
The column group denotes its position in the periodic table, hence electronic configuration. The simplified reactions, listed in the next column, can also be read in detail from the Pourbaix diagrams of the considered element in water. Finally the column potential indicates the electric potential of the element measured against a H-electrode in aqueous, pH 7 solution. All missing elements in this table are either not metals or have a negative standard potential. Antimony and polonium are considered metalloids and thus can not be noble metals. Also chemists and metallurgists − consider copper and bismuth not noble metals because they easily oxidize due to the reaction O2 + 2 H2O + 4 e 4 − OH (aq) +0.40 V which is possible in moist air. Silver and copper film over and oxidize easily and readily, thus the copper sheets with a patina of oxidation used in architectural designs and the resultant market for a myriad of silver polishing compounds. The film over of Silver is due to its high sensibility to hydrogen sulfide. Chemically patina is caused by an attack of oxygen in wet air and by CO afterward.[6] On the other hand, rhenium coated mirrors are said to be very durable,[6] despite the fact that 2 rhenium and technetium are said to tarnish slowly in moist atmosphere.[12]
Noble metal
See also • Base metal • Precious metal
References [1] [2] [3] [4]
A. Holleman, N. Wiberg, "Lehrbuch der Anorganischen Chemie", de Gruyter, 1985, 33. edition, p. 1486 Die Adresse für Ausbildung, Studium und Beruf (http://www.uni-protokolle.de/Lexikon/Edelmetall.html) "Dictionary of Mining, Mineral, and Related Terms", Compiled by the American Geological Institute, 2nd edition, 1997 Scoullos, M.J., Vonkeman, G.H., Thornton, I., Makuch, Z., "Mercury - Cadmium - Lead: Handbook for Sustainable Heavy Metals Policy and Regulation",Series: Environment & Policy, Vol. 31, Springer-Verlag, 2002 [5] The New Encyclopedia Britannica, 15th edition, Vol. VII, 1976 [6] A. Holleman, N. Wiberg, "Inorganic Chemistry", Academic Press, 2001 [7] Everett Collier, "The Boatowner’s Guide to Corrosion", International Marine Publishing, 2001, p. 21 [8] Hüger, E.; Osuch, K. (2005). "Making a noble metal of Pd". EPL (Europhysics Letters) 71: 276. doi:10.1209/epl/i2005-10075-5. [9] S. Fuchs, T.Hahn, H.G. Lintz, "The oxidation of carbon monoxide by oxygen over platinum, palladium and rhodium catalysts from 10−10 to 1 bar", Chemical engineering and processing, 1994, V 33(5), pp. 363-369 (http://cat.inist.fr/?aModele=afficheN&cpsidt=3322977) [10] D. R. Lidle editor, "CRC Handbook of Chemistry and Physics", 86th edition, 2005 [11] A. J. Bard, "Encyclopedia of the Electrochemistry of the Elements", Vol. IV, Marcel Dekker Inc., 1975 [12] R. D. Peack, "The Chemistry of Technetium and Rhenium", Elsevier, 1966
Notes • R. R. Brooks, "Noble metals and biological systems: their role in Medicine, Mineral Exploration, and the Environment", CRC Press, 1992
External links • noble metal - chemistry (http://www.britannica.com/EBchecked/topic/416979/noble-metal) Encyclopædia Britannica, online edition • To see which bands cross the Fermi level, the Fermi surfaces of almost all the metals can be found at the Fermi Surface Database (http://www.phys.ufl.edu/fermisurface/) • The following article might also clarify the correlation between band structure and the term noble metal: Hüger, E.; Osuch, K. (2005). "Making a noble metal of Pd". EPL (Europhysics Letters) 71: 276. doi:10.1209/epl/i2005-10075-5.
130
Nonmetal
Nonmetal Nonmetal, or non-metal, is a term used in chemistry when classifying the chemical elements. On the basis of their
general physical and chemical properties, every element in the periodic table can be termed either a metal or a nonmetal. (A few elements with intermediate properties are referred to as metalloids). The elements generally regarded as nonmetals are: • • • • • •
hydrogen (H) In Group 14: carbon (C) In Group 15 (the pnictogens): nitrogen (N), phosphorus (P) Several elements in Group 16, the chalcogens: oxygen (O), sulfur (S), selenium (Se) All elements in Group 17 - the halogens All elements in Group 18 - the noble gases
There is no rigorous definition for the term "nonmetal" - it covers a general spectrum of behaviour. Common properties considered characteristic of a nonmetal include: • • • • • •
poor conductors of heat and electricity when compared to metals they form acidic oxides (whereas metals generally form basic oxides) in solid form, they are dull and brittle, rather than metals which are lustrous, ductile or malleable usually have lower densities than metals they have significantly lower melting points and boiling points than metals non-metals have high electronegativity
They also have a negative valence, compared to the positive valence of metals. Only eighteen elements in the periodic table are generally considered nonmetals, compared to over eighty metals, but nonmetals make up most of the crust, atmosphere and oceans of the earth. Bulk tissues of living organisms are composed almost entirely of nonmetals. Most nonmetals are monatomic noble gases or form diatomic molecules in their elemental state, unlike metals which (in their elemental state) do not form molecules at all.
Metallisation at huge pressures Nevertheless, even these 18 elements tend to become metallic at large enough pressures (see nearby periodic table at ~300 GPa).
131
Platinum group
132
Platinum group H Li
He Be
B
C
N
O
F
Na Mg
Al
Si
P
S
Cl Ar
As
Se Br Kr
K
Ca Sc
Ti
Rb
Sr
Zr Nb Mo Tc
Ru Rh
Pd Ag Cd In
Sn
Sb
Te
Hf Ta
Pb
Bi
Po At Rn
Y
Cs Ba La Fr
*
V
Cr Mn Fe
Co
Ni
Cu Zn Ga Ge
Ne
W
Re
Os
Ir
Pt
Au Hg Tl
Ra Ac ** Rf Db Sg
Bh
Hs
Mt
Ds
Rg Cn
*
Ce Pr Nd Pm Sm Eu
** Th Pa
U
I
Xe
Gd Tb Dy Ho Er Tm Yb Lu
Np Pu Am Cm Bk Cf
Es Fm Md No Lr
Platinum group metals
The platinum group metals (abbreviated as the PGMs; alternatively, the platinoids, platidises, platinum group or platinum metals) sometimes collectively refers to six metallic elements clustered together in the periodic table. These elements are all transition metals, lying in the d-block (groups 8, 9, and 10, periods 5 and 6). The six platinum group metals are ruthenium, rhodium, palladium, osmium, iridium, and platinum. They have similar physical and chemical properties, and tend to occur together in the same mineral deposits.[1]
History Naturally occurring platinum and platinum-rich alloys have been known by pre-Columbian Americans for a long time.[2] Though the metal was used by pre-Columbian peoples, the first European reference to platinum appears in 1557 in the writings of the Italian humanist Julius Caesar Scaliger (1484 – 1558) as a description of a mysterious metal found in Central American mines between Darién (Panama) and Mexico ("up until now impossible to melt by any of the Spanish arts").[2] The Spaniards named the metal platina ("little silver") when they first encountered it in Colombia. They regarded platinum as an unwanted impurity in the silver they were mining.[2] [3]
Properties The platinum metals have outstanding catalytic properties. They are highly resistant to wear and tarnish, making platinum, in particular, well suited for fine jewelry. Other distinctive properties include resistance to chemical attack, excellent high-temperature characteristics, and stable electrical properties. All these properties have been exploited for industrial applications.[4]
Sources Platinum Sperrylite (platinum arsenide, PtAs2) ore is a major source of this metal. A naturally occurring platinum-iridium alloy, platiniridium, is found in the mineral cooperite (platinum sulfide, PtS). Platinum in a native state, often accompanied by small amounts of other platinum metals, is found in alluvial and placer deposits in Colombia, Ontario, the Ural Mountains, and in certain western American states. Platinum is also produced commercially as a by-product of nickel ore processing. The huge quantities of nickel ore processed
Platinum group makes up for the fact that platinum makes up only two parts per million of the ore. South Africa, with vast platinum ore deposits in the Merensky Reef of the Bushveld complex, is the world's largest producer of platinum, followed by Russia.[5] [6] Platinum and palladium are also mined commercially from the Stillwater igneous complex in Montana, USA. Osmium Iridiosmium is a naturally occurring alloy of iridium and osmium found in platinum-bearing river sands in the Ural Mountains and in North and South America. Trace amounts of osmium also exist in nickel-bearing ores found in the Sudbury, Ontario region along with other platinum group metals. Even though the quantity of platinum metals found in these ores is small, the large volume of nickel ores processed makes commercial recovery possible.[6] [7] Iridium Metallic iridium is found with platinum and other platinum group metals in alluvial deposits. Naturally occurring iridium alloys include osmiridium and iridiosmium, both of which are mixtures of iridium and osmium. It is recovered commercially as a by-product from nickel mining and processing.[6] Ruthenium Ruthenium is generally found in ores with the other platinum group metals in the Ural Mountains and in North and South America. Small but commercially important quantities are also found in pentlandite extracted from Sudbury, Ontario and in pyroxenite deposits in South Africa.[6] Rhodium The industrial extraction of rhodium is complex as the metal occurs in ores mixed with other metals such as palladium, silver, platinum, and gold. It is found in platinum ores and obtained free as a white inert metal which is very difficult to fuse. Principal sources of this element are located in river sands of the Ural Mountains, in North and South America and also in the copper-nickel sulfide mining area of the Sudbury Basin region. Although the quantity at Sudbury is very small, the large amount of nickel ore processed makes rhodium recovery cost effective. However, the annual world production in 2003 of this element is only 7 or 8 tons and there are very few rhodium minerals.[8] Palladium Palladium is found as a free metal and alloyed with platinum and gold with platinum group metals in placer deposits of the Ural Mountains of Eurasia, Australia, Ethiopia, South and North America. However it is commercially produced from nickel-copper deposits found in South Africa and Ontario, Canada. The huge volume of nickel-copper ore processed makes this extraction profitable in spite of its low concentration in these ores.[8]
Production The production of pure platinum group metals normally starts from residues of the production of other metals with a mixture of several of those metals. One typical starting product is the anode residue of gold or nickel production. The differences in chemical reactivity and solubility of several compounds of the metals under extraction are used to separate them.[4] A first step is to dissolve all the metals in aqua regia forming their respective nitrates. If silver is still present, this is then separated by forming insoluble silver chloride. Rhodium sulfate is separated after the salts have been melted together with sodium hydrogensulfate and leached with water. The residue is then melted together with sodium peroxide, which dissolves all the metals and leaves the iridium. The two remaining metals, ruthenium and osmium, form ruthenium and osmium tetroxides after chlorine has been added to solution. The osmium tetroxide is then dissolved in alcoholic sodium hydroxide and separated from the ruthenium tetroxides. All of these metals' final chemical compounds can ultimately be reduced to the elemental metal using hydrogen.[4]
133
Platinum group
Production in nuclear reactors Significant quantities of platinum group metals – Ruthenium, Rhodium and Palladium are formed as fission products in nuclear reactors.[9] With escalating prices and increasing global demand, reactor produced noble metals are emerging as an alternative source. Various reports are available on the possibility of recovering fission noble metals from spent nuclear fuel.[10] [11] [12] Recently there is an upsurge in the recovery of valuable fission products which ref lects in the form of articles in leading scientific journals. Palladium has been of special interest due to its less complex behavior when compared to rhodium and ruthenium. The special interest in palladium may be also due to its widespread application in chemical catalysis and the electronics industry. Several research groups are exploring the possibility of recovering palladium by various methods like direct electrolysis of high-level liquid waste,[13] [14] using room temperature ionic liquids [15] (RTILs) as electrolyte for nuclear fuel dissolution and recovery, solvent extraction, ion exchange, etc. Room [16] temperature ionic liquids have been employed to recover rhodium, and ruthenium [17] also recently.
See also • Platinum group metals in Africa • Merensky Reef
External links • • • •
Platinum Today: The world's leading authority on platinum group metals [18] Platinum Group Spot Prices [19] USGS page on PGM's [20] Platinum Metals Review: the quarterly E-journal of research on the platinum metals and of developments in their application in industry [21]
References [1] Harris, D. C.; Cabri L. J. (1991). "Nomenclature of platinum-group-element alloys; review and revision". The Canadian Mineralogist 29 (2): 231 – 237. [2] Weeks, M. E. (1968). Discovery of the Elements (7 ed.). Journal of Chemical Education. pp. 385 – 407. ISBN 0848685792. OCLC 23991202. [3] Woods, Ian (2004). The Elements: Platinum (http://books.google.com/?id=hy2WcbKpXSkC&printsec=frontcover). Benchmark Books. ISBN 978-0761415503. . [4] Hunt, L. B.; Lever, F. M. (1969). "Platinum Metals: A Survey of Productive Resources to industrial Uses" (http://www. platinummetalsreview.com/pdf/pmr-v13-i4-126-138.pdf). Platinum Metals Review 13 (4): 126 – 138. . Retrieved 2009-10-02. [5] Xiao, Z.; Laplante, A. R. (2004). "Characterizing and recovering the platinum group minerals —a review". Minerals Engineering 17: 961 – 979. doi:10.1016/j.mineng.2004.04.001. [6] "Platinum – Group Metals" (http://minerals.usgs.gov/minerals/pubs/commodity/platinum/platimcs07.pdf) (PDF). U.S. Geological Survey, Mineral Commodity Summaries. January 2007. . Retrieved 2008-09-09. [7] Emsley, J. (2003). "Iridium" (http://books.google.com/?id=j-Xu07p3cKwC&pg=PA202). Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 201 – 204. ISBN 0198503407. . [8] Chevalier, Patrick (?). "Mineral Yearbook: Platinum Group Metals" (http://www.nrcan-rncan.gc.ca/mms-smm/busi-indu/cmy-amc/ content/2004/71.pdf). Natural Resources Canada. . Retrieved 2008-10-17. [9] R. J. Newman, F. J. Smith (1970). "Platinum Metals from Nuclear Fission – an evaluation of their possible use by the industry" (http://www. platinummetalsreview.com/dynamic/article/view/pmr-v14-i3-088-092). Platinum Metals Review 14 (3): 88. . [10] Zdenek Kolarik, Edouard V. Renard (2003). "Recovery of Value Fission Platinoids from Spent Nuclear Fuel; PART I: general considerations and basic chemistry" (http://www.platinummetalsreview.com/dynamic/article/view/pmr-v47-i2-074-087). Platinum Metals Review 47 (2): 74. . [11] Kolarik, Zdenek; Renard, Edouard V. (2005). "Potential Applications of Fission Platinoids in Industry" (http://www. platinummetalsreview.com/dynamic/article/view/49-2-79-90). Platinum Metals Review 49: 79. doi:10.1595/147106705X35263. . [12] Zdenek Kolarik, Edouard V. Renard (2003). "Recovery of Value Fission Platinoids from Spent Nuclear Fuel; PART II: Separation process" (http://www.platinummetalsreview.com/dynamic/article/view/pmr-v47-i3-123-131). Platinum Metals Review 47 (3): 123. .
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Platinum group
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[13] Jayakumar, M; Venkatesan, K; Srinivasan, T; Rao, P (2009). "Studies on the feasibility of electrochemical recovery of palladium from high-level liquid waste". Electrochimica Acta 54: 1083. doi:10.1016/j.electacta.2008.08.034. [14] Pokhitonov, Yu. A.; Romanovskii, V. N. (2005). "Palladium in Irradiated Fuel. Are There Any Prospects for Recovery and Application?". Radiochemistry 47: 1. doi:10.1007/s11137-005-0040-7. [15] Jayakumar, M; Venkatesan, K; Srinivasan, T (2007). "Electrochemical behavior of fission palladium in 1-butyl-3-methylimidazolium chloride". Electrochimica Acta 52: 7121. doi:10.1016/j.electacta.2007.05.049. [16] Jayakumar, M; Venkatesan, K; Srinivasan, T (2008). "Electrochemical behavior of rhodium(III) in 1-butyl-3-methylimidazolium chloride ionic liquid". Electrochimica Acta 53: 2794. doi:10.1016/j.electacta.2007.10.056. [17] Jayakumar, M; Venkatesan, K.A.; Srinivasan, T.G.; Vasudeva Rao, P.R. (2008). "Electrochemical behavior of ruthenium (III), rhodium (III) and palladium (II) in 1-butyl-3-methylimidazolium chloride ionic liquid". Electrochimica Acta 54: 2747. doi:10.1016/j.electacta.2009.06.043. [18] http://www.platinum.matthey.com/ [19] http://www.kitco.com/market/ [20] http://minerals.usgs.gov/minerals/pubs/commodity/platinum/ [21] http://www.platinummetalsreview.com/dynamic/
Post-transition metal Group # 12 13 14 15 Period 4
30 31 Zn Ga
5
48 49 50 Cd In Sn
6
80 81 82 83 Hg Tl Pb Bi
Atomic numbers show state at STP
Solids Liquids
In chemistry, the term post-transition metal is used to describe the category of metallic elements to the right of the transition elements on the periodic table[1] [2] . IUPAC defines "transition elements" as either the elements in groups 3 – 11 or the elements in groups 3 – 12[3] . According to the first definition, post-transition metals include group 12 —zinc, cadmium, and mercury. This collection of elements is illustrated by the element boxes to the right. Occasionally germanium, antimony, and/or polonium are also included as metals, although these are usually considered to be metalloids. According to the second definition of transition elements, group 12 would not be included as a post-transition metal. An examination of textbooks and monographs in 2003 revealed that both definitions are used with roughly equal frequency[4] . In the 1950s, most inorganic chemistry textbooks defined transition elements as excluding group 11 —copper, silver, and gold in addition to group 12[4] . A third definition of post-transition metal that includes group 11 and group 12 elements is no longer recommended by IUPAC[3] but is still used on occasion[5] [6] .
Post-transition metal
136
Poor metals The trivial name poor metals is sometimes applied to the metallic elements in the p-block of the periodic table. Their melting and boiling points are generally lower than those of the transition metals and their electronegativity higher, and they are also softer. They are distinguished from the metalloids, however, by their significantly higher boiling points in the same row. "Poor metals" is not a rigorous IUPAC-approved nomenclature, but the grouping is generally taken to include aluminium, gallium, indium, tin, thallium, lead, and bismuth. Occasionally germanium, antimony, and polonium are also included, although these are usually considered to be metalloids or "semi-metals". Elements 113, 114, 115, and 116, which are currently allocated the systematic names ununtrium, ununquadium, ununpentium, and ununhexium, would likely exhibit properties characteristic of poor metals; however sufficient quantities of them have not yet been synthesized to examine their chemical properties. 13
14
15
16
Al Aluminium
Ga
Ge
Gallium
Germanium
In
Sn
Sb
Indium
Tin
Antimony
Tl
Pb
Bi
Po
Thallium
Lead
Bismuth
Polonium
Uut
Uuq
Uup
Uuh
ununtrium ununquadium ununpentium ununhexium
External links • Patent-Invent - Poor Metals Quick Facts [7] • Royal Armouries - Poor Metals [8]
References [1] General Chemistry: Principles and Structure (http://books.google.com/books?&id=UVlGAAAAYAAJ&dq=post-transition+metal+ brady&q=post-transition+"just+to+the+right"#search_anchor) (5th ed.), by James E. Brady, page 96. Published by Wiley, 1990. ISBN 0471621315, 9780471621317 [2] Instant Notes in Inorganic Chemistry (http://books.google.com/books?id=8yQOhvD3tWcC&pg=PA185#v=onepage&q=&f=false) (2nd ed.), by P.A. Cox, page 185 – 186. Published by Garland Science/BIOS Scientific Publishers, 2004. ISBN 1859962890, 9781859962893 [3] Nomenclature of Inorganic Chemistry, IUPAC Recommendations (2005) (http://www.iupac.org/publications/books/rbook/ Red_Book_2005. pdf) IR 3-6.2 p 51 [4] William B. Jensen (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table". Journal of Chemical Education 80 (8): 952 – 961. [5] Introductory solid state physics (http://books.google.com/books?id=w6eAC9y47_4C&pg=PA216#v=onepage&q=&f=false) (2nd ed.),
by H.P. Myers, page 216. Published by Taylor & Francis, 1997. ISBN 074840659X, 9780748406593 [6] Bioinorganic Chemistry (http://books.google.com/books?id=tCOyYeekTSEC&pg=PA69#v=onepage&f=false), by K. Hussain Reddy, page 69. Published by New Age International, 2003. ISBN 8122414370, 9788122414370 [7] http://www.chemistry.patent-invent.com/chemistry/poor_metals.html [8] http://www.royalarmouries.org/extsite/view.jsp?sectionId=2884
Transactinide element
Transactinide element In chemistry, transactinide elements (also, transactinides, or super-heavy elements) are the chemical elements with atomic numbers greater than those of the actinides, the heaviest of which is lawrencium (103).[1] [2] Transactinide elements are also transuranic elements, that is, have an atomic number greater than that of uranium (92), an actinide. The further distinction of having an atomic number greater than the actinides is significant in several ways: • The transactinide elements all have electrons in the 6d subshell in their ground state (and thus are placed in the d-block). The last actinide, lawrencium, also has one electron in the 6d subshell. • Except for dubnium, even the longest-lasting isotopes of transactinide elements have extremely short half-lives, measured in seconds, or smaller units. • The element naming controversy involved the first five or six transactinide elements. These elements thus used three-letter systematic names for many years after their discovery had been confirmed. (Usually the three-letter names are replaced with two-letter names relatively shortly after a discovery has been confirmed.) Transactinides are radioactive and have only been obtained synthetically in laboratories. None of these elements has ever been collected in a macroscopic sample. Transactinide elements are all named after nuclear physicists and chemists or important locations involved in the synthesis of the elements. Chemistry Nobelist Glenn T. Seaborg who first proposed the actinide concept which led to the acceptance of the actinide series also proposed the existence of a transactinide series ranging from element 104 to 121 and a superactinide series approximately spanning elements 122 to 153. The transactinide seaborgium is named in his honor. The term transactinide is an adjective, and is not commonly used alone as a noun to refer to the transactinide elements.
List of the transactinide elements • 104 Rutherfordium, Rf • 105 Dubnium, Db • 106 Seaborgium, Sg • 107 Bohrium, Bh • 108 Hassium, Hs • 109 Meitnerium, Mt • 110 Darmstadtium, Ds • 111 Roentgenium, Rg • 112 Copernicium, Cn • 113 Ununtrium, Uut* • 114 Ununquadium, Uuq* • 115 Ununpentium, Uup* • 116 Ununhexium, Uuh* • 117 Ununseptium, Uus* • 118 Ununoctium, Uuo* * The synthesis of these elements has not been officially attested by IUPAC, while in several cases previous syntheses have been confirmed by other institutions or other methods. The names and symbols given are provisional as no names for the elements have been agreed on.
137
Transactinide element
138
See also • Transuranium element • Bose-Einstein condensate (also known as Superatom) • Island of stability
References [1] IUPAC Provisional Recommendations for the Nomenclature of Inorganic Chemistry (2004) (http://www.iupac.org/reports/provisional/ abstract04/connelly_310804.html) (online draft of an updated version of the " Red Book " IR 3-6) [2] Morss, Lester R.; Edelstein, Norman M.; Fuger, Jean, eds (2006). The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands: Springer. ISBN 13978-1-4020-3555-5.
Transuranium element In chemistry, transuranium elements (also known as transuranic elements) are the chemical elements with atomic numbers greater than 92 (the atomic number of uranium). None of these elements are stable; they decay radioactively into other elements.
Overview Of the elements with atomic numbers 1 to 92, all but four (technetium, promethium, astatine, and francium) occur in easily detectable quantities on Earth, having stable, or very long half-life isotopes, or are created as common products of the decay of uranium. All of the elements with higher atomic numbers, however, have been first discovered in the laboratory, other than neptunium and plutonium. They are all radioactive, with a half-life much shorter than the age of the Earth, so any atoms of these elements, if they ever were present at the Earth's formation, have long since decayed. Trace amounts of neptunium and plutonium form in some uranium-rich rock, and small amounts are produced during atmospheric tests of atomic weapons. The Np and Pu generated are from neutron capture in uranium ore with two subsequent beta decays ( 238U → 239U → 239Np → 239Pu).
Periodic table with elements colored according to the half-life of their most stable isotope.
Stable elements; Radioactive elements with very long-lived isotopes. Their half-lives of over four million years confer them very small, if not negligible radioactivities; Radioactive elements that may present low health hazards. Their most stable isotopes have half-lives between 800 and 34.000 years. Because of this, they usually have some commercial applications; Radioactive elements that are known to pose high safety risks. Their most stable isotopes have half-lifes between one day and 103 years. Their radioactivities confers them little potential for commercial uses; Highly radioactive elements. Their most stable isotopes have half-lifes between one day and several minutes. They pose severe health risks. Few of them receive uses outside basic research; Extremely radioactive elements. Very little is known about these elements due to their extreme instability and radioactivity.
Transuranium element Those that can be found on Earth now are artificially generated synthetic elements, via nuclear reactors or particle accelerators. The half lives of these elements show a general trend of decreasing with atomic number. There are exceptions, however, including dubnium and several isotopes of curium. Further anomalous elements in this series have been predicted by Glenn T. Seaborg, and are categorised as the “island of stability.” Heavy transuranic elements are difficult and expensive to produce, and their prices go up rapidly with atomic number. As of 2008, weapons-grade plutonium cost around $4,000/gram (or roughly 150 times more than gold),[1] and californium cost $60,000,000/gram.[2] Due to production difficulties, none of the elements beyond californium have industrial applications or were ever produced in macroscopic quantities. Transuranic elements that have not been discovered, or have been discovered but are not yet officially named, use IUPAC's systematic element names. The naming of transuranic elements is a source of controversy.
Discovery and naming of transuranium elements The majority of the transuranium elements were produced by three groups: • A group at the University of California, Berkeley, under three different leaders: • Edwin Mattison McMillan, first to produce a transuranium element: • 93. neptunium, Np, named after the planet Neptune, as it follows uranium and Neptune follows Uranus in the planetary sequence (1940). • Glenn T. Seaborg, next in order, who produced: • 94. plutonium, Pu, named after the dwarf planet Pluto, following the same naming rule as it follows neptunium and Pluto follows Neptune in the pre-2006 planetary sequence (1940). • 95. americium, Am, named because it is an analog to europium, and so was named after the continent where it was first produced (1944). • 96. curium, Cm, named after Pierre and Marie Curie, famous scientists who separated out the first radioactive elements (1944). • 97. berkelium, Bk, named after the city of Berkeley, where the University of California, Berkeley is located (1949). • 98. californium, Cf, named after the state of California, where the university is located (1950). • Albert Ghiorso, who had been on Seaborg's team when they produced curium, berkelium, and californium, took over as director to produce: • 99. einsteinium, Es, named after the theoretical physicist Albert Einstein (1952). • 100. fermium, Fm, named after Enrico Fermi, the physicist who produced the first controlled chain reaction (1952). • 101. mendelevium, Md, named after the Russian chemist Dmitri Mendeleev, credited for being the primary creator of the periodic table of the chemical elements (1955). • 102. nobelium, No, named after Alfred Nobel (1956). • 103. lawrencium, Lr, named after Ernest O. Lawrence, a physicist best known for development of the cyclotron, and the person for whom the Lawrence Livermore National Laboratory and the Lawrence Berkeley National Laboratory (which hosted the creation of these transuranium elements) are named (1961). • A group at the Joint Institute for Nuclear Research in Dubna, Russia (then the Soviet Union) who produced: • 104. rutherfordium, Rf, named after Ernest Rutherford, who was responsible for the concept of the atomic nucleus (1966). • 105. dubnium, Db, an element that is named after the city of Dubna, where the JINR is located. Also known in Western circles as "hahnium" in honor of Otto Hahn (1968).
139
Transuranium element
140
• 106. seaborgium, Sg, named after Glenn T. Seaborg. This name caused controversy because Seaborg was still alive, but eventually became accepted by international chemists (1974). • 107. bohrium, Bh, named after the Danish physicist Niels Bohr, important in the elucidation of the structure of the atom (1981). • A group at the Gesellschaft für Schwerionenforschung (Society for Heavy Ion Research) in Darmstadt, Hessen, Germany, under Peter Armbruster, who produced: • 108. hassium, Hs, named after the Latin form of the name of Hessen, the German Bundesland where this work was performed (1984). • 109. meitnerium, Mt, named after Lise Meitner, an Austrian physicist who was one of the earliest scientists to become involved in the study of nuclear fission (1982). • 110. darmstadtium, Ds named after Darmstadt, Germany, the city in which this work was performed (1994). • 111. roentgenium, Rg named after Wilhelm Conrad Röntgen, discoverer of X-rays (1994). • 112. copernicium, Cn named after astronomer Nicolaus Copernicus (1996).
List of the transuranic elements •
Actinides • • • • • • • • • • •
93 neptunium Np 94 plutonium Pu 95 americium Am 96 curium Cm 97 berkelium Bk 98 californium Cf 99 einsteinium Es 100 fermium Fm 101 mendelevium Md 102 nobelium No 103 lawrencium Lr
•
Transactinide elements • • • • • • • • • • • • • • •
104 rutherfordium Rf 105 dubnium Db 106 seaborgium Sg 107 bohrium Bh 108 hassium Hs 109 meitnerium Mt 110 darmstadtium Ds 111 roentgenium Rg 112 copernicium Cn 113 ununtrium Uut* 114 ununquadium Uuq* 115 ununpentium Uup* 116 ununhexium Uuh* 117 ununseptium Uus* 118 ununoctium Uuo*
*The existence of these elements has been confirmed, however the names and symbols given are provisional as no names for the elements have been agreed on.
Super-heavy atoms Super-heavy atoms, (super heavy elements , commonly
abbreviated SHE), are the transactinide elements beginning with rutherfordium (atomic number 104). They have only been made artificially, and currently serve no useful purpose because their short half-lives cause them to decay after a few minutes to just a few milliseconds, which also makes them extremely hard to study.[3] [4] Super-heavy atoms have all been created during the latter half of the 20th century and are continually being
Position of the super-heavy elements in the periodic table.
Transuranium element created during the 21st century as technology advances. They are created through the bombardment of elements in a particle accelerator, for example the nuclear fusion of californium-249 and carbon-12 creates rutherfordium. These elements are created in quantities on the atomic scale and no method of mass creation has been found.[3]
See also • • • • •
Bose-Einstein condensate (also known as Superatom) Island of stability Minor actinides Waste Isolation Pilot Plant, repository for transuranic waste Extension of the periodic table beyond the seventh period
Further reading • Annotated bibliography for the transuranic elements [5] from the Alsos Digital Library for Nuclear Issues. • Transuranium elements [6] [7] • Super Heavy Elements network official website (network of the European integrated infrastructure initiative EURONS) • Prof. Amnon Marinov's Site with related publications [8] [9] • Darmstadium and beyond
References [1] "Price of Plutonium" (http://hypertextbook.com/facts/2008/AndrewMorel.shtml). The Physics Factbook. . [2] Rodger C. Martin and Steven E. Kos. "Applications and Availability of Californium-252 Neutron Sources for Waste Characterization" (http:/ /www.ornl.gov/~webworks/cpr/pres/108701_.pdf) (pdf). . [3] Heenen, P. H.; Nazarewicz, W. (2002). "Quest for superheavy nuclei". Europhysics News 33: 5. doi:10.1051/epn:2002102. [4] Greenwood, N. N. (1997). "Recent developments concerning the discovery of elements 100-111". Pure and Applied Chemistry 69: 179. doi:10.1351/pac199769010179. [5] http://alsos.wlu.edu/qsearch.aspx?browse=science/Transuranium+Elements [6] http://web.fccj.org/~ethall/uranium/uranium.htm [7] http://www.transfermium.net/ [8] http://www.phys.huji.ac.il/~marinov/index.htm [9] http://pubs.acs.org/cen/80th/darmstadtium.html
141
Transition metal
142
Transition metal The term transition metal (sometimes also called a transition element) has two possible meanings: • In the past it referred to any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. All elements in the d-block are metals. (In actuality, the f-block is also included in the form of the lanthanide and actinide series.) • The modern, IUPAC definition[1] states that a transition metal is "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." Group 12 elements are not transition metals in this definition. Jensen[2] has reviewed the historical usage of the terms transition element (or metal) and d-block. The word "transition" was first used to describe the elements now known as the d-block by the English chemist Charles Bury in 1921, who referred to a transition series of elements during the change of an inner layer of electrons (for example n=3 in the 4th row of the periodic table) from a stable group of 8 to one of 18, or from 18 to 32. [3]
Classification In the d -block the atoms of the elements have between 1 and 10 d electrons. The following table shows IUPAC definition of what is called "transition metal". Group
3
4
5
6
7
8
Mn 25 Fe 26
9
10
11
12
Co 27
Ni 28
Cu 29
Zn 30
Period 4 Sc 21 Ti 22
V 23
Cr 24
Period 5 Y 39 Zr 40
Nb 41
Mo 42 Tc 43
Ru 44
Rh 45
Pd 46
Ag 47
Cd 48
Period 6
Hf 72
Ta 73
W 74
Re 75
Os 76
Ir 77
Pt 78
Au 79
Hg 80
Period 7
Rf 104 Db 105
Sg 106
Bh 107 Hs 108 Mt 109 Ds 110 Rg 111 Cn 112
Atoms of scandium and yttrium have a single d electron in the outermost shell, and thus are mostly considered transition metals. However, as all their compounds contain the ions Sc3+ and Y3+ in which there are no d electrons, these elements are not universally considered transition metals. This is disputed by people saying that classification must be lead by neutral atoms properties, and fact zirconium and titanium compounds also don't contain d-electrons. The two vacant places in group 3 is due to the fact that, for period 6, the place is disputed between lanthanum and lutetium and for period 7, between actinium and lawrencium . To prevent this, IUPAC placed all of these 4 as lanthanoids/actinoids, which makes both series 15 elements long, despite fact there are only 14 f-electrons possible and then f-block can't be any longer. This means IUPAC lanthanoids/actinoids contains one d-element. Modern scientists usually claim Lu and Lr as d-block elements rather La and Ac, however, both variants are widely used. The electronic structure of transition metal atoms can be written, with a few minor exceptions, as [Ng]ns2(n-1)d m, as the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals the situation is reversed so that the s electrons have higher energy. Consequently an ion such as Fe 2+ has no s electrons, it has the electronic configuration [Ar]3d6 as compared with the configuration of the atom, [Ar]4s23d6. According to IUPAC, Zinc, cadmium, mercury and copernicium are transition metals, although some say they are not.[2] as they have the electronic configuration [Ng]d 10s2, with no incomplete d shell.[4] People who count Zn, Cd and Hg as post-transition mention that in the oxidation state +2 the ions have the electronic configuration [ Ng ] d10. and while compounds in the +1 oxidation state, such as Hg 2+, are known there are no unpaired electrons because of 2 the formation of a covalent bond between the two atoms of the dimer. However, it is opposed by opinion that d-block must be equal by its content to transition metals, because if mercury has no incompleteness in d-orbital, thus ytterbium has no incompleteness in f-orbital and is a transition metal. Also, judging on some separated properties
Transition metal (like Irving-Williams series of stability constants of complexes) is also mostly denied.
Characteristic properties There are a number of properties shared by the transition elements that are not found in other elements, which results from the partially filled d shell. These include • the formation of compounds whose colour is due to d - d electronic transitions • the formation of compounds in many oxidation states, due to the relatively low reactivity of unpaired d electrons.[5] • the formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are also paramagnetic (e.g. nitric oxide, oxygen)
Coloured compounds Colour in transition-series metal compounds is generally due to electronic transitions of two principal types. • charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a From left to right, aqueous solutions of: Co(NO3)2 (red); K2Cr2O7 ligand-to-metal charge-transfer (LMCT) transition. (orange); K2CrO4 (yellow); NiCl2 (turquoise); CuSO4 (blue); KMnO4 These can most easily occur when the metal is in a (purple). high oxidation state. For example, the colour of chromate, dichromate and permanganate ions is due to LMCT transitions. A metal-to ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily oxidised. Mercuric iodide, HgI2, is red because of a MLCT transition. As this example shows, charge transfer transitions are not restricted to transition metals.[6] • d -d transitions. An electron jumps from one d-orbital to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using crystal field theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on Tanabe-Sugano diagrams. In centrosymmetric complexes, such as octahedral complexes, d -d transitions are forbidden by the Laporte rule and only occur because of vibronic coupling in which a molecular vibration occurs together with a d-d transition. Tetrahedral complexes have somewhat more intense colour because mixing d and p orbitals is possible when there is no centre of symmetry, so transitions are not pure d-d transitions. The molar absorptivity (ε) of bands caused by d-d − − − transitions are relatively low, roughly in the range 5-500 M 1cm 1 (where M = mol dm 3).[7] Some d -d transitions 5 are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese(II), which has a d configuration in which all five electron has parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. In fact many compounds of manganese(II) appear almost colourless. The −1 −1 2+ spectrum of [Mn(H2O)6] shows a maximum molar absorptivity of about 0.04 M cm in the visible spectrum.
143
Transition metal
144
Oxidation states A characteristic of transition metals is that they exhibit two or more oxidation states, usually differing by one. For − example, compounds of vanadium are known in all oxidation states between −1, such as [V(CO)6] , and +5, such as − VO3 . 4
Main group elements in groups 13 to 17 also exhibit multiple oxidation states. The "common" oxidation states of these elements typically differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No compound of Ga(II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. The only compounds in which gallium has a − formal oxidation state of +2 are dimeric compounds, such as [Ga2Cl6]2 , which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom. [8] Thus the main difference in oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons. The maximum oxidation state in the first row transition metals is equal to the number of valence electron from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second and third rows the maximum − occurs with ruthenium and osmium (+8). In compounds such as [MnO4] and OsO4 the elements achieve a stable octet by forming four covalent bonds. −
The lowest oxidation states are exhibited in such compounds as Cr(CO) 6 (oxidation state zero) and [Fe(CO)4]2 (oxidation state −2) in which the 18-electron rule is obeyed. These complexes are also covalent.
Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution the ions are hydrated by (usually) six water molecules arranged octahedrally.
Magnetism Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.[9] In octahedral complexes with between four and seven d electrons both high spin and low spin states are possible. Tetrahedral − transition metal complexes such as [FeCl ]2 are high spin because the crystal field splitting is small so that the 4 energy to be gained by virtue of the electrons being in lower energy orbitals is always less that the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d 6 and square-planar d 8 complexes. In these cases, crystal field splitting is such that all the electrons are paired up. Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Antiferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.
See also • Inner transition element, a name given to any member of the f-block • Ligand field theory a development of crystal field theory taking covalency into account • Post-transition metal
References [1] International Union of Pure and Applied Chemistry. " transition element (http://goldbook.iupac.org/T06456.html)". Compendium of Chemical Terminology Internet edition. [2] Jensen, William B. (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table" (http://www.uv.es/~borrasj/ingenieria_web/ temas/tema_1/lecturas_comp/p952.pdf). Journal of Chemical Education 80 (8): 952 – 961. doi:10.1021/ed080p952. . [3] "Langmuir's theory of the arrangement of electrons in atoms and molecules" C.R. Bury, J. Amer. Chem. Soc. 43, p.1602-1609 (1921) [4] Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley.
Transition metal [5] Matsumoto, Paul S (2005). "Trends in Ionization Energy of Transition-Metal Elements" (http://www. jce.divched.org/Journal/Issues/ 2005/Nov/abs1660.html). Journal of Chemical Education 82: 1660. doi:10.1021/ed082p1660. . [6] T.M. Dunn in Lewis, J.; Wilkins,R.G. (1960). Modern Coordination Chemistry. New York: Interscience., Chapter 4, Section 4, "Charge Transfer Spectra", pp. 268-273. [7] Orgel, L.E. (1966). An Introduction to Transition-Metal Chemistry, Ligand field theory (2nd. ed.). London: Methuen. [8] Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419 p. 240 [9] Figgis, B.N.; Lewis, J. (1960). Lewis, J. and Wilkins, R.G.. ed. The Magnetochemistry of Complex Compounds. Modern Coordination Chemistry. New York: Interscience. pp. 400 – 454.
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146
See also Table of nuclides A table of nuclides or chart of nuclides is a two-dimensional graph in which one axis represents the number of neutrons and the other represents the number of protons in an atomic nucleus. Each point plotted on the graph thus represents the nuclide of a real or hypothetical chemical element. This system of ordering nuclides can offer a greater insight into the characteristics of isotopes than the better-known periodic table, which shows only elements instead of each of their isotopes.
A chart of nuclides (cut into three parts for better presentation).
Description and utility A chart or table of nuclides (capitalization optional) is a simple map to the nuclear, or radioactive, behaviour of nuclides, as it distinguishes the isotopes of an element. It contrasts with a periodic table, which only maps their chemical behavior, since isotopes of the same element do not differ chemically. Nuclide charts organize isotopes along the X axis by their numbers of neutrons and along the Y axis by their numbers of protons, out to the limits of the neutron and proton drip lines. This representation was first published by Giorgio Fea in 1935,[1] and expanded by Emilio Segrè in 1945 or G. Seaborg. In 1958, Walter Seelmann-Eggebert and Gerda Pfennig published the first edition of the Karlsruhe Nuclide Chart. Its 7th edition was made available in 2006. Today, one finds several nuclide charts, four of them have a wide distribution: the Karlsruhe Nuclide Chart, the Strasbourg Universal Nuclide Chart, the Chart of the Nuclides from the JAEA and the Nuclide Chart from Knolls Atomic Power Laboratory.[2] It has become a basic tool of the nuclear community.
Trends in the Chart of Nuclides 5
6
7
He
8
Li
9
6
7
H
8
He
9
Li
7
9
He
10
8
10
He
9
H
Ne
11
10
B
11
C
12
13
14
F
10
11
B
12
C
13
14
15
F
16
Li
11
12
B
13
C
14
15
16
F
17
18
Na
11
Li
12
13
B
14
C
15
16
17
F
18
19
Na
20
12
Li
13
14
B
15
C
16
17
18
F
19
20
Na
21
Be Be Be Be Be
N N N N N
O O O O O
Ne Ne Ne Ne
Na
12 Mg
Mg Mg
13 Al 22
14
Al Si
• Isotopes are nuclides with the same number of protons but differing numbers of neutrons; that is, they have the same atomic number and are therefore the same chemical element. Isotopes neighbor each other vertically e.g. Carbon-12, Carbon-13, and Carbon-14 in the sample table above. • Isotones are nuclides with the same number of neutrons but differing number of protons. Isotones neighbor each other horizontally. Example: Carbon-14, Nitrogen-15, Oxygen-16 in the sample table above. • Isobars are nuclides with the same number of nucleons, i.e. mass number, but different numbers of protons and different number of neutrons. Isobars neighbor each other diagonally from lower-left to upper-right. Example:
Table of nuclides
• • • •
•
147
Carbon-14, Nitrogen-14, Oxygen-14 in the sample table above. Beyond the neutron drip line along the lower left, nuclides decay by neutron emission. Beyond the proton drip line along the upper right, nuclides decay by proton emission. Drip lines have only been established for some elements. The island of stability is a hypothetical region of the table of nuclides that contains isotopes far more stable than other transuranic elements. There are no stable atoms having an equal number of protons and neutrons in their nuclei with atomic number greater than 20 as can be readily "read" from the chart. Nuclei of greater atomic number require an excess of neutrons for stability. There are no stable atoms having atomic number greater than Z=82 (lead)[3] . Atoms with atomic numbers of 82 and lower all have stable isotopes, with the exceptions of technetium (Z=43) and promethium (Z=61).
Available representations Charts of the nuclides article
description
Table of nuclides (complete)
Presents the data via a single, contiguous chart that requires both vertical and horizontal scrolling to view all its contents (262 kB total HTML download).
Table of nuclides (segmented, wide)
Presents the data via four separate charts, each typically with 30 elements. Depending on the browser, no horizontal scrolling is required in window widths of at least 1225 to 1440 pixels (311 kB total HTML download).
Table of nuclides (segmented, narrow)
Presents the data via eight separate charts, each typically with 15 elements. Horizontal scrolling is not required for all but the smallest computer monitors (321 KB total HTML download).
Table of nuclides (sorted by half-life)
Presents the data in a one-dimensional list where all nuclides are sorted by their half-life, including specific mass excess and decay-modes, no horizontal scrolling is required (95 kB total HTML download).
Table of nuclides (combined) Provides both the eight-chart, segmented presentation and the single, contiguous chart. Provides quick-jump hyperlinks to jump between the two. Features expanded introductory text for first-time readers. (588 kB total HTML download). File:NuclideMap.PNG
Single image (not HTML) of the National Nuclear Data Center chart from the NuDat 2 database (3.9 MB)
Articles on isotopes of an element article
description
Index to isotope pages
A periodic table that provides links to a separate article on each element and its isotopes.
Isotope lists
A page that provides data on the isotopes of each element in groups of 24 elements.
Table of nuclides
148
See also • List of nuclides. Presents information in list form, and in order of stability, for the 905 nuclides which are stable, or are radioactive with half lives greater than 60 minutes.
External links • Karlsruhe Nuclide Chart [4] • Universal Nuclide Chart
[5]
• Interactive Chart of Nuclides (Brookhaven National Laboratory) [6] • The Lund/LBNL Nuclear Data Search [7] • Another example of a Chart of Nuclides from Korea [8] • YChartElements [19] dynamic periodic table and chart of the nuclides, a Yoix application
•
Compact Chart of Nuclides (non-standard representation with elements along a diagonal) 70x70. [9]
[20] The Live Chart of Nuclides - IAEA • in Java [20] or HTML [21] [10] • Map of the Nuclides
References [1] Georgio Fea. Il Nuovo Cimento 2 (1935) 368 [2] "What We Do: The Chart of Nuclides" (http://www.knollslab.com/nuclides.html). Knolls Atomic Power Laboratory. . Retrieved 14 May 2009. [3] Holden,CRC Handbook of Chemistry and Physics, 90th Edition §11 [4] http://www.karlsruhenuclidechart.net [5] http://www.nucleonica.net/unc.aspx [6] http://www.nndc.bnl.gov/chart [7] http://nucleardata.nuclear.lu.se/nucleardata/toi/ [8] http://atom.kaeri.re.kr [9] http://adavidstubbs.home.comcast.net/~adavidstubbs/Quark/Isotope_table_(compact).htm [10] http://t2.lanl.gov/data/map.html
Island of stability
149
Island of stability The island of stability is a term from nuclear physics that describes the possibility of elements with particularly stable "magic numbers" of protons and neutrons. This would allow certain isotopes of some transuranium elements to be far more stable than others; that is, to decay much more slowly (with half-lives of at least minutes or days, compared to fractions of a second; some have even suggested the possibility of half-lives on the order of millions of years[1] ).
3-dimensional rendering of the theoretical Island of Stability.
History The idea of the island of stability was first proposed by Glenn T. Seaborg. The hypothesis is that the atomic nucleus is built up in "shells" in a manner similar to the electron shells in atoms. In both cases shells are just groups of quantum energy levels that are relatively close to each other. Energy levels from quantum states in two different shells will be separated by a relatively large energy gap. So when the number of neutrons and protons completely fill the energy levels of a given shell in the nucleus, the binding energy per nucleon will reach a local maximum and thus that particular configuration will have a longer lifetime than nearby isotopes that do not have filled shells.[2] A filled shell would have "magic numbers" of neutrons and protons. One possible magic number of neutrons for spherical nuclei is 184, and some possible matching proton numbers are 114, 120 and 126 – which would mean that the most stable spherical isotopes would be ununquadium-298, unbinilium-304 and unbihexium-310. Of particular note is Ubh-310, which would be "doubly magic" (both its proton number of 126 and neutron number of 184 are thought to be magic) and thus the most likely to have a very long half-life. (The next lighter doubly-magic spherical nucleus is lead-208, the heaviest stable nucleus and most stable heavy metal.) Isotopes of elements in the range between 110 through 114 have been found to decay more slowly than isotopes of nuclei nearby in the periodic table. However, recent research indicates that large nuclei are deformed, causing magic numbers to shift. Hassium-270 is now believed to be doubly-magic nucleus, with deformed magic numbers 108 and 162. Its half-li fe may be as high as 23 seconds.[3] [4]
Island of stability
150
Half-lives of large isotopes Fermium is the heaviest element that can be produced in a nuclear reactor. The stability (half-life of the longest-lived isotope) of elements generally decreases from element 101 to element 109 and then approaches an island of stability with longer-lived isotopes in the range of elements 111 and 114.[5] The longest-lived observed isotopes are shown in the following table.
Periodic table with elements colored according to the half-life of their most stable isotope.
Stable elements. Radioactive elements with half-lives of over four million years. Half-lives between 800 and 34,000 years. Half-lives between 1 day and 103 years. Half-lives ranging between several minutes and 1 day. Extremely radioactive elements with half-lives less than a minute.
Known isotopes of elements 100 through 118 [5] [6] Number
Name
Longest-lived
Half-life
Article
measured isotope 100
Fermium
257
101 days
Isotopes of fermium
101
Mendelevium
258
52 days
Isotopes of mendelevium
102
Nobelium
259
58 minutes
Isotopes of nobelium
103
Lawrencium
262
3.6 hours
Isotopes of lawrencium
104
Rutherfordium
267
1.3 hours
Isotopes of rutherfordium
105
Dubnium
268
29 hours
Isotopes of dubnium
106
Seaborgium
271
1.9 minutes
Isotopes of seaborgium
107
Bohrium
270
61 seconds
Isotopes of bohrium
108
Hassium
277
16.5 minutes Isotopes of hassium
109
Meitnerium
278
~8 seconds
Isotopes of meitnerium
110
Darmstadtium
281
11 seconds
Isotopes of darmstadtium
111
Roentgenium
281
22.8 seconds Isotopes of roentgenium
112
Copernicium
285
29 seconds
113
Ununtrium
286
19.6 seconds Isotopes of ununtrium
Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut
Isotopes of copernicium
Island of stability
151 114
Ununquadium
289
2.6 seconds
Isotopes of ununquadium
115
Ununpentium
289
220 ms
Isotopes of ununpentium
116
Ununhexium
293
61 ms
Isotopes of ununhexium
117
Ununseptium
294
78 ms
Isotopes of ununseptium
118
Ununoctium
294
0.89 ms
Isotopes of ununoctium
Uuq Uup Uuh Uus Uuo
(Note that for elements 109-118 the longest-lived known isotope is always the heaviest one discovered, making it likely that there are still longer-lived isotopes among the undiscovered heavier ones)
The half-lives of elements in the island are uncertain due to the small number of atoms manufactured and studied to date. Many physicists think they are relatively short, on the order of minutes, hours, or perhaps days. However, some 9 theoretical calculations indicate that their half-lives may be long (some calculations put it on the order of 10 years).[7] It is possible that these elements could have unusual chemical properties, and, if long-lived enough, various applications (such as targets in nuclear physics and neutron sources). However, the isotopes of several of these elements still have too few neutrons to be stable. The island of stability still has not been reached, since the island's "shores" are more neutron rich than nuclides that have been experimentally produced. The alpha-decay half-lives of 1700 nuclei with 100 ≤ Z ≤ 130 have been calculated in a quantum tunneling model with both experimental and theoretical alpha-decay Q-values.[8] [9] [10] [11] [12] [13] The theoretical calculations are in good agreement with the available experimental data.
Island of relative stability 232
Th (thorium), 235U and 238U (uranium) are the only naturally occurring isotopes beyond bismuth that are relatively stable over the current lifespan of the universe. Bismuth was found to be unstable in 2003, with an 19 α-emission half-life of 1.9×10 years for Bi-209. All other isotopes beyond bismuth are relatively or very unstable. So the main periodic table ends at bismuth, with an island at thorium and uranium. Between bismuth and thorium there is a "sea of instability", which renders such elements as astatine, radon, and francium extremely short-lived relative to all but the heaviest elements found so far. Current theoretical investigation indicates that in the region Z=106 – 108 and N≈160 – 164, a small ‘island/peninsula’ might be stable with respect to fission and beta decay, such superheavy nuclei undergoing only alpha decay. [9] [10] [11] Also, 298114 is not the center of the magic island as predicted earlier. [14] On the contrary, the nucleus with Z=110, N=183 appears to be near the center of a possible 'magic island' (Z=104 – 116, N≈176 – 186). In the N≈162 region the beta-stable, fission survived 268106 is predicted to have alpha-decay half-life ~3.2hrs that is greater than that (~28s) of the deformed doubly-magic 270108.[15] The superheavy nucleus 268106 has not been produced in the laboratory as yet (2009). For superheavy nuclei with Z>116 and N≈184 the alpha-decay half-lives are predicted to be less than one second. The nuclei with Z=120, 124, 126 and N=184 are predicted to form spherical doubly-magic nuclei and be stable with respect to fission. [16] Calculations in a quantum tunneling model show that such superheavy nuclei would undergo alpha decay within microseconds or less.[9] [10] [11]
Synthesis problems Manufacturing nuclei in the island of stability may be very difficult, because the nuclei available as starting materials do not deliver the necessary sum of neutrons. So for the synthesis of isotope 298 of element 114 by using plutonium and calcium, one would require an isotope of plutonium and one of calcium, which have together a sum of at least 298 nucleons (more is better, because at the nuclei reaction some neutrons are emitted). This would require, for example, the use of calcium-50 and plutonium-248 for the synthesis of element 114. However these isotopes (and heavier calcium and plutonium isotopes) are not available in weighable quantities. This is also the case for other
Island of stability
152
target-projectile combinations. However it may be possible to generate the isotope 298 of element 114, if the multi-nucleon transfer reactions would work in low-energy collisions of actinide nuclei.[17] One of these reactions may be: 248
Cm + 238U → 298Uuq + 186W + 2 1n
Quest for the island of stability "We search for the island of stability because, like Mount Everest, it is there. But, as with Everest, there is profound emotion, too, infusing the scientific search to test a hypothesis. The quest for the magic island shows us that science is far from being coldness and calculation, as many people imagine, but is shot through with passion, longing and romance." [18]
—Oliver Sacks
See also • Island of stability: Ununquadium — Unbinilium — Unbihexium • Table of nuclides — a visualization of the island of stability • Periodic table and Periodic table (extended)
External links • • • • • • • • • • • •
Hunting the biggest atoms in the universe [19] (July 23, 2008) The hunt for superheavy elements [20] (April 7, 2008) The synthesis of spherical superheavy nuclei in 48Ca induced reactions [21] (needs login so can not access !) Uut and Uup Add Their Atomic Mass to Periodic Table [22] (Feb 2004) New elements discovered and the island of stability sighted [23] (Aug 1999 - includes report on article later retracted) First postcard from the island of nuclear stability [24] (1999) Second postcard from the island of stability [25] (Oct 2001) Superheavy Elements "Island of Stability" [26] (single text slide - undated) Superheavy elements [27] (Jul 2004 Yuri Oganessian of JINR ) Can superheavy elements (such as Z=116 or 118) be formed in a supernova? Can we observe them? [28] NOVA - Island of Stability [29] New York Times Editorial by Oliver Sacks regarding the Island of Stability theory [30] (Feb 2004 re 113 and 115)
References [1] "Superheavy Element 114 Confirmed: A Stepping Stone to the Island of Stability" (http://www.physorg. com/news173028810.html). . Retrieved 11 October 2009. [2] "Shell Model of Nucleus" (http://hyperphysics.phy-astr.gsu.edu/hbase/nuclear/shell.html). HyperPhysics. Department of Physics and Astronomy, Georgia State University. . Retrieved 22 January 2007. [3] Dvořák, Jan (2007-07-12). "PhD. Thesis: Decay properties of nuclei close to Z = 108 and N = 162" (http://deposit.ddb.de/cgi-bin/ dokserv?idn=985213566&dok_var=d1&dok_ext=pdf&filename=985213566. pdf). Technischen Universität München. . [4] Dvorak, J.; Brüchle, W.; Chelnokov, M.; Dressler, R.; Düllmann, Ch.; Eberhardt, K.; Gorshkov, V.; Jäger, E. et al. (2006). "Doubly Magic Nucleus Hs162108270". Physical Review Letters 97 (24): 242501. doi:10.1103/PhysRevLett.97.242501. PMID 17280272. [5] Emsley, John (2001). Nature's Building Blocks ((Hardcover, First Edition) ed.). Oxford University Press. pp. (pages 143,144,458). ISBN 0198503407. [6] Alexandra Witze (April 6, 2010). ["http://www.sciencenews.org/view/generic/id/57964/title/Superheavy_element_117_makes_debut_" "Superheavy element 117 makes debut"]. "". Retrieved April 6, 2010. [7] Moller Theoretical Nuclear Chart 1997 (http://ie.lbl.gov/toipdf/theory.pdf) [8] P. Roy Chowdhury, C. Samanta, and D. N. Basu (January 26, 2006). "α decay half-lives of new superheavy elements" (http://link.aps.org/ doi/10.1103/PhysRevC.73.014612). Phys. Rev. C 73: 014612. doi:10.1103/PhysRevC.73.014612. .
Island of stability [9] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). "Predictions of alpha decay half lives of heavy and superheavy elements". Nucl. Phys. A 789: 142 – 154. doi:10.1016/j.nuclphysa.2007.04.001. [10] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Search for long lived heaviest nuclei beyond the valley of stability" (http://link. aps.org/doi/10.1103/PhysRevC.77.044603). Phys. Rev. C 77: 044603. doi:10.1103/PhysRevC.77.044603. . [11] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Nuclear half-lives for α -radioactivity of elements with 100 < Z < 130". At. Data & Nucl. Data Tables 94: 781. doi:10.1016/j.adt.2008.01.003. [12] P. Roy Chowdhury, D. N. Basu and C. Samanta (January 26, 2007). "α decay chains from element 113" (http://link.aps.org/doi/10.1103/ PhysRevC.75.047306). Phys. Rev. C 75: 047306. doi:10.1103/PhysRevC.75.047306. . [13] Chhanda Samanta, Devasish Narayan Basu, and Partha Roy Chowdhury (2007). "Quantum tunneling in 277112 and its alpha-decay chain". Journal of the Physical Society of Japan 76: 124201 – 124204. doi:10.1143/JPSJ.76.124201. [14] Sven Gösta Nilsson, Chin Fu Tsang, Adam Sobiczewski, Zdzislaw Szymaski, Slawomir Wycech, Christer Gustafson, Inger-Lena Lamm, Peter Möller and Björn Nilsson (February 14, 1969). "On the nuclear structure and stability of heavy and superheavy elements". Nuclear Physics A 131 (1): 1 – 66. doi:10.1016/0375-9474(69)90809-4. [15] J. Dvorak, W. Brüchle, M. Chelnokov, R. Dressler, Ch. E. Düllmann, K. Eberhardt, V. Gorshkov, E. Jäger, R. Krücken, A. Kuznetsov, Y. Nagame, F. Nebel,1 Z. Novackova, Z. Qin, M. Schädel, B. Schausten, E. Schimpf, A. Semchenkov, P. Thörle, A. Türler, M. Wegrzecki, B. 270 Wierczinski, A. Yakushev, and A. Yeremin (2006). "Doubly Magic Nucleus 108 Hs-162" (http://scitation.aip.org/getabs/servlet/ GetabsServlet?prog=normal&id=PRLTAO000097000024242501000001&idtype=cvips&gifs=yes). Phys. Rev. Lett. 97 (24): 242501. doi:10.1103/PhysRevLett.97.242501. PMID 17280272. . [16] S. Cwiok, P.-H. Heenen and W. Nazarewicz (2005). "Shape coexistence and triaxiality in the superheavy nuclei" (http://www.phys.utk. edu/witek/fission/utk/Papers/natureSHE.pdf) (PDF). Nature 433 (7027): 705. doi:10.1038/nature03336. PMID 15716943. . [17] Zagebraev, V; Greiner, W (2008). "Synthesis of superheavy nuclei: A search for new production reactions" (http://arxiv.org/pdf/0807. 2537v1). Physical Review C 78: 034610. doi:10.1103/PhysRevC.78.034610. . [18] "Greetings From the Island of Stability" (http://www.nytimes.com/2004/02/08/opinion/08SACK.html?ex=1391576400& en=68476e9da837f91f&ei=5007&partner=USERLAND&pagewanted=all), Opinion in the New York Times, February 8, 2004 [19] http://www.newscientist.com/article/mg19926661.200-hunting-the-biggest-atoms-in-the-universe.html?full=true [20] http://arxivblog.com/?p=350 [21] http://159.93.28.88/popeko/e114_287.html [22] http://www.radiochemistry.org/periodictable/elements/115.html [23] http://www.ias.ac.in/currsci/aug10/articles9.htm [24] http://www.cerncourier.com/main/article/39/7/18 [25] http://www.cerncourier.com/main/article/41/8/17 [26] http://imglib.lbl.gov/ImgLib/COLLECTIONS/BERKELEY-LAB/SEABORG-ARCHIVE/index/96B05658.html [27] http://physicsweb.org/articles/world/17/7/7 [28] http://curious.astro.cornell.edu/question.php?number=599 [29] http://www.pbs.org/wgbh/nova/sciencenow/3313/02.html [30] http://www.nytimes.com/2004/02/08/opinion/08SACK.html?ex=1391576400&en=68476e9da837f91f&ei=5007& partner=USERLAND
153
Article Sources and Contributors
Article Sources and Contributors Periodic table Source: http://en.wikipedia.org/w/index.php?oldid=382320968 Contributors: 129.186.19.xxx, 158.252.248.xxx, 1993 lol, 203.109.250.xxx, 64.26.98.xxx, A-giau, A. di M.,
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Brane.Blokar, Burntsauce, Chris Dybala, Claidheamohmor, CryptoDerk, Dandelions, Darnir redhat, Dpvwia, Eddideigel, Edgar181, Eras-mus, Eric119, Femto, Fonzy, French user, Gene Nygaard, Gerard1894, Gurkha711, Hairy Dude, HighKing, Hugo-cs, Ica irns, Inter, Isnow, Jag123, KevinXcore, LukeSurl, Margoz, Michael Hardy, MidnightBlueMan, Nergaal, NuclearWarfare, Obliterator, Ojigiri, Philip Trueman, Puchiko, Res2216firestar, Reyk, R fc1394, Rjwilmsi, Roentgenium111, Sapna3654, Sbyrnes321, Scerri, Scorpion451, Seano12345, Se glea, Silverxxx, Stephenb, Stone, Titus III, Tonyrex, Ver sus22, Vsmith, West Brom 4ever, Woohookitty, Xerxes314, Yekrats, 95 anonymous edits Group XV Source: http://en.wikipedia.org/w/index.php?oldid=380875622 Contributors: 1297, Afisch80, Ahoerstemeier, Ajraddatz, Altermattc, Andrea105, AngelOfSadness, Antonio Lopez,
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EamonnPKeane, Emperorbma, Gentgeen, Höyhens, IvanLanin, J'88, Looxix, Mav, N2e, Numbo3, Pjstewart, Rich Farmbrough, SimonMayer, Smack, Sombrero, Suisui, Tarquin, The Anome, Urod, WFPM, Yankeesrule3, आशीष भटनागर , 19 anonymous edits s-block Source: http://en.wikipedia.org/w/index.php?oldid=380133449 Contributors: Ahoerstemeier, Balcer, Betterusername, Brane.Blokar, Bryan Derksen, Choihei, Dcljr, Donarreiskoffer,
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Xanthine at en.wikipedia Later version(s) were uploaded by McLoaf at en.wikipedia.
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