Frederick C. Strong Ill
Stevens Institute of Technology Hoboken, New Jersey
I
The Atomic Form Periodic Tdde
I
T o an outsider, it must appear that the favorite parlor game of many chemists is the devising of new forms of the periodic table of the elements. Being on the inside, we realize that many of the novel forms proposed are based on some different guiding principle that the author feels is important. A certain amount of dissatisfaction with the conventional short and long forms is usually the reason why a chemist enters the game. This author finds two sources of dissatisfaction. One is the inability of conventional tables to cope geometrically with the increasing length of the periods, which necessitates placing the lanthanide and actinide series in footnotes to the table. This was not as serious in the pre-plutonium era when the rare earth elements were rare and the actinides not recognized as a similar series. With pure lauthanidss becoming increasingly available, and the new actinides assuming world-shaking proportions, footnotes hardly seem to be appropriate positions for these two important series. The other inadequacy of conventional tables is their failure to make sufficientlyclear to students the relationship between atomic structure and physical and chemical properties. While the connection is there, i t is not readily apparent. The reason for this is quite obvious historically, since the conventional tables originated with Mendeleev in 54 B.Q.M. (Before Quantum Mechanics). While he put the chemists far ahead of the physicists in understanding the atom, it must be conceded that the physicists have caught up. As yet, the two periodic tables in commonest use do not reflect this concession. Some Solutions to the Problem
Two geometrical arrangements that provide more space for elements as the periods become longer, obviating the necessity for footnotes, are the spiral and circular forms. The spiral form, both twodimensional and three-dimensional, has been proposed many times. One disadvantage is that the periods run into one another without separation. Another is that this form does not connect atomic structure and periodicity. Numerous circular tables are listed in Mazurs' book.' Since the one about to be described is essentially the long table bent into a circle, no great originality is claimed for its geometry, though its group numbering is thought to be unique and more useful. However, the choice of arrangement is more than a matter of geo-
metrical convenience; correlation between a table in concentric circles and the structure of the atom is possible. Therefore the author likes to call it an atomic form table. Cabra12recognized this point in calling his table the Natural Classijications of the Elements. The Atomic Form
Before assembling an atomic form table, it is well to recognize that there is a choice of a guiding principle to make. It is necessary to decide which is more fundamental atomic structure or chemical properties. R e solving to be completely objective, we must champion the side of physics and admit that the relation is best conceived as one of cause and effect,with atomic structure determining chemical properties. Consequently, i t is desirable to construct first a table that is periodic with respect to atomic structure. Chemical periodicity will then be assured and a notation useful to chemists can be superimposed. Figure 1 shows how the atomic form is begun, the element symbol being placed in the position of the added electron. This process is continued (Fig. 2) up to cal-
Figure 1.
Starting tho table.
cium, which can be deduced to have the structure 2-8-8-2 by counting the symbols preceding it. The symbol of scandium (2-8-9-2) cannot be put in the shell of the added electron without making it appear to follow magnesium. Therefore it is written in the same arc as calcium and the region enclosedwithin an outer boundary marked 11-Aato indicate that, like calcium, i t has two electrons in its outer shell. Of course, this actually r e sults in a conventional long-form table bent into a circle but with the lanthanides and actinides inserted in their rightful places. However, the group numbering (Roman numerals) is different, being (with a few minor exceptions) the number of electrons in the outer shell. For chemical purposes (next section) it is best to consider that the inert elements have no electrons in their outer shell (Group 0). Chemical Considerations
Presented in part at the Delawere Valley ACS Regional Meeting, Philadelphia, Pa., February 16, 1956, and more completely at the 132nd Meeting of the American Chemical Society, New Yark, September, 1957.
So far, this provides a purely physical table, but the chemical significance of Group 11-A is striking. With a
' MAZURS,E. G., "Types of Graphic Representation of the Periodic System of the Elements," E. Mazurs, La Grange, Ill., 1957.
ZBol. did&ticoescola agron. Elisen Maciel, Inst. agron. Sal (Pe lotas, Brad) No. 1, 3-6 (1951). a A is used because there is aleo a Group 11-B.
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few exceptions, all elements in this group exhibit an oxidation state of +2. This fact is not brought out by group numbers in conventional tables. Pursuing this point further, it is also a fact that +2 is the lowest oxidation state (other than zero) of the elements in this group. Similarly, 1is the lowest oxidation state of the elements in 11-B, while the other group numbers are the highest oxidation states for their elements (+4 for Pb, +5 for Bi, etc.). This suggests adding a second (arabic) group number to give the range of oxidation states. I n the lefthand side of Group 11-A, this will he the conventional group number. Lower limits to the oxidation states of the metals in Groups I11 to VI are placed along the dotted metal-non-metal boundary. Recurring subgroup numbers in Group 11-A are differentiated by letters: 3A, 3B, 3C, 4.4, 4B, etc. On lookmg a t the table, a student can deduce that the oxidation states of manganese (Group 11-A:7) range from +2 to +7, osmium (Group 11-A:8) from +2 to +8, hydrogen (Group I-C:-1) from +1 to -1, etc. Exceptions can he found but are not numerous, for in-
+
Figure 2.
Periodic table, "atomic form.'.
stance, +6 is the upper oxidation limit for iron (not +8) and + 3 for gold, (not +2). Anot,her type of exception, caused by forcing the elements into an oversimplified pattern, is that the ground states of Cr, Nh, and Mo have only one electron in their outer orbits. However, the chemical fact that is of more significanceis that they lose two electrons but not just one (in stable compounds). Hydrogen, always a misfit, is placed in line with the halogens because it has a range of oxidation states that is analogous to the halogen range of + 7 to - 1 (Group VII: -I), but limited by its atomic structure, to a maximum of 1. On the basis of range of oxidation states, placing hydrogen in group I-A: 1with the alkali metals would he illogical. Similarly, the inert element group is best designated as 0: 0, rather than VIII :0. The lanthanide and actinide group numbering requires further comment. While exceptions to the indicated ranges of oxidation states occur, the enlarged snbgroups, e.g., 11-A :6A, bring together similar elemeut,~: U (+2, +3, +4, +5, +6), NP (+3, +4, +5, +el, Pu (+3, +4, +5, +6).
+
Copyright 1958.
Volume 36, Number
7, July 1959
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