User:Double sharp/Periodic table/Old history section

Early systemisation attempts
The notion of a chemical element was first defined by the Irish chemist Robert Boyle in 1680, as "those primitive and simple Bodies of which the mixt ones are said to be composed, and into which they are ultimately resolved." Another definition was given by French chemist Antoine Lavoisier in 1781 in his textbook Traité Élémentaire de Chimie (Elementary Treatise of Chemistry): a substance whose smallest units cannot be broken down into a simpler substance. Lavoisier's book contained a list of "simple substances" that Lavoisier believed could not be broken down further, which formed the basis for the modern list of elements. Lavoisier's list also included 'light' and 'caloric', which at the time were believed to be material substances.

Starting in 1801, the English chemist John Dalton began to reintroduce the atomic theory into science. Although Lavoisier had ignored the possible existence of atoms, Dalton assumed that each element was in fact composed of its own kind of atom, and that atoms could not transmute into one another, contrary to the old alchemical ideas. Drawing on data of the composition of compounds, he assigned atomic weights to the elements. For example, Lavoisier had found that water is composed of 85% oxygen and 15% hydrogen; he assumed that water would have the formula HO, and thus expected that an oxygen atom weighed 85/15 = 5.5 times the weight of a hydrogen atom. This provided a justification for the law of definite proportions: whenever two elements react to form different compounds, the ratios of the amounts reacting are always ratios of whole numbers. However, some of his atomic weights were wrong because he did not yet have the understanding of valence, as well as the understanding that the molecules of gases such as hydrogen and oxygen were in fact each composed of two atoms. This resulted in a problem of differing atomic weights that was finally resolved by Italian chemist Stanislao Cannizzaro in 1860.

The discovery of atomic weights, and the determination that many of them were close to integer multiples of that of hydrogen, led to the 1815 hypothesis of William Prout that all elements were literally built out of hydrogen; later and more accurate determination of atomic weights were far enough away from integers to discount this hypothesis, although the idea is in some sense similar to the modern idea of a nucleus composed of individual protons and neutrons.

Soon after the notion of a chemical element was defined, attempts were made to organise them. In 1817, German physicist Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups triads. Chlorine, bromine, and iodine formed a triad; as did calcium, strontium, and barium; lithium, sodium, and potassium; and sulfur, selenium, and tellurium. Today, all these triads form part of modern-day groups. German chemist Leopold Gmelin worked with this system, and by 1843 he had identified more triads, as well as some groups of four. Although various chemists were able to identify more and more relationships between small groups of elements, they had yet to build one scheme that encompassed them all.

Mendeleev
The periodic table is generally considered to have been discovered in 1869, by the Russian chemist Dmitri Mendeleev. Although other chemists had found some other versions of the periodic system at about the same time, Mendeleev was the most dedicated to developing and defending his system, and it was his system that most impacted the scientific community. On 17 February 1869, Mendeleev began arranging the elements and comparing them using Cannizzaro's corrected atomic weights. He began with a few elements, and over the course of the day his system grew till it encompassed most of the known elements. After finding a consistent arrangement, his printed table appeared the next month in the journal of the Russian Chemical Society. In some cases, an element appeared to be missing, and Mendeleev boldly predicted that an unknown element would fit in that place, specifically pointing out two entries "? = 68" and "? = 70" in his publication that should be the missing heavier homologues of aluminium and silicon, respectively. Unlike others (such as Julius Lothar Meyer, who had arrived at a periodic system at about the same time), Mendeleev did not use valence as his classification tool (because many elements showed multiple valences); he preferred to consider the elements as basic substances, and hence ordered them by atomic weight. He closed his article with eight points summarising the periodic law:
 * 1) The elements, if arranged according to their atomic weights, exhibit an evident periodicity of properties.
 * 2) Elements which are similar as regards their chemical properties have atomic weights which are either of nearly the same value (e.g., platinum, iridium, osmium), or which increase regularly (e.g., potassium, rubidium, caesium).
 * 3) The arrangement of the elements, or of groups of elements, in the order of their atomic weights corresponds to their so-called valences as well as, to some extent, to their distinctive chemical properties—as is apparent among other series—in that of lithium, beryllium, [boron], carbon, nitrogen, oxygen, and [fluorine].
 * 4) The elements which are 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 should be between 65 and 71.
 * 7) The atomic weight of an element may sometimes be amended by a knowledge of those of contiguous elements. Thus, the atomic weight of tellurium must lie between 123 and 126, and cannot be 128.
 * 8) Certain characteristic properties of the elements can be foretold from their atomic weights.

The next year, Mendeleev relocated several more elements (indium, cerium, thallium, and uranium) on the basis of their atomic weights being likely in error; in every case the updated value accords with modern knowledge. He also added a third predicted element, that he expected would be a missing heavier analogue of boron. In 1871, Mendeleev published a long article, including an updated form of his table, that made his predictions for unknown elements explicit. To give provisional names to his predicted elements, Mendeleev used the prefixes eka-, dvi- or dwi-, and tri-, from the Sanskrit names of digits 1, 2, and 3, depending upon whether the predicted element was one, two, or three places down from the known element of the same group in his table: thus, his first predicted elements were called eka-boron, eka-aluminium, and eka-silicon.

In 1875, the French chemist Paul-Émile Lecoq de Boisbaudran, working without knowledge of Mendeleev's prediction, discovered a new element in a sample of sphalerite, and named it gallium. He isolated the element and began determining its properties. Mendeleev, reading de Boisbaudran's publication, sent a letter claiming that gallium was his predicted eka-aluminium. Although Lecoq de Boisbaudran was initially sceptical (and suspected that Mendeleev was trying to take credit for his discovery), he later admitted that Mendeleev was correct. All the properties de Boisbaudran found matched Mendeleev's predictions, except the density: de Boisbaudran had measured 4.7 g/cm3. Mendeleev suggested that de Boisbaudran remeasure the density, and this time he found a density of 5.9 g/cm3, exactly what Mendeleev had predicted. (The modern value is 5.904 g/cm3).

In 1879, the Swedish chemist Lars Fredrik Nilson discovered a new element, which he named scandium. Like de Boisbaudran before him, he was unaware of Mendeleev's prediction. His compatriot Per Theodor Cleve recognised that scandium matched the properties of eka-boron, and notified Mendeleev. In 1886, the German chemist Clemens Winkler discovered another new element, which he named germanium; his compatriot Hieronymous Theodor Richter recognised it as Mendeleev's eka-silicon. In 1889, Mendeleev noted at the Faraday Lecture to the Royal Institution in London that he had not expected to live long enough "to mention their discovery to the Chemical Society of Great Britain as a confirmation of the exactitude and generality of the periodic law".

In 1894, Lord Rayleigh and William Ramsay discovered the noble gas argon. This had not been foreseen by the periodic law, and it did not seem to fit. Mendeleev at first suggested that argon might not be an element, but rather a triatomic allotrope of nitrogen; but measurements soon ruled out this possibility. The next year, helium was found on Earth, and three years after that, the remaining stable noble gases neon, krypton, and xenon were discovered. (The last natural noble gas, radon, was discovered in 1899.) In 1900, Ramsay suggested to Mendeleev that these noble gases in fact belonged in an eighth main group, between the halogens and the alkali metals. Mendeleev agreed, and considered the result to have been a "magnificent survival" of the periodic law against a "critical test".

A problem remained, however, in the rare earth elements, which at that point were little studied. In the known rare earth elements, atomic weight appeared to increase, but with no concomitant increase in valence. In 1878, the Czech chemist Bohuslav Brauner read Mendeleev's 1871 paper, and attempted to fit the rare earths in the table himself; lanthanum would go in group III, cerium in group IV, and "didymium" (then thought to be an element) in group V. However, didymium would not show a higher oxide of valence 5 that this placement would require. The 1885 discovery of Austrian chemist Carl Auer von Welsbach that didymium was in reality a mixture of two elements, which he called praseodymium and neodymium, did not solve the problem. After much fruitless investigation, in 1902 Brauner abandoned the idea of placing the rare earths among the main groups all other elements belonged in, and suggested his "asteroidal hypothesis". Just as between Mars and Jupiter there was not one planet but a whole asteroid belt, so Brauner felt that between lanthanum and tantalum, all the rare earths would appear, taking up one space in group IV of the periodic table. This hypothesis soon gained ground, although there was some dispute on whether the rare earths belonged all in group IV, all in group III, or perhaps should be spread over groups III and IV or even groups II, III, and IV.

Moseley: atomic numbers
The discovery of X-rays in 1895 (by German physicist Wilhelm Röntgen), and of radioactivity in 1896 (by French physicist Henri Becquerel), made it possible to probe the internal structure of the atom. In 1897, J. J. Thomson discovered the electron. Becquerel had determined that radioactivity was an intrinsic property of uranium, and Polish physicist and chemist Marie Curie noticed that thorium exhibited it as well. She and her husband Pierre Curie then went on to analyse pitchblende, an ore of uranium, and found two more strongly radioactive elements: polonium and radium. The New Zealand physicist Ernest Rutherford also investigated radioactivity, and in 1900–3, he and Frederick Soddy found that radioactive decay involved the transmutation of elements. In 1911, Rutherford discovered that the positive charges in an atom (balancing the negative charges of the electrons) were concentrated in its centre, thus discovering the atomic nucleus.

Amateur Dutch physicist Antonius van den Broek proposed in 1913 that the nuclear charge determined the placement of elements in the periodic table; Rutherford and Soddy praised his contribution, and Rutherford coined the word "atomic number" for this nuclear charge.

In 1913, English physicist Henry Moseley using X-ray spectroscopy confirmed van den Broek's proposal experimentally. Moseley determined the value of the nuclear charge of each element and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge. Nuclear charge is identical to proton count and determines the value of the atomic number (Z) of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley's research immediately resolved discrepancies between atomic weight and chemical properties. Mendeleev had placed tellurium before iodine to fit periodic trends, but tellurium was heavier than iodine, so he was convinced that one of the atomic weights had to be wrong; but the anomaly did not go away with more and more accurate measurements. Moseley clarified the situation by proving that the atomic number of tellurium (52) was lower than that of iodine (53). In particular, the determination of atomic numbers also clarified the order of chemically similar rare-earth elements. The French chemist Georges Urbain, who specialised in the rare earths, challenged Moseley to determine the rare earths present in a sample; he was astonished to receive in one hour the correct answer, which had taken him months to find by chemical means.

Moseley predicted, in 1913, that the only elements still missing between aluminium (Z = 13) and gold (Z = 79) were Z = 43, 61, 72, and 75, all of which were later discovered. The atomic number is now the absolute definition of an element and gives a factual basis for the ordering of the periodic table. After Moseley was killed in World War I, the Swedish physicist Manne Siegbahn continued Moseley's work for elements heavier than gold, and found that the heaviest known element at the time, uranium, had atomic number 92. The missing elements in this range were found to be 85 and 87. The discovery of elements had now ceased to be an open-ended task: up to uranium, it was now known exactly which elements remained to be discovered.

The discovery of radioactivity also came with the discovery of isotopes. Many of the new radioactive elements were found to be chemically identical, posing a problem for the periodic table. Although some (such as van den Broek) tried to incorporate every one of them with its own place, Soddy realised in 1913 that these species had the same nuclear charge and simply different atomic weights, and should occupy the same space in the periodic table. He thus named them "isotopes" after the Greek for "same place". This idea was confirmed in 1914, when the American chemist Theodore W. Richards found that lead exhibits significant variation in atomic weights (as it is the final decay product of thorium and uranium, and lead produced by decay of heavier elements shows different isotopic abundances from primordial lead).

Electron shells and quantum mechanics
Rutherford's discovery of the nucleus posed a new problem: by classical physics, atoms should not be stable, because the negatively charged electrons would spiral into the positively charged nucleus. Moreover, the optical spectra of the elements showed discrete lines, but by classical physics, electrons should be able to take on any energy level, and so the spectra should have been continuous.

These problems were solved by the Danish physicist Niels Bohr, who applied Max Planck's idea of quantisation to the atom. He concluded that the energy levels of electrons were quantised: only a discrete set of stable energy states were allowed. Bohr then attempted to understand periodicity through electron configurations, surmising that the outer electrons should be responsible for the chemical properties of the elements, and that moving to the next element involved adding one more electron. In 1913, he produced the first electronic periodic table. However, he could not yet theoretically justify the electron configurations of the elements, and was forced to use chemical and spectroscopic considerations. The configurations he obtained from these, and again in 1921–3 when he produced an improved version, do not always match the modern ones. Further problems arose, such that some physicists such as Werner Heisenberg and Wolfgang Pauli began to question the notion of electron orbits.

In 1923, Bohr wrote to Pauli for help. Pauli's scheme, with the four quantum numbers, generally resolved the problems. He formulated his exclusion principle, which stated that no two electrons could have the same four quantum numbers. This explained the lengths of the periods in the periodic table (2, 8, 18, and 32), which corresponded to the number of electrons that each shell could occupy. In 1925, Friedrich Hund arrived at configurations close to the modern ones.

British chemist Charles Bury is credited with the first use of the term "transition metal" in 1921 to refer to elements between the main-group elements of groups II and III. Based on the work of American chemists Gilbert N. Lewis and Irving Langmuir, he explained the chemical properties of transition elements as a consequence of the filling of an inner subshell rather than the valence shell: the second-outermost for the transition elements, and the third-outermost for the inner transition elements.

The arguments of Bohr and Bury, as well as the chemical arguments of Austrian chemist Friedrich Paneth, proved of relevance with regard to the discovery of element 72. Urbain claimed to have found element 72 as a rare earth in 1911, but Bohr, Bury, and Paneth believed that element 72 should be a zirconium homologue. Dirk Coster and Georg von Hevesy were thus motivated to search for the new element in zirconium ores, where they found the real element 72 in 1922; they named it hafnium, after the Latin name for Copenhagen, where the element had been discovered (and also Bohr's hometown).

The Aufbau principle, which describes the electron configurations of the elements, was first empirically observed by Erwin Madelung in 1926 and published in 1936. As such, it is often called the Madelung rule. He noted that:
 * 1) electrons fill orbitals in the order of increasing n + ℓ, a sum of two quantum numbers;
 * 2) when two orbitals of the same n + ℓ value are available, electrons preferentially fill that with increasing n.

In 1962 the Russian agricultural chemist Vsevolod Klechkovsky proposed a theoretical explanation for the first part of the Madelung rule, based on the statistical Thomas–Fermi model of the atom. Theoretical justification of the full Madelung rule came in 1971 from Demkov and Ostrovsky, as well as in 1979 from D. Pan Wong.

The first artificial elements and the actinide concept
By 1925, the pool of missing elements from hydrogen to uranium had shrunk still further: the last two stable elements 72 (hafnium) and 75 (rhenium) had been found, so only Z = 43, 61, 85, and 87 remained missing. The element with Z = 43 held a special place; it was Mendeleev's eka-manganese, the fourth of Mendeleev's predictions to which he had assigned an atomic weight (100, later revised to 99). Many chemists thus tried to find it in nature, but none were successful. It eventually became the first element to be synthesised artificially via nuclear reactions rather than discovered in nature; Italian chemists Emilio Segrè and Carlo Perrier analysed a sample of molybdenum (the preceding element) that had been bombarded with deuterons, and isolated from it the new element 43. Initially, they did not name their discovery, as it was felt that artificial elements produced in invisible quantities should not count. However, by 1947, the situation had changed massively, as synthetic long-lived plutonium (element 94) had since been stockpiled; as such, Paneth proposed to end the discrimination between natural and synthetic isotopes. Segrè and Perrier then named their discovery technetium, after the Greek word for "artificial". Of the remaining infra-uranium elements, elements 61 (promethium) and 85 (astatine) were likewise produced artificially; element 87 (francium) became the last element to be discovered in nature, by French chemist Marguerite Perey. It was later determined that technetium, promethium, and astatine also occur in nature, but only in vanishingly small trace quantities from the disintegration of uranium.

In 1934, Italian physicist Enrico Fermi and his team attempted to produce transuranium elements by bombarding uranium with neutrons; they thought that they were detecting transuranium elements, but in 1939, Otto Hahn, Lise Meitner, Fritz Strassmann, and Otto Frisch discovered that the "new elements" were actually fission products: products of the splitting of uranium into two smaller nuclei.

It was generally agreed that a second f-block series should start in the fragmentary seventh period, but not where. Due to their chemistry, it was thought that the elements actinium through uranium formed part of a fourth d-block transition series. As part of the ongoing research on nuclear fission, the American physicist Edwin McMillan once again bombarded uranium with neutrons in 1939, finding a new half-life of 2.3 days. McMillan and Segrè analysed this chemically, working under the assumption that element 93 would be a heavier homologue of rhenium, but did not find any rhenium-like chemistry; they thus assumed that the 2.3-day half-life was another fission product. But as more information about nuclear fission came in, this idea became more remote, and the next year McMillan and Philip Abelson once again analysed the unknown half-life, finding that it was not like any known element and was more like uranium than a rare earth. McMillan and Abelson thus discovered the first transuranium element, which they named neptunium as Neptune is the next planet after Uranus in the Solar System.

McMillan continued working on his new element, as he realised that the decay of element 93 would have to produce the new element 94. He confided his work with his colleague, American chemist Glenn T. Seaborg. McMillan was then sent away to join the military as the United States was planning to join World War II. Seaborg and his colleagues then continued the work, isolating element 94, plutonium, the next year. Because neptunium and plutonium showed chemistry similar to that of uranium, they concluded that the second f-block series started at uranium, creating a "uranide" series analogous to the lanthanides. Seaborg's team then attempted to continue, assuming that elements 95 and 96 would likewise show uranium-like behaviour, but could not identify them. Thus Seaborg suggested in 1944 that there was an actinide series analogous to the lanthanides; later that year his team found elements 95 and 96, and found them to be similar to their lanthanide homologues europium and gadolinium. By analogy, they named the new elements americium and curium. Seaborg's proposed revision was considered bold at the time; renowned inorganic chemists advised him not to publish it, as they feared it would ruin his reputation. He published anyway after the war ended, and the actinide concept enabled the correct identification of all the subsequent elements. Nevertheless, some debate continues about exactly where the f-block elements start and end, even after the existence of a second f-block series has become universally accepted.

Completing the seventh row
It had historically been thought that elements beyond 104 would be impossible, because the electronic repulsion between the protons would overpower the strong nuclear force between them. But in 1949, German physicists Maria Goeppert Mayer and Johannes Hans Daniel Jensen et al. independently derived the nuclear shell model: like the electrons, the nucleons are arranged in shells, and elements corresponding to a closed shell should experience greater stability. This should create an island of stability of superheavy elements. The next closed-shell proton number has been variously estimated to be between 114 and 126, and there is still no consensus on where exactly it is;  nevertheless, the known superheavy elements to 118 provide evidence that this theory is correct, as without the shell closures they would not exist.

The synthesis of new elements proved uneventful until element 101, with all discoveries being made by the team of Seaborg and Albert Ghiorso at the Lawrence Berkeley National Laboratory (LBNL) in Berkeley, California, United States, from 1949 to 1955. However, a significant controversy arose with the succeeding elements 102 through 106 in the 1960s and 1970s, as competition arose with a new team at the Joint Institute for Nuclear Research (JINR) in Dubna in the Soviet Union, led by Georgy Flyorov. Each team claimed discovery, and in some cases each proposed their own name for the element, creating an element naming controversy that lasted decades. IUPAC at first adopted a hands-off approach, preferring to wait and see if a consensus would be forthcoming. Unfortunately, it was also the height of the Cold War, and it became clear after some time that this would not happen. As such, IUPAC decided on a systematic naming scheme in 1979 that would provide neutral temporary names for these elements until it was clear who would be officially awarded the discovery. IUPAC and the International Union of Pure and Applied Physics (IUPAP) then created a Transfermium Working Group (TWG, fermium being element 100) in 1985 to arbitrate the discoveries, which by then stretched to element 109 after the work of scientists at the Gesellschaft für Schwerionenforschung (Society for Heavy Ion Research) in Darmstadt, West Germany.

The TWG held meetings with delegates from the three competing institutes; in 1990, they established criteria on recognition of an element, and in 1991, they finished the work on assessing discoveries and disbanded. These results were published in 1993. LBNL criticised the report, while JINR and GSI endorsed it. In 1994, IUPAC published a recommendation on naming the disputed elements. This recommendation was criticized by the LBNL for several reasons. Firstly, their suggestions were scrambled: the names rutherfordium and hahnium, originally suggested by Berkeley for elements 104 and 105, were respectively reassigned to elements 106 and 108. Secondly, elements 104 and 105 were given names favored by JINR, despite earlier recognition of LBNL as an equal co-discoverer for both of them. Thirdly and most importantly, IUPAC rejected the name seaborgium for element 106, having just approved a rule that an element could not be named after a living person (Seaborg being still alive), even though the 1993 report had given the LBNL team the sole credit for its discovery. After a significant amount of public pressure, both from the United States and elsewhere, IUPAC yielded and suggested a new set of names in 1997, which avoided scrambling suggestions and included seaborgium; this finally was accepted. By this point, the GSI had gone on to discover three new elements after the deliberations of the TWG had concluded: elements 110, 111, and 112.

The final six elements on the periodic table were discovered from 1999 to 2009 at the JINR (now in Russia), now led by Yuri Oganessian; and in the case of element 113, also at the Riken research institute in Wakō, Saitama, Japan. For the elements 110–118 whose discovery postdated the deliberations of the TWG, IUPAC and IUPAP set up new Joint Working Parties (JWP) to evaluate discoveries of new elements according to the TWG's criteria; after priority was assigned, the elements were officially added to the periodic table, and the discoverers were invited to propose their names. By 2016, this had occurred for all 118 known elements, therefore completing the periodic table's first seven rows. This was unprecedented in the periodic table's history.

Most recently, there have been complaints about the scientific quality of the JWP reports, along with the roles that IUPAC and IUPAP should play in the recognition of elements. This stems from the discovery of new elements having shifted from being a matter of chemistry to being a matter of nuclear physics; all the elements from nobelium (element 102) onward were identified physically through their radioactive decay, rather than by their chemistry. In response to this, some new rules have been instituted for the discoveries of elements beyond 118, that would begin an eighth row on the periodic table.

In celebration of the periodic table's 150th anniversary, the United Nations declared the year 2019 as the International Year of the Periodic Table, celebrating "one of the most significant achievements in science". Today, the periodic table is among the most recognisable icons of chemistry.

Chronology of element discovery
The elements with no date listed for discovery were known either in antiquity or to the alchemists.