The Periodic Table: A Historical Survey

Transcription

1 The Periodic Table: A Historical Survey Jack Fergusson Formerly Reader in Chemistry, Department of Chemistry, University of Canterbury ( jb.fergusson@xtra.co.nz) Probably more than anything else the periodic table symbolizes the chemist. Its development began with the increase in the number of elements discovered. By the mid-eighteenth century 15 elements were known, and by the time Antoine L. Lavoisier ( ) published his textbook 25 elements were known to chemists. Lavoisier published a list of 33 elements in 1789 (Lavoisier, 1790, pp ), of which 21 were true elements. By the early part of the nineteenth-century the number of elements had increased dramatically. Jöns J. Berzelius s 1827 table of atomic weights contained 48 elements (Spronsen, 1969, p ). There were, however, some misgivings among some scientists about the increasing complexity and number of elements (Crosland, 1971, p.186). For example, Humphry Davy ( ) considered the number of true elements was small, yet ironically he discovered six (Hudson, 1992, p.97). Most chemists accepted the increase, and by 1860 multiplicity of the elements was recognized (Hudson, 1992, p.125). It was almost axiomatic that chemists would then look for some systematic way of classifying the elements. The result, the Periodic Table, which was developing in the 1860s, has been described by Dennis Rouvray as one of the most important discoveries ever made ; however, it was said to have aroused unrelenting and sometimes biting criticism (Rouvray, 2004, p.1). Eric Scerri called it a big idea in chemistry (Scerri, 2007, p.xiii). Atomic Weights In addition to the discovery of more and more elements there were two essentials required before classification was possible: firstly, some knowledge of their chemical properties; and secondly, some numerical ordering, which in the nineteenth century was knowledge of the weights of the elements. By the end of the eighteenth century chemists were familiar with the concept of equivalent weights for compounds, that is the weight of a reagent that reacts with a known weight of another reagent. This particularly applied to the neutralization of acids or bases. J.B. Richter ( ) during discovered the constancy of the neutralization of a base (or acid) and in 1802 E.G. Fischer ( ) published a list of equivalent weights of acids using a value of 1000 for sulphuric acid as his reference point (Spronsen, 1969, p.43; Scerri, 2007, pp.31-32). At the same time, chemists were making more use of the chemical balance. M Klaproth ( ), for example, did not ignore the discrepancies that appeared when chemical analytical results did not add up to 100%, and because of this he discovered the elements uranium and zirconium (as their oxides) (Partington, 1962, pp ; Hudson, 1992, pp.77-78). The concept of the weight of elements was embodied in John Dalton s ( ) atomic theory (1807-8) where he postulated that atoms of the same element were identical and had a particular weight, and that atoms of different elements had different weights. The first list of relative weights of ultimate particles was recorded by Dalton in 1803 on page 247 of his notebook, and later the same year in a paper to the Manchester Philosophical Society. The list included six elements H, 1; C, 4.3; N, 4.2; O, 5.5; P, 7.2; and S, 14.4 (Dalton et al. 1899, p.26; Partington, 1962, pp , p.789; Leicester & Klickstein, 1952, p.221; Spronsen, 1969, pp.45-48). The 1803 paper was on the absorption of gases by water and other liquids). By 1810 Dalton had increased his list to 32 elements but at least 13 of the atomic weights were in doubt (Spronsen, 1969, pp.45-48). Obtaining a list of relative atomic weights was an achievement, even though some values were wrong because of a lack of knowledge of chemical composition or valency or oxidation state. Dalton assumed if two elements (A and B) only formed one compound its composition would be AB, unless some factor suggested otherwise; if they formed two compounds their compositions would be AB and either AB2 or A2B (Dalton et al. 1899, p.30; Dalton, 1971, p ; Leicester & Klickstein, 1952, p.217; Brock, 1992, p.317). Based on this Dalton et al. (1899, p.30) wrote: From the application of these rules, to the chemical facts already well ascertained we deduce the following conclusions; 1 st. That water is a binary compound of hydrogen and oxygen, [HO] and the relative weights of the two elementary atoms are as 1:7, nearly; 2 nd. That ammonia is a binary compound of hydrogen and azote, [NH] and the relative weights of the two atoms are as 1:5, nearly In the years that followed the number of atomic weights determined increased rapidly, keeping pace with the discovery of new elements. The main person involved was Berzelius ( ) who published 2

2 Chemistry Education in New Zealand May 2010 five lists of atomic weights between 1815 and 1845 (Partington, 1964, pp ; Spronsen, 1969, pp.46-50; Hudson, 1992, pp.86-89). His 1845 list contained 51 elements of which 41 atomic weights were correct or close to correct (Spronsen, 1969, pp.47-49). The values for the alkali metals were incorrect because it was assumed their compounds with oxygen had the stoichiometry MO (Partington, 1948, p.206). Berzelius made use of two principles in obtaining his atomic weights: the law of specific heats enunciated by P.L. Dulong ( ) and A.T. Petit ( ) (Leicester & Klickstein, 1952, pp ; Partington, 1964, pp ), and the law of isomorphism postulated by E. Mitscherlich ( ) (Leicester & Klickstein, 1952, pp ; Partington, 1964, pp ).the law of specific heats, now known to be an approximation, was that the product of the atomic weights and specific heats of elements was a constant. Mitscherlich s law stated that when two compounds are isomorphous (e.g., Na 2 HPO 4 12H 2 O and Na 2 HAsO 4 12H 2 O) they crystallize in similar forms and therefore they usually have similar chemical formulae. Hence, if the number of atoms of one element in one of the compounds is known (in the example above, one P atom) then the number of atoms of the analogous element in the second compound can be inferred (that is one As atom). For gases only, the law of combining volumes, namely that gaseous reactants combined in simple ratios of volumes, proposed by J.L. Gay-Lussac ( ) in 1809 (Dalton et al., 1893, pp.8-24; Leicester & Klickstein, 1952, pp ) also provided a way of obtaining atomic weights. In 1811 an Italian chemist Amedeo Avogadro ( ) proposed a far-reaching principle that equal volumes of all gases at the same temperature and pressure contained the same number of particles (atoms or molecules) (Dalton et al., 1893, pp.28-51; Leicester & Klickstein, 1952, pp ; Avogadro, 1971, pp ): The first hypothesis to present itself in this connection, and apparently even the only admissible one, is the supposition that the number of integral molecules in any gases is always the same for equal volumes, or always proportional to the volumes. Setting out from this hypothesis it is apparent that we have the means of determining very easily the relative masses of the molecules of substances obtainable in the gaseous state This important concept languished for 50 years before being considered seriously by chemists. Dalton rejected Gay Lussac s law and Avogadro s principle because he could not see how atoms of different elements, which were of different sizes, could fit in the same space (Spronsen, 1969, pp.48-49). Berzelius made use of Avogadro s principle for some of his calculations, but he and Dalton could not accept that atoms of the same element could bond together, that is gaseous oxygen was O and not O 2 (Spronsen, 1969, pp.49-50; Hudson, 1992, p.88; Laing, 2004, pp ). But in the main Avogadro s principle was either overlooked or ignored. It is considered that partly this was because of the problems in understanding his difficult terminology, (Hudson, 1992, p.86) and also because he could have presented more experimental data. (Partington, 1964, p.217). Also after about 1840 most chemists tended to use equivalent weights because they were experimentally measurable quantities and therefore considered more reliable (the weight would be better called experimental weights; Spronsen, 1969, pp.50-51; Partington, 1964, p.214; Leicester & Klickstein, 1952, p.406). By the mid-nineteenth century the situation was rather confused and F.A. Kekule ( ) suggested a Congress for chemists to sort things out: this occurred at Karlsruhe in 1860 (Spronsen, 1969, pp.42-43; Partington, 1964, pp ; Laing, 2004, pp ). The result was that nothing was sorted out, except that Stanislao Cannizzaro ( ) vigorously promoted Avogadro s principle at the Congress, and left the participants with a paper on the concept (Crosland, 1971, p.222; Hudson, 1992, pp ; Partington, 1964, pp ; Leicester & Klickstein, 1952, pp ; Laing, 2004, pp This action finally led to the acceptance of the principle (it is not unusual in science for new proposals to be ignored and eventually accepted owing to the activity of another person or persons, or even a clearer presentation of the original proposal), which at last put atomic weights on a firm footing. One of the more significant contributions in determining accurate atomic weights was by J.S. Stas ( ), who during established combining ratios accurately through analysis and synthesis of compounds in very pure states (Partington, 1964, p.877). The scene was now set for others to try and find some system in the arrangement of the elements making use of their atomic weights. Early Classifications of the Elements One of the first mentions of a classification was by Lavoisier who, in his text published in 1789, listed elements as either metals or non-metals (Lavoisier, 1790, pp ). In 1812 Davy observed similarity in the chemical properties of the elements and indicated the similarity between platinum and gold, antimony and tellurium, and the similarity between potassium, sodium and barium. He wrote that 3

3 [t]here is likewise a chain of graduations of resemblance which can be traced throughout the whole series of metallic bodies. (Davy, 1812, pp ). The absence of a concept of valency, or oxidation state, meant that vertical relationships (in terms of the modern periodic table) between the elements were first discerned (Spronsen, 1969, pp.53-56). The same valency was usually assumed vertically and even if wrong the trends were clear. Horizontal relationships had to wait until accurate atomic weights were available, that is until around A significant step was taken by Johann W. Döbereiner ( ) in 1817 and 1829 (only six atomic weights were correct in 1817), who discovered the triads (Dobereiner, 1829); Leicester & Klickstein, 1952, pp ), that is a group of three elements whose middle member had an atomic weight which was the arithmetic mean of the other two, for sulphur selenium and tellurium the atomic weight (AW) of Se 1/2(AW of S + AW of Te), i.e. ( )/ (Döbereiner compared the result with the empirically found atomic weight for selenium of 79). Extension of the triads was hampered by the confusion over equivalent and atomic weights. However, L. Gmelin ( ) in 1827, 1843 and 1852 and J.B.A. Dumas ( ) in 1851 and 1859 did extend the concept (Spronsen, 1969, pp.69-71, p.74; Scerri, 2007, pp.44-48). Dumas searched for a homologous series principle (as found in organic chemistry) in inorganic chemistry, and compared the difference in the atomic weights of neighbouring members of different triads (Weeks & Leicester, 1968; Spronsen, 1969, pp.74-75, 85-86; Dobereiner, 1829; Partington, 1948, pp ; Leicester, 1965, pp ; Partington, 1964). Numerous other chemists investigated similar mathematical relations between atomic weights, including M. von Pettenkofer ( ), P. Kremers (b. 1827), J.H. Gladstone ( ), J.P. Cooke ( ), E. Lenssen (b. 1837), W. Odling ( ), and M. Carey Lea ( ) (Spronsen, 1969, pp.63-96). The Periodic Table During the 1860s a number of attempts were made to produce a classification table (Kaji, 2004, pp ). In 1862 Alexandre E.B. de Chancourtois ( ), a mineralogist, proposed a helical arrangement set at 45 around a cylinder, divided into 16 segments (he selected 16 because the atomic weight of oxygen was 16). Elements with similar chemical properties lay on the same vertical line; he called his diagram Vis tellurique as tellurium was located at the centre of the screw. Chancourtois s work was generally ignored, partly because his published work did not contain diagrams of the helix, owing to a problem with the publisher (Sponsen, 1969, pp ; Kaji, 2004, p.93; Rouvray, 2004, pp.19-21; Scerri, 2007, pp.69-71). John A.R. Newlands ( ) in listed the elements in increasing order of atomic weights and observed the similarity in physical and chemical properties at intervals of eight (or multiples of 8) elements. He called the result the law of octaves, and his 1864 table is given in Table 1. Newlands found that he only got a satisfactory result when he used the correct atomic weights of Cannizzaro. The announcement of his work in 1866 was received with ridicule at the English Chemical Society and the society would not publish his paper. However, 23 years later in 1887 Newlands received the Davy Medal from the Royal Society for his work (Weeks & Leicester, 1968; Newlands, 1971, pp ; Leicester, 1965, p.192; Partington, 1964; Rouvray, 2004, pp.21-22; Kaji, 2004, pp.94-95). Scerri outlines in some detail the events and mythology surrounding Newland s law of octaves (Scerri, 2007, pp.78-80). William Odling ( ) in 1864 and in 1865 produced tables of the elements, based on increasing atomic weights, but made exceptions where it seemed chemically reasonable. His table (Table 2) (Spronsen, 1969, pp ; Kaji, 2004, p.95; Rouvray, 2004, pp.23-25) has similarities to that of Mendeleev s (there are numerous English spellings of Mendeleev s name including Mendeleeff, Mendeleyev, and Mendelejeff ) 1869 table. Odling was the first to invert the order of tellurium (129) and iodine (127) and place Te after Se and I after Br. Gustavus D. Hinrichs ( ) in 1867 produced a circular periodic table with the elements in the same group radiating out on spokes from the centre. This was called the natural classification of the elements, and had a biological flavour. Two years later he produced a more compact tabular version (Table 3) (Spronsen, 1969, pp ; Rouvray, 2004, pp.24-25; Kaji, 2004, p.96). Finally in 1868 Dmitri I. Mendeleev ( ) a Russian chemist (Posin, 1948, pp ; Strathern, 2000) and J. Lothar Meyer ( ), a German chemist, independently drew up the basis of the modern periodic table. Meyer produced his first table in 1864 in the first edition of his book Die Modernen Theorien der Chemie, a popular book which was based on the periodic properties of the elements. He based his table on increasing atomic weights and the valency of the elements (see Table 4) (Spronsen, 1969, pp ; Leicester, 1965, pp ). A new table in 1868 was not published until 1872 because of delays in getting the second edition of his 4

4 Chemistry Education in New Zealand May 2010 Table 1: Arrangement of the Elements According to Newlands (1865) H 1* F 8 Cl 15 Co, Ni 22 Br 29 Pd 36 I 42 Li 2 Na 9 K 16 Cu 23 Rb 30 Ag 37 Cs 44 Tl 53 G 3 Mg 10 Ca 17 Zn 25 Sr 31 Cd 38 Bo 4 Al 11 Cr 19 Y 24 Ce, La Ba V Pt, Ir Pb U 40 Ta 46 Th 56 C 5 Si 12 Ti 18 In 26 Zr 32 Sn 39 W 47 Hg 52 N 6 P 13 Mn 20 As 27 O 7 S 14 Fe 21 Se 28 * Ordinal (order) number Di, Mo Ro, Ru Table 2: Arrangement of the Elements According to Odling (1864) 34 Sb 41 Nb 48 Bi Te 45 Au 49 Os 51 Ro 104 Pt 97 Ru 104 Ir 197 Pd Os 199 H 1 Ag 108 Au Zn 65 Cd 112 Hg 200 L 7 Tl 203 G 9 Pb 207 B 11 Al 27.5 U 120 C 12 Si 28 Sn 118 N 14 P 31 As 75 Sb 122 Bi 210 O 16 S 32 Se 79.5 Te 129 F 19 Cl 35.5 Br 80 I 127 Na 23 K 39 Rb 85 Cs 133 Mg 24 Ca 40 Sr 87.5 Ba 137 Ti 50 Zr 89.5 Ta 138 Th Ce 92 Cr 52.5 Mo 96 V 137 Mn 55 W 184 Fe 56 Co 59 Ni 59 Cu 63.5 Table 3: Hinrichs Tabular form of the Periodic Table (1869) Name Group Symbol of the Elements Pantoïds γ H Kaloïds Kζ Li Na Ka Rb - Chalcoïds Xζ - - Ca Sr Ba Cadmoïds Kδ (Be?) Mg Zn Cd Pb Hydrargoïds Yγ Hg Cuproïids Kν - - Cu Ag Au Ferroïds Σν (Be?) Al Cr Mn Fe Ni Co Rh Ir -Ur Molybdoïds Mλ Bo - Mo Wo Titanoïds Tτ C Si Ti Pd Pt? Sn Nicboïds Nβ - - Va Nb Ta Phosphoïds φ N P P As As Sb Bi Sulphoïds θ O S Se Te - Chloroïds χ Fl Cl Br Io - Pantoïds γ H

5 book published. Also in late 1869 and in 1870 Meyer published another table and a plot of atomic volumes of the atoms against atomic weight. The plot clearly demonstrated the periodic relationship between the elements (Spronsen, 1969, pp ). Fig. 1, a plot of atomic radius (using Bragg-Slater radii) (Bragg, 1920; Slater, 1964) against atomic number, displays the periodic trends. The peaks occur for the alkali metals (except for the last high point). Fig. 1 Atomic radii versus atomic number radius (pm) (Bragg-Slater radii) Atomic Li F Na Cl K Br Rb Atomic Number Mendeleev devised the best table. He ordered the elements in terms of their chemical properties and the weights of the atoms of the elements. Mendeleev found that the elements followed each other in an orderly way with increasing atomic weight. The system that he published in 1869 in his textbook Principles of Chemistry is given in Table 5 (Mendeléeff, 1905, pp.xvii-xviii; Mendeléeff, 1899, Mendeléeff, 1905, pp ; Strathern, 2000, pp ; Kaji, 2004, pp ). The table shows a close similarity to one produced in 1869/70 by Meyer (Spronsen, 1969, p.129). There are a number of mistakes in the tables, but considering what was known at the time the results are remarkable. In later years Mendeleev improved his table a number of times, and published these in later editions of his book, which was translated into other languages (Mendeléeff, 1905, vol. 2, pp.1-47). I Cs Tl There are some gaps in Mendeleev s table (marked by a?), for example, between Si and Sn and between Al and Ur (indium). Mendeleev called these missing elements eka-silicon and eka-aluminium (eka is the Sanskrit numeral one). He then went on in 1871 to predict the properties of the elements, which turned out to be germanium and gallium respectively. The information in Table 6 shows how accurately Mendeleev predicted the properties of the then undiscovered germanium (Spronsen, 1969, p.139; Rouvray, 2004, pp.32-34; Gordin, 2004, pp.56-65). Mendeleev had to make a decision about the placement of the element Te (tellurium) as shown by the question mark written beside its atomic weight. It should follow I (J) (iodine) on the basis of atomic weights, but should precede iodine because of its similar chemical properties to selenium. The problem also occurs in two other places in the periodic table at Co/Ni and Ar/K. Mendeleev made his decision correctly on the basis of chemistry and suggested the atomic weight of Te was in error; in fact, this was not the case because of the numerous isotopes of Te that were unknown at the time. Mendeleev enunciated a simple law of periodicity and an extended form of table in 1869: the properties of the elements, as well as the forms and properties of their compounds, are in periodic dependence on, or (expressing ourselves algebraically) form a periodic function of, the atomic weights of the elements. (Mendeléeff, 1905, vol. 2, pp.17-18, 490). Table 7 is based on the tables in the 1905 English edition of his book (Mendeléeff, 1905, vol. 1, p.xvii; Mendeléeff, 1905, vol. 2, p.21). The problem over the three pairs that seemed to be out of place with respect to their atomic weights was resolved after H.G.J. Mosley ( ) discovered in the significance of the atomic number (number of positive charges, that is protons, in the nucleus: 52 for Te and 53 for I) as being Table 4: Arrangement of the Elements According to Meyer (1864) 4-werting* 3-werting 2-werting 1-werting 1-werting 2-werting Li = 7.03 Be = 9.3? C = 12.0 N = O = Fl = Na = Mg = 24.0 Si = 28.5 P = 31.0 S = Cl = 5.46 K = Ca = 40.0 As = 75.0 Se = 78.8 Br = Rb = 85.4 Sr = 87.6 Sn = Sb = Te = J = Cs = Ba = Pb = Bi = Tl = 204? 4-werting 6-werting 4-werting 4-werting 4-werting 2-werting Ti = 48 Mo = 92 Mn = 55.1 Ni = 58.7 Co = 58.7 Zn = 65 Cu = 63.5 Fe = 56.0 Zr = 90 Vd = 137 Ru = Rh = Pd = 106 Cd = Ag = Ta = W = 184 Pt = Ir = Os = Hg = Au = * werting = valency 6

6 Chemistry Education in New Zealand May 2010 Table 5: Mendeleev s Periodic Table of 1869* I II III IV V VI Ti = 50 Zr = 90? = 180 V = 51 Nb = 94 Ta = 192 Cr = 52 Mo = 95 W = 196 Mn = 55 Rh = Pt = Fe = 56 Ru = Ir = 198 Ni = Co = 59 Pd = Os = 199 H = 1 Cu = 63.4 Ag = 108 Hg = 200 Be = 9.4 Mg = 24 Zn = 65.2 Cd = 112 B = 11 Al = 27.4? = 68 Ur = 116 Au = 197? C = 12 Si = 28? = 70 Sn = 118 N = 14 P = 31 As = 75 Sb = 122 Bi = 210 O = 16 S = 32 Se = 79.4 Te = 128? F = 19 Cl = 35.5 Br = 80 J = 127 Li = 7 Na = 23 K = 39 Rb = 85.4 Cs = 133 Tl = 204 Ca = 40 Sr = 87.6 Ba = 137 Pb =207? = 45 Ce = 92?Eu = 56 La = 94?Yt = 60 Di = 95?In = 75.6 Th = 118? * Some symbols are different from those used today Table 6: Mendeleev s Predicted Properties of Germanium and Actual Values Property Predicted property of eka-silicon, Es Property Values for germanium discovered in 1886 Element Atomic weight 72 Atomic weight Specific gravity 5.5 Specific gravity 5.47 (20 ) Molar volume 13 cc Molar volume cc Colour Dirty grey Colour Grey white Specific heat Specific heat Heat in air White EsO 2 Heat in air White GeO 2 Reaction in acid Slight Reaction in acid None with HCl Preparation EsO 2 + Na Preparation GeO 2 + C Preparation K 2 EsF 6 + Na Preparation K 2 GeF 6 + Na Dioxide Dioxide Refractory Dioxide Refractory Specific gravity 4.7 Specific gravity Molar volume 22 cc Molar volume 22 cc Tetrachloride Boiling point <100*C Boiling point 86 C Specific gravity 1.9 Specific gravity Molar volume 113 cc Molar volume cc Tetramethyl derivative Boiling point 160 C Boiling point 160 C Specific gravity 0.96 Specific gravity <1.00 7

7 more fundamental than atomic weights (Mendeléeff, 1905, vol. 1, pp.xvii-xviii; Mendeléeff, 1905, vol. 2, p.21). Also, after the discovery of isotopes it became clear why tellurium had a higher atomic weight than iodine. Tellurium has eight naturally occurring isotopes listed below with their abundances (Weast & Astle, 1981, pp.b ). 120 Te Te 123 Te 124 Te 125 Te 126 Te 128 Te 130 Te 0.1% 2.5% 0.9% 4.6% 7.0% 18.7% 31.7% 34.5% The weighted mean is atomic mass units. For iodine, the next element in the periodic table there is only one isotope 127 I, whose atomic weight is , hence iodine is lighter than naturally occurring tellurium. Therefore, the atomic number the number of protons is the fundamental number and determines the order of the elements in the periodic table, and not the atomic weight as originally used. It is the atomic numbers 52 and 53 that distinguish Te and I. A number of people, after the publication of the Mendeleev table, attempted other ways of portraying the periodicity of the elements, some being quite inventive and others bizarre, and often with the aim of emphasising a certain aspect. Chapter 6 of van Spronsen s book The Periodic System of the Chemical Elements lists a number of attempts (Spronsen, 1969, pp ). Some used an extended table, as A. Werner ( ) in 1905, which is close to the modern table as shown in outline form in Table 8. In the 1930s it was accepted that any elements higher than uranium (transuranic) would be homologues of the third row transition elements, Re, Os, Ir and Pt (Sime, 1996, pp ). This was in spite of Niels Bohr s prediction in 1922 that transuranics would be homologues of the lanthanides (Spronsen, 1969, pp ), as indeed they are. Also, experimental evidence suggested nuclear reaction only changed the number of protons or neutrons slightly (Sime, 1996, p.164). These beliefs led physicists and chemists astray when they were investing the reaction of uranium and thorium with neutrons that produced nuclear fission and not transuranic elements as was thought (Graetzer & Anderson, 1971). A summary of the various attempts at producing a periodic table is shown in Table 9. Modern Basis of the Periodic Table Joseph J. Thomson ( ) (known by his friends as JJ) studied cathode rays from 1883 and showed that they were negatively charged particles, eventually called electrons (Chayut, 1991). Ernest Rutherford s ( ) work on the structure of the nucleus, Bohr s work on electron structure and orbitals, followed by quantum mechanics (L. de Broglie, W. Heisenberg, E. Schrodinger) (Levine, 1974, p.1-15) made clear the significance of the electron for the chemical properties of the elements and the structure of the periodic table. The electron configuration of the elements is the modern basis of the periodic law. The arrangement of electrons in atoms can be written in the shorthand form (nl) x where n is the principal quantum number 1, 2, 3,, and l the orbital quantum number, 0, 1, 2, (n-1), and x the number of electrons in the orbital. The electronic arrangement for the first 10 elements is given below where each element in the list has one more electron than the preceding member. Element Electron Element Electron configuration configuration H 1s 1 He 1s 2 Li 1s 2 2s 1 Be 1s 2 2s 2 B 1s 2 2s 2 2p 1 C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 A modern form of the periodic table is given in Table 10, where the electron configuration of the outermost (valency) orbital is given at the head of each column. A row of elements across the table is called a period, e.g., Li, Be, B, C, N, O, F, Ne is the first short period. The list of the elements down a row is called a group. Elements in the same group have similar chemical properties, which arises from the fact that they have similar outer electron configurations, e.g., for the group headed by nitrogen the outer configuration is ns 2 np 3, where n is the principal quantum number: Element At. Electron configuration No. N 7 1s 2 2s 2 2p 3 P 15 1s 2 2s 2 2p 6 3s 2 3p 3 As 33 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3 Sb 51 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 3 Bi 83 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 6s 2 6p 3 Group 1 elements, the alkali metals, have the ns 1 outer electron configuration, and the group 17 elements, the halogens, have the ns 2 np 5 outer electron configuration. Modern Periodic Table Showing the Elements Valence Electrons The International Union of Pure and Applied Chemistry (IUPAC) suggestion for distinguishing the groups in the table is to number from the left 1 to 18. So the group headed by boron, B, is group 13,

8 Chemistry Education in New Zealand May 2010 Table 7: Periodic Table According to Mendeleev (1905 edition)* Typical elements I II III IV V VI VII H Li Be B C N O F Na Even elements Odd elements I II III IV V VI VII VIII I II III IV V VI VII Mg Al Si P S Cl K Ca Ti V Cr Mn Fe Co Ni Cu Zn Ga As Se Br Rb Sr Yt Zr Nb Mo Ru Rh Pd Ag Cd In Sn Sb Te J Cs Ba La Ce Er Di? Ta W Os Ir Pt Au Hg Tl Pb Bi Th U * An adaptation of the three tables in Mendeleev s book, atomic weights are not included. Table 8: Extended periodic table s-block elements Lanthanides Actinides Transition metals Trans-actinides p-block elements Super actinides Table 9: Summary of Development of the Periodic Table Date Person Country Contribution Number of elements Early groupings of elements 1789 Lavoisier Fr. Metals/nonmetals Döbereiner Ger. Triads of elements , 43 Gmelin Ger. Triads of elements 50, Pettinkofer Ger. Equivalent weights , 57 Dumas Fr. Triads and bigger groups , 56 Kremers Ger. Mathematical series Gladstone Eng. Triads, similar atomic weights Cooke USA Groups of six Lenssen Ger. Triads of elements Odling Eng. Triads and bigger groups Mercer Eng. Parallels in groups Carey Lea USA Based around At Wt =45 58 Periodic system 1862 de Chancourtois Fr. 3D helical representation Newlands Eng. Law of octaves Odling Eng. Periodic system Hinrichs USA Atomic weight relationships Meyer Ger. Periodic system, atomic volume Mendeleev Rus. Periodic system, chemical knowledge 63 Later developments 1882 Bayley Eng. Long forms of PT Bassett Eng Long forms of PT Thomsen Den. Long forms of PT Werner Swiz. Long forms of PT Mosley Eng. Atomic number 86 9

9 and the group headed by chromium, Cr, is group 6. An earlier numbering scheme used Roman numerals and sub-categories a and b. For example group 1 was Ia and group 11 was Ib, group 15 was Va. The recent suggestion is much simpler and preferred. The actinides range from actinium Ac (atomic number 89) to lawrencium Lw (103). Periodic tables differ in the placement of La and Ac: some, as in Table 10, place them at the beginning of the third and fourth row transition metal series respectively, others place them at the beginning of the lanthanides and actinides respectively. In the latter case, Lu and Lr would be at the beginning of the third and fourth row transition metals respectively (Scerri, 2007, p.21). Elements 104 to 113 have been called transactinides, and elements the super-actinides or super-heavy elements (Krebs, 2006, pp , ; Ede, 2006, pp ; Seaborg, 1963). Some of the elements 104 to 111 are named after people who worked in radiochemistry, such as Rutherford (Rf, 104), Seaborg (Sg, 106), Bohr (Bh, 107), Meitner (Mt, 109), and Roentgen (Rg, 111), others after place names (e.g., Dubnium (Db, 105), was originally called Hahnium, Hn, Dubna is in Russia). The naming of the later elements has been described as a soap opera as scientists and committees have competed for names (Karol, 2004, pp ; Scerri, 2007, pp.8-9). IUPAC have used transitional names for the elements, until an element s discovery is confirmed and a name agreed on. The transitional names make use of the Latin roots for numbers, so 106 (before deciding on seaborgium) was un-nil-hex that is unnilhexium (Unh) (Krebs, 2006, pp ). All of these elements are produced in minute amounts and have very short half-lives, except element 114 (ununquadium, Uuq), which has a half-life of 30 seconds and is the first element of an island of stability. The island of stability concept, suggested by G.T. Seaborg (b. 1912) (Karol, 2004, pp ), arises from the theory of the nuclear configuration of neutrons and protons (nucleons) where the proposed nuclear energy shells are full (not unlike the electron shells of an atom) (Karol, 2004, pp ). The periodic table can be considered as the chemists handbook summarizing the chemistry of the elements, and pointing chemists toward new research ideas. In 1912 Alfred Stock asked the simple question: does boron have similar hydrogen chemistry to its neighbour carbon? (Stock, 1933, Preface). It was found not to be similar but did reveal a new type of compound between boron and hydrogen, the boranes. Also, in the process of handling the lighter members of the boranes, such as B 2 H 6, which are colourless gases and some of which spontaneously flammable, Stock developed the vacuum line, a technique for handling the compounds (Stock, 1933, pp ). In addition, because he and his collaborators handled a lot of mercury in the vacuum line, he also listed the effects of mercury poisoning that they experienced (Stock, 1933, p.203). References Avogadro, A Essay on a manner of determining the relative masses of the elementary molecules of bodies etc. In M. P. Crosland (Ed.), The science of matter: a historical study: selected readings, Harmondsworth: Penguin. Bragg, W.L The arrangement of atoms in crystals, philosophical magazine and journal of science (sixth series) 40, 236, Brock, W.H The fontana history of chemistry, London: Fontana Press. Chayut, M J. J. Thomson: the discovery of the electron and the chemists, Annals of science 48, 533. Crosland, M.P The science of matter: a historical study, selected readings, Harmondsworth: Penguin. Dalton, J A New System of Chemical Philosophy, In M. P. Crosland (Ed.), The science of matter: a historical study: selected readings, Harmondsworth: Penguin. Dalton, J., Gay-Lussac, J-L., & Avogadro, A Foundations of the molecular theory: comprising papers and extracts, Edinburgh: The Alembic Club. Dalton, J., Wollaston, W.H., & Thomson, T Foundations of the atomic theory: comprising papers and extracts, Edinburgh: The Alembic Club. Davy, H Elements of chemical philosophy 1, part 1, pp Quoted in Crosland, 1971, p.193. Dobereiner, J.W An attempt to group elementary substances according to their analogies. Quoted in Crosland, 1971, pp Ede, A The chemical element: a historical perspective, Westport, Conn.: Greenwood Press. Gordin, M.D The short happy life of mendeleev s periodic law. In D.H. Rouvray & R.B. King (Eds), the periodic table: into the 21 st century, Baldock, Herts: Research Studies Press Graetzer, H.G. & Anderson, D.L The discovery of nuclear fission: a documentary history, New York: Van Nostrand Reinhold Company. J. Hudson, J The history of chemistry, London: Chapman and Hall. Kaji, M Discovery of the periodic law: mendeleev and other researchers on element classification in the 1860s. In D.H. Rouvray & R.B. King (Eds), the periodic table: into the 21 st century, Baldock, Herts: Research Studies Press. Karol, P.J The heavy elements. In D.H. Rouvray & 10

10 Chemistry Education in New Zealand May 2010 Table 10: Modern Periodic Table showing the elements valence electrons n ns 1 ns 2 (n-1)d 1-10 s 2 np 1 np 2 np 3 np 4 np 5 np 6 s-block Transition Metals, d-block p-block Group Number st Short Period 2 2 nd Short Period 3 1 st Long Period 4 2 nd Long Period 5 3 rd Long Period 6 7 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 He 2 Be B C N O F Ne Mg 12 Ca 20 Sr 38 Ba 56 Ra 88 Al 13 Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Y 39 La 57 Ac 89 Zr 40 Hf 72 Rf 104 Nb 41 Ta 73 Db 105 Mo 42 W 74 Sg 106 Tc 43 Re 75 Bh 107 Ru 44 Os 76 Hs 108 Rh 45 Ir 77 Mt 109 Pd 46 Pt 78 Ds 110 Ag 47 Au 79 Rg 111 Cd 48 Hg 80 Uub 112 In 49 Tl 81 Si 14 Ge 32 Sn 50 Pb 82 Uuq 114 P 15 As 33 Sb 51 Bi 83 S 16 Se 34 Te 52 Po 84 Cl 17 Br 35 I 53 At 85 Ar 18 Kr 36 Xe 54 Rn 86 4f d 0-1 6s 2 Lanthanides 5f d 0-1 7s 2 Actinides Ce 58 Th 90 Pr 59 Pa 91 Nd 60 U 92 Pm 61 Np 93 Sm 62 Pu 94 Eu 63 Am 95 Gd 64 Cm 96 Tb 65 Bk 97 Dy 66 Cf 98 Ho 67 Es 99 Er 68 Fm 100 Tm 69 Md 101 Yb 70 No 102 Lu 71 Lr

11 R.B. King (Eds), the periodic table: into the 21 st century, Baldock, Herts: Research Studies Press. Krebs, R.E The history and use of our earth s chemical elements, 2 nd ed., Westport, Conn.: Greenwood Press. Laing, M Patterns in the periodic table - old and new. In D.H. Rouvray & R.B. King (Eds), the periodic table: into the 21 st century, Baldock, Herts: Research Studies Press. Lavoisier, Antoine-Laurent Elements of chemistry: in a new systemic order, containing all the modern discoveries, trans. Robert Kerr, 1965, New York: Dover Publications, 1965 (1790). Leicester, H.M The historical background of chemistry, New York: John Wiley & Sons. Leicester, H.M. & Klickstein, H.S A source book in chemistry , Cambridge, Mass.: Harvard University Press. Levine, I.N Quantum chemistry, 2 nd. ed., Boston: Allyn and Bacon. Mendeléeff, D The periodic law of the chemical elements (Faraday lecture), journal of the chemical society, transactions, 55, Mendeléeff, D The principles of chemistry, ed. Thomas A. Pope, trans. George Kamensky, 3 rd English ed., two vols, vol. 1, London: Longmans, Green and Co. Mosley, H.G.J The high-frequency spectrum of the elements, philosophical magazine series 6, 26, 257. Newlands, J Letter to the Editor of Chemical news. In M.P. Crosland (Ed.), The science of matter: a historical survey: selected readings, Harmondsworth: Penguin. Partington, J.R A short history of chemistry, 2 nd ed., London: Macmillan. Partington, J.R A history of chemistry, vol. III, London: Macmillan. Partington, J.R A history of chemistry, vol. IV, London: Macmillan. Posin, D.Q Mendeleyev: the story of a great scientist, New York and Toronto: Whittlesey House; Mc- Graw Hill Book Company. Rouvray, D.H Fact and fable in the story of the table. In D.H. Rouvray & R.B. King (Eds), the periodic table: into the 21 st century, Baldock, Herts: Research Studies Press. Scerri, E.R The periodic table: its story and significance, Oxford: Oxford University Press. Seaborg, G.T Man-made transuranic elements, Englewood Cliffs, N. J.: Prentice Hall. Sime, R.L Lise Meitner: a life in physics, Berkeley: University of California Press. Slater, J.C Atomic radii in crystals, the journal of chemical physics, 41, 10: Spronsen, J.W. van The periodic system of chemical elements: a history of the first hundred years, Amsterdam: Elsevier. Stock, A Hydrides of boron and silicon, Ithaca and New York: Cornell University Press. Strathern, P Mendeleyev s dream: the quest for the elements, Harmondsworth: Penguin. Weast, R.C. & Astle, M.J. (eds.), Crc handbook of chemistry and physics, Boca Raton: CRC Press. Weeks, M.E., & Leicester, H.M Discovery of the elements, 7 th ed., Easton: Journal of Chemical Education. ChemScrapes I m telling you, you re NUTS...they ll never buy it! Brendan Burkett 12