History of the periodic table

The periodic table is an arrangement of the chemical elements, which are organized on the basis of their atomic numbers, electron configurations and recurring chemical properties. Elements are presented in order of increasing atomic number. The standard form of the table consists of a grid with rows called periods and columns called groups.

The history of the periodic table reflects over two centuries of growth in the understanding of chemical properties, with major contributions made by Antoine-Laurent de Lavoisier, Johann Wolfgang Döbereiner, John Newlands, Julius Lothar Meyer, Dmitri Mendeleev, and Glenn T. Seaborg.[1][2]

Early history

A number of physical elements (such as platinum, mercury, tin and zinc) have been known from antiquity, as they are found in their native form and are relatively simple to mine with primitive tools.[3] Around 330 BCE, the Greek philosopher Aristotle proposed that everything is made up of a mixture of one or more roots, an idea that had originally been suggested by the Sicilian philosopher Empedocles. The four roots, which were later renamed as elements by Plato, were earth, water, air and fire. Similar ideas about these four elements also existed in other ancient traditions, such as Indian philosophy.

First categorizations

The history of the periodic table is also a history of the discovery of the chemical elements. The first person in history to discover a new element was Hennig Brand, a bankrupt German merchant. Brand tried to discover the Philosopher's Stone—a mythical object that was supposed to turn inexpensive base metals into gold. In 1669 (or later), his experiments with distilled human urine resulted in the production of a glowing white substance, which he called "cold fire" (kaltes Feuer).[4] He kept his discovery secret until 1680, when Robert Boyle rediscovered phosphorus and published his findings. The discovery of phosphorus helped to raise the question of what it meant for a substance to be an element.

In 1661, Boyle defined an element as "those primitive and simple Bodies of which the mixt ones are said to be composed, and into which they are ultimately resolved."[5]

Lavoisier's Traité Élémentaire de Chimie (Elementary Treatise of Chemistry), which was written in 1789 and first translated into English by the writer Robert Kerr, is considered to be the first modern textbook about chemistry. Lavoisier defined an element as a substance that cannot be broken down into a simpler substance by a chemical reaction.[6] This simple definition served for a century and lasted until the discovery of subatomic particles. Lavoisier's book contained a list of "simple substances" that Lavoisier believed could not be broken down further, which included oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc and sulfur, 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. He classified these substances into metals and non metals. While many leading chemists refused to believe Lavoisier's new revelations, the Elementary Treatise was written well enough to convince the younger generation. However, Lavoisier's descriptions of his elements lack completeness, as he only classified them as metals and non-metals.

In 1815, the English physician and chemist William Prout noticed that atomic weights seemed to be multiples of that of hydrogen.[7]

In 1817, Johann Wolfgang Döbereiner, a chemist, began to formulate one of the earliest attempts to classify the elements.[8] 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.[9]

Definition of Triad law:-"Chemically analogous elements arranged in increasing order of their atomic weights formed well marked groups of three called Triads in which the atomic weight of the middle element was found to be generally the arithmetic mean of the atomic weight of the other two elements in the triad.

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

In 1860, a revised list of elements and atomic masses was presented at a conference in Karlsruhe. It helped spur creation of more extensive systems. The first such system emerged in two years.[10]

Comprehensive formalizations

French geologist Alexandre-Emile Béguyer de Chancourtois noticed that the elements, when ordered by their atomic weights, displayed similar properties at regular intervals. In 1862, he devised a three-dimensional chart, named the "telluric helix", after the element tellurium, which fell near the center of his diagram.[11][12] 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. The original paper from Chancourtois in Comptes Rendus de l'Académie des Sciences did not include a chart and used geological rather than chemical terms. In 1863, he extended his work by including a chart and adding ions and compounds.[13]

The next attempt was made in 1864. British chemist John Newlands presented a classification of the sixty-two known elements. Newlands noticed reoccurring trends in physical properties of the elements at reoccurring intervals of multiple of eight in mass number;[14] based on this observation, he produced a classification of these elements into eight groups. Each group displayed a similar progression; Newlands likened these progressions to the progression of notes within a musical scale.[15][16][17][12] Newlands's table did not leave any gaps for possible future elements and in some cases had to put two elements into the same position within the same octave. Newlands's table was ridiculed by some of his contemporaries. The Chemical Society refused to publish his work. The president of the Society William Odling defended Society's decision by saying that such 'theoretical' topics might be controversial;[18] there was even harsher opposition from within the Society, suggesting the elements could have been just as well listed alphabetically.[10] Later that year, Odling suggested a table of his own[19] but failed to get recognition following his role in opposing Newlands's table.[18]

German chemist Lothar Meyer also noted the sequences of similar chemical and physical properties repeated at periodic intervals. According to him, if the atomic weights were plotted as ordinates and the atomic volumes as abscissas—the curve obtained a series of maximums and minimums—the most electropositive elements appearing at the peaks of the curve in the order of their atomic weights. In 1864, a book of his was published; it contained an early version of the periodic table containing 28 elements, classified elements into six families by their valence—for the first time, elements had been grouped according to their valence. Works on organizing the elements by atomic weight until then had been stymied by inaccurate measurements of the atomic weights.[20] In 1868, he revised his table, but this draft table was published only after his death. In a paper dated December 1869 which appeared early in 1870, Meyer published a new periodic table of 55 elements, in which the series of periods are properly ended by an element of the alkaline earth metal group. The paper also included a line chart of relative atomic volumes, which illustrated periodic relationships of physical characteristics of the elements, and which assisted Meyer in deciding where elements should appear in his periodic table. By this time he had already seen the publication of Mendeleev's first periodic table, but his work appears to have been largely independent.[3]

Russian chemist Dmitri Mendeleev arranged the elements by atomic mass, corresponding to relative molar mass. It is sometimes said that he played "chemical solitaire" on long train journeys, using cards with various facts about the known elements.[21] Another possibility is that he was inspired in part by the periodicity of the Sanskrit alphabet, which was pointed out to him by his friend an linguist Otto von Böhtlingk.[22] Mendeleev used the trends he saw to suggest that atomic weights of some elements were incorrect and accordingly changed their placing: for instance, he figured there was no place for a trivalent uranium with the mass of 120 in his work, and he doubled both the atomic weight and valency of uranium, suggesting it was a hexavalent element with the atomic weight of 240. Mendeleev also figured some spots in his ordering had no element to match, and he left gaps on account of future discoveries of these elements, using the elements before and after those missing ones to predict their properties. In 1869, he finalized his first work and had it published.[23][24] Mendeleev also sent it to a number of well-known chemists, including Meyer; this preceded Meyer's first comprehensive periodic table which he published a few months later, acknowledging Mendeleev's priority.[25] Mendeleev continued to improve his ordering; in 1870, it gained a tabular shape, and in 1871, it was titled "periodic table". Some changes also occurred with new revisions, with some elements changing positions.

The first of Mendeleev's predictions was confirmed in 1875, when gallium was discovered; its properties were close to Mendeleev's predictions for what he termed eka-aluminium. Two more of his predictions were confirmed within another decade.[12] Mendeleev was even able to correct some initial measurements with his predictions.[26] Later chemists used this to justify Mendeleev's table.[10]


In 1882, both Meyer and Mendeleev received the Davy Medal from the Royal Society in recognition of their work on the periodic law.

The importance of Newlands' analysis was eventually recognized by the Chemistry Society with a Gold Medal five years after they recognized Mendeleev's work. It was not until the following century, with Gilbert N. Lewis's valence bond theory (1916) and Irving Langmuir's octet theory of chemical bonding (1919), that the importance of the periodicity of eight would be accepted.[27][28][29]

Pre-atomic theory developments

English chemist Henry Cavendish, the discoverer of hydrogen in 1766, discovered that air is composed of more gases than nitrogen and oxygen.[30] He recorded these findings in 1784 and 1785; among them, he found a then-unidentified gas less reactive than nitrogen.[31] Although helium was discovered in 1868, its chemistry was not investigated at that time. In 1895, William Ramsay and Lord Rayleigh isolated argon from air and determined that it was a new element. Following this discovery, Ramsay noted that an entire group of gases, the noble gases, was missing from the periodic table. Using fractional distillation to separate air, Ramsay discovered three more noble gases in 1898: neon, krypton, and xenon. Although Mendeleev's table predicted several undiscovered elements, it did not predict the existence of noble gases. Mendeleev added them to the table as Group 0 in 1902, without disturbing the basic concept of the periodic table.[31]

In addition to the predictions of scandium, gallium, and germanium that quickly were realized, Mendeleev's 1871 table left many more spaces for undiscovered elements, though he did not provide detailed predictions of their properties. In total, he predicted eighteen elements, though only half corresponded to elements that were later discovered.[26] He observed that an entire row of his table appeared to be missing between cerium and tantalum, attributing this anomaly to the presumed nature of these elements. This arrangement was used to maintain consistency with the eight-column structure, as well as the placement of the elements from osmium to gold, which were known to be analogs of ruthenium through silver. Nevertheless, he predicted the atomic weights for some of these elements, ranging from 140 for eka-molybdenum to 175 for eka-caesium. While elements with these atomic weights were later discovered, their chemistry did not correspond to the gaps in Mendeleev's table.[32] Several heavier analogs were correctly predicted, such as tri-manganese (rhenium) and dvi-iodine (astatine), though they were later found to occupy the positions of dvi-manganese and eka-iodine respectively. This imprecise use of prefixes was caused by a break in the table's structure—the apparent "dead zone" between cerium and tantalum— itself a consequence of Mendeleev's strict adherence to the eight-column structure.[32]

By 1904, Mendeleev's table rearranged several elements, and included the noble gases along with most other newly discovered elements. It still had the dead zone, and a row zero was added above hydrogen and helium to include coronium and the ether, which were widely believed to be elements at the time.[32] Although the Michelson-Morley Experiment in 1887 cast doubt on the possibility of a luminiferous ether as a space-filling medium, physicists set constraints for its properties.[33] Mendeleev believed it to be a very light gas, with an atomic weight several orders of magnitude smaller than that of hydrogen. He also postulated that it would rarely interact with other elements, similar to the noble gases of his group zero, and instead permeate substances at a velocity of 2,250 kilometres (1,400 mi) per second.[34]

In 1905, Swiss chemist Alfred Werner resolved the dead zone of Mendeleev's table. He determined that the rare earth elements (lanthanides), 13 of which were known, lay within that gap. Although Mendeleev knew of lanthanum, cerium, and erbium, they were previously unaccounted for in the table because the number of lanthanides and their exact order were not known. This was in part a consequence of their similar chemistry and imprecise determination of their atomic masses. Combined with the lack of a known group of similar elements, this rendered the placement of the lanthanides in the periodic table difficult.[35] This discovery led to a restructuring of the table and the first appearance of the 32-column form.[32]

Atomic theory and isotopes

Four radioactive elements were known in 1900: radium, actinium, thorium, and uranium. These radioactive elements (termed "radioelements") were accordingly placed at the bottom of the periodic table, as they were known to have greater atomic weights than stable elements, although their exact order was not known. Researchers believed there were still more radioactive elements yet to be discovered, and during the next decade, the decay chains of thorium and uranium were extensively studied. Many new radioactive substances were found, including the noble gas radon, and their chemical properties were investigated.[12] By 1912, almost 50 different radioactive substances had been found in the decay chains of thorium and uranium. American chemist Bertram Boltwood proposed several decay chains linking these radioelements between uranium and lead. These were thought at the time to be new chemical elements, substantially increasing the number of known "elements" and leading to speculations that their discoveries would undermine the concept of the periodic table.[26] For example, there was not enough room between lead and uranium to accommodate these discoveries, even assuming that some discoveries were duplicates or misidentifications. It was also believed that radioactive decay violated one of the central principles of the periodic table, namely that chemical elements could not undergo transmutations and always had unique identities.[12]

Frederick Soddy and Kazimierz Fajans found in 1913 that although these substances emitted different radiation,[36] many of these substances were identical in their chemical characteristics, so shared the same place on the periodic table.[37][38] They became known as isotopes, from the Greek isos topos ("same place").[12][39] Austrian chemist Friedrich Paneth cited a difference between "real elements" (elements) and "simple substances" (isotopes), also determining that the existence of different isotopes was mostly irrelevant in determining chemical properties.[26]

Following Charles Glover Barkla's discovery of characteristic X-rays emitted from metals in 1906, English physicist Henry Moseley considered a possible correlation between x-ray emissions and physical properties of elements. Moseley, along with Charles Galton Darwin, Niels Bohr, and George de Hevesy, proposed that the atomic mass (A) or nuclear charge (Z) may be mathematically related to physical properties.[40] The significance of these atomic properties was determined in the Geiger-Marsden experiment, in which the atomic nucleus and its charge were discovered.[41]

In 1913, amateur Dutch physicist Antonius van den Broek was the first to propose that the atomic number (nuclear charge) determined the placement of elements in the periodic table. He correctly determined the atomic number of all elements up to atomic number 50 (tin), though made several errors with heavier elements. However, Broek did not have any method to experimentally verify the atomic numbers of elements; thus, they were still believed to be a consequence of atomic weight, which remained in use in ordering elements.[40]

Moseley was determined to test Broek's hypothesis.[40] After a year of investigation of the Fraunhofer lines of various elements, he found a relationship between the X-ray wavelength of an element and its atomic number.[42] With this, Moseley obtained the first accurate measurements of atomic numbers and determined an absolute sequence to the elements, allowing him to restructure the periodic table. Moseley's research immediately resolved discrepancies between atomic weight and chemical properties, where sequencing strictly by atomic weight would result in groups with inconsistent chemical peroperties. For example, his measurements of X-ray wavelengths enabled him to correctly place argon (Z = 18) before potassium (Z = 19), cobalt (Z = 27) before nickel (Z = 28), as well as tellurium (Z = 52) before iodine (Z = 53), in line with periodic trends. The determination of atomic numbers also clarified the order of chemically similar rare earth elements; it was also used to confirm that Georges Urbain's claimed discovery of a new rare earth element (celtium) was invalid, earning Moseley acclamation for this technique.[40]

Swedish physicist Karl Siegbahn continued Moseley's work for elements heavier than gold (Z = 79), and found that the heaviest known element at the time, uranium, had atomic number 92. In determining the largest identified atomic number, gaps in the atomic number sequence were conclusively determined where an atomic number had no known corresponding element; the gaps occurred at atomic numbers 43, 61, 72, 75, 85, and 87.[40]

In the 1910s and 1920s, pioneering research into quantum mechanics led to new developments in atomic theory and small changes to the periodic table. The Bohr model was developed during this time, and championed the idea of electron configurations that determine chemical properties. Bohr proposed that elements in the same group behaved similarly because they have similar electron configurations, and that noble gases had filled valence shells;[43] this forms the basis of the modern octet rule. This research then led Austrian physicist Wolfgang Pauli to investigate the length of periods in the periodic table in 1924. Mendeleev asserted that there was a fixed periodicity of eight, and expected a mathematical correlation between atomic number and chemical properties;[44] Pauli demonstrated that this was not the case. Instead, the Pauli exclusion principle was developed. This states that no electrons can coexist in the same quantum state, and showed, in conjunction with empirical observations, the existence of four quantum numbers and its consequence on the order of shell filling.[43] This determines the order in which electron shells are filled and explains periodicity of the periodic table.

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. He explained the chemical properties of transition elements as a consequence of the filling of an inner subshell rather than the valence shell. This proposition, based upon the work of American chemist Gilbert N. Lewis, suggested the appearance of the d subshell in period 4 and the f subshell in period 6, lengthening the periods from 8 to 18 and then 18 to 32 elements.[45]

Looking for expansions and the end of the periodic table

As early as 1913, Bohr's research on electronic structure led physicists such as Johannes Rydberg to extrapolate the properties of undiscovered elements heavier than uranium. Many agreed that the next noble gas after radon would most likely have the atomic number 118, from which it followed that the transition series in the seventh period should resemble those in the sixth. Although it was thought that these transition series would include a series analogous to the rare earth elements, characterized by filling of the 5f shell, it was unknown where this series began. Predictions ranged from atomic number 90 (thorium) to 99, many proposing a beginning beyond the known elements at or beyond atomic number 93. The elements from actinium to uranium were instead believed to form a fourth series of transition metals because of their high oxidation states; accordingly, they were placed in groups 3 through 6.[46]

In 1940, neptunium and plutonium were the first transuranic elements to be discovered; they were placed in sequence beneath rhenium and osmium, respectively. However, preliminary investigations of their chemistry suggested a greater similarity to uranium than to lighter transition metals, challenging their placement in the periodic table.[47] During his Manhattan Project research in 1943, Glenn T. Seaborg experienced unexpected difficulties in isolating the elements americium and curium, as they were believed to be part of a fourth series of transition metals. Seaborg wondered if these elements belonged to a different series, which would explain why their chemical properties, in particular the instability of higher oxidation states, were different from predictions.[47] In 1945, against the advice of colleagues, he proposed a significant change to Mendeleev's table: the actinide series.[48]

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 and included in the periodic table. The actinide series is the second row of the f-block (5f series), and comprises the elements from actinium to lawrencium. In both the actinide and lanthanide series, an inner electron shell is being filled.

Seaborg's subsequent elaborations of the actinide concept theorized a series of superheavy elements in a transactinide series comprising elements from 104 to 121 and a superactinide series of elements from 122 to 153.[47] He proposed an extended periodic table with an additional period of 50 elements (thus reaching element 168); this eighth period was derived from an extrapolation of the Aufbau principle and placed elements 121 to 138 in a g-block, in which a new g subshell would be filled.[49] Seaborg's model, however, did not take into account relativistic effects resulting from high atomic number and electron orbital speed. Burkhard Fricke in 1971[50] and Pekka Pyykkö in 2010[51] used computer modeling to calculate the positions of elements up to Z = 172, and found that the positions of several elements were different from those predicted by Seaborg. Although models from Pyykkö, Fricke, and Nefedov et al.[52] generally place element 172 as the next noble gas, there is no clear consensus on the electron configurations of elements beyond 120 and thus their placement in an extended periodic table. It is now thought that because of relativistic effects, such an extension will feature elements that break the periodicity in known elements, thus posing another hurdle to future periodic table constructs.[51]

The discovery of tennessine in 2010 filled the last remaining gap in the seventh period. Any newly discovered elements will thus be placed in an eighth period.

Despite the completion of the seventh period, experimental chemistry of some transactinides has been shown to be inconsistent with the periodic law. In the 1990s, Ken Czerwinski at University of California, Berkeley observed similarities between rutherfordium and plutonium and dubnium and protactinium, rather than a clear continuation of periodicity in groups 4 and 5. More recent experiments on copernicium and flerovium have yielded inconsistent results, some of which suggest that these elements behave more like the noble gas radon rather than mercury and lead, their respective congeners. As such, the chemistry of many superheavy elements has yet to be well-characterized, and it remains unclear whether the periodic law can still be used to extrapolate the properties of undiscovered elements.[2][53]

See also


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