|Periodic Table of the Elements||He
Transition Metals ( B )
Alphabetically :Actinium (89)
Aluminium (13, UK)
Aluminum (13, US)
Caesium (55, UK)
Cesium (55, US)
Chemical elements are uniquely characterized by their atomic numbers (Z) (now defined as the number of protons in a nucleus of that element).
Before Mendeleev (1834-1907) made that notion clear with an early version of the above table, in 1869, many basic facts puzzled chemists. They could only work with mass numbers obtained by measuring the weight ratios of low-pressure gases, using Avogadro's law (1811). Mass-numbers can be useful to identify elements but they don't provide reliable clues to guess the relations between them. They are only approximately integers and they do not always increase with Z...
For example, Scheele (1742-1786) had discovered, in 1774, that muriatic acid was composed of hydrogen and something else, which he guessed to be an oxide of some new element with mass number 19.45, tentatively dubbed "muriaticum" (the prejudice of that era, formulated by Lavoisier in 1777, was that all acids should contain oxygen).
Disproving Lavoisier's preconceptions and Scheele's guess, Humphry Davy (1778-1829) showed, in 1810, that "muriaticum" doesn't exist. The "oxide of muriaticum" discovered by Scheele was actually a new element (Cl = Chlorine, of mass 35.45).
The reason why mass-numbers are close to integers for many important elements but not for some others, including chlorine, wasn't understood until the notion of isotope emerged, in 1913.
The great confusion of a bygone era is entirely resolved by the periodic table of elements whose structure is worth comitting to memory:
For the Lanthanides (elements 57-71) see
Martyn Poliakoff's video
An easier American mnemonic for the Lanthanides (elements 57-71) is:
The above grouping by column is supplemented by the important distinctions listed below.
Early textbooks (before the 1950's) were excluding groups 11 and 12 from the transition metals. Elements of group 11 are now universally recognized as transition metals, so are elements of group 12 under the simple convention adopted here, following many modern authors who view d-block and transition metals as strictly synonymous. The current definition from the IUPAC is controversial; it would classify Zinc and Cadmium as post-transition metals while Mercury (and Copernicium) should be considered transition metals, because of the recent (2007) synthesis of mercury tetrafluoride (introducing a new oxidation state for Mercury that has been given a relativistic explanation which doesn't apply to Zinc or Cadmium).
Actinides are normally classified as rare earths because of their obvious chemical similarities with lanthanides, without the endorsement of the IUPAC (this issue is relatively unimportant, because of the lack of chemical uses of actinides outside of nuclear engineering).
As the above table demonstrates, the franctic search for elements 85 and 87 (before WWII) once left a few nomenclature debris. It also left at least one textbook example of pathological science (perceived observations at the threshold of detectability that turn into pseudoscience). The so-called (imaginary) Allisson effect was advocated by Fred C. Allison (1882-1974) well beyond the call of scientific duty, even after his method was disproved by H.G. MacPherson (of UC Berkeley) in 1934... A case worth studying.
The current official procedures for enacting the names of new elements were adopted well after the settlement (1997) of a major naming controversy (the Transfermium War) about elements 104, 105 and 106.
In 1997, element 109 (Meitnerium, Mt) was named after Lise Meitner (1878-1968). The rare honor was widely perceived has some kind of posthumus apology for not having shared the Nobel prize (Chemistry 1944) awarded to Otto Hahn for their joint work. This was also a way to honor the Hahn/Meitner team, as the name Hahnium (for element 105) had to be permanently dropped to avoid further confusion after the naming war.
On 19 February 2010 (the 537 th anniversary of Copernicus' birth) element 112 was officially given the name that had been under review since July 2009: Copernicium (Cn). Its temporary name in the IUPAC system was Ununbium (Uub). Alternately, it can be identified as eka-mercury (or eka-hydrargyrum, eka-Hg) the same way element 111 (Roentgenium, Rg) was formerly known as eka-gold. Mendeleev himself introduced the prefix "eka-" to name any undiscovered element after whatever appears above it in the periodic table (such elements are chemically similar).
On 12 August 2012, a Japanese team at RIKEN (Rikagaku Kenkyujo = Institute for physical and chemical research) has published their observation of a decay chain of an atom of Uut-278, including the well-known alpha-decay of Db-262 into Lr-258 which clearly identifies element 113 as the source of that decay chain. This is construed as a definite discovery of 113, for which more ambiguous results had been obtained at RIKEN in 2004 and 2005 (and also at Livermore and Dubna between 2003 and 2005). It's widely expected that the 2012 results will give naming rights to the Japanese, who had mentioned four possible names for element 113 (Japonium, Rikenium or Nishinarium). Since March 2016, most bets seem to be on Nipponium (symbol Np).
On 2015-12-30, the IUPAC has announced that they have completed their final review of the discovery claims for elements 113, 115, 117 and 118 and found them to be satisfactory. The IUPAC is now inviting the respective teams of discovers to propose names and symbols for the new elements. The scientific community at large will then be given a chance to offer comments before the names are finalized before the end of 2016. This milestone is heralded at the completion of the periodic table (or, at least, its first seven lines, up to and including the soon-to-be-properly-named element 118). If nothing else, that's important for typographical and aesthetic reasons!
On 2016-06-08, the next-to-last step took the form of a press release by the IUPAC announcing the four proposed names and the opening of a public discussion about them, scheduled to end in November 2016.
The quantum state of an electron around a nucleus is fully described in terms of the following four quantum numbers :
The Pauli Exclusion Principle states that two electrons cannot be in the same quantum state. They must differ in at least one of the values of the above 4 quantum numbers. This implies that a subshell (n,l) may contain no more than 2(2l+1) electrons, as tabulated above (the total number within the whole shell is at most 2n2 ).
The minimal energy of the electronic cloud surrounding a lone nucleus is achieved when electrons occupy available subshell room in the order at left, starting with the 1s subshell. This simplified version of the Aufbau principle explains the structure of the periodic table of elements, where elements with similar chemical properties are listed in the same column: The chemical properties of an element depend mostly on the valence electrons located in the outermost subshell(s) which are usually the least favored energetically (with the reservations noted below, in the case of "f" subshells).
The electronic configuration around a nucleus may be summarized by listing all nonempty subshells in the above order of increasing energies (1s, 2s, 2p, 3s, etc.) with a superscript indicating the number of electrons in each. The repartition of electrons into orbitals of the same subshell is usually ignored.
A complementary term symbol is sometimes added to better describe the ground configuration. It may be obtained using Hund's Rule, a set of empirical recipes due to Friedrich Hund (1896-1997). Electrons avoid pairing up on the same orbital unless all the orbitals of the subshell are occupied.
For brevity, the configuration of a noble gas may be denoted by its bracketed symbol [as a prefix] in the electronic configuration of subsequent elements. Note that all subshells of noble gases are full. Chemical inertness is due to an outter shell containing a total of 8 electrons (except for helium).
In the periodic table, successive "transition metals" correspond to the "filling" of a "d" subshell (from 1 to 10 electrons). Adding 1 to 14 electrons to the empty "f" subshell of Lanthanum yields the other elements of the Lanthanide series (Z = 58 to 71) whose chemical similarity with Lanthanum may be explained by stating that the "f" subshell corresponds to orbitals that are "closer" to the nucleus than those of the previous "s" subshell, so "f" electrons are less likely fo be valence electrons (the same situation repeats with Actinium and the Actinides series, from Z = 89 to 103). This geometrical explanation should not be taken too literally...
For completeness, it should be noted that the energy levels of some subshells are so close that the pairing of electrons may lead to a few exceptions (in particular for Cr and Cu) in the application of the simplified Aufbau principle presented above.
Theoretically, the chemistry of even highly radioactive heavy elements can be precisely computed from the well-known laws of quantum electrodynamics (as if radioactivity didn't exist). In practice, such a computation is way beyond our current abilities and chemistry remains an experimental science... Nevertheless, the ingenuity of experimentalists is considerable and some chemical facts have been obtained for an element like Hassium (Z = 108) although only about 40 atoms of it have ever been observed. This is made possible by the fact that Hassium happens to have a surprisingly stable isotope (Hs-277) with a half-life of more than 10 minutes, as explained in an excellent video interview of Martyn Poliakoff (part of The Periodic Table of Videos, produced by Brady Haran).