Up to now, only the option of hydrogen placement in a single location has been considered. To avoid settling upon one option or the other, some early periodic classifications placed hydrogen above both halogen and alkali metal Groups [14]. An interesting 18-column table of this type was devised by LeRoy in 1931 (Figure 3.3) [15]. Of relevance to discussions in later chapters, LeRoy (and other contemporary chemists) placed boron and aluminum in Group 3A (now Group 3), not Group 3B (now Group 13).
The dominance of hydrogen was repeated more cleanly in a pyramidal Periodic Table with hydrogen at its apex [16]. The dual linkage of hydrogen to lithium and fluorine became very popular. Some commercial Periodic Tables and some of those within textbooks, locate an “H” symbol to head both Group 1 and Group 17. In fact, many classrooms and lecture halls in schools, colleges, and universities are adorned with this version (Figure 3.4). This “dodging the choice” is actually worse pædogically, with students coming to believe that hydrogen is both an alkali metal and a halogen.
Table 3.1 Summary of placement reasons in Group 1 and Group 17
Figure 3.3 The 1931 LeRoy version of the Periodic Table (from Ref. [15]).
Elemental Dual Identity
Whether the Reader is supportive of the placement of hydrogen in Group 1 and Group 17, it raises an important point that is usually overlooked: that is, can an element be placed in two locations? It is a crucial philosophical point in this book, as in several instances in later chapters, elements seem to fit in more than one location. Though others before him actually adopted a dual location for an element, Rich seems to have been the first to emphasize that dual (or triple) identity/location of an element was a very significant and fundamental conceptual and philosophical break from the idea that each element occupied a single location [17].
Figure 3.4 The first three Periods of an earlier version of the Fisher ScientificTM wall chart.
Hydrogen as a Member of Group 13 or 14
If the element does not fit well at either extreme of the Periodic Table, where, then? The earliest mention of placing hydrogen in the middle of the Periodic Table was in 1893 by Rang. He devised one of the first 18-member forms of the Periodic Table, numbering the Groups from I to VIII and then I to VII again. He placed the symbol for hydrogen at the head of the second Group III (Figure 3.5). Then in the caption, he noted [18]:
H may not be exactly in its true place, still it cannot be very far from it.
In 1964, Sanderson proposed that hydrogen fitted better in the middle of the Periodic Table, specifically, over carbon. His reasons were that the electronegativity of hydrogen was closer to that of the Group 14 elements and that hydrogen had half-filled outer electron shells. He was careful to suggest that, even though hydrogen should be placed over carbon, it needed to be in a “separate independent position” [19]. Perhaps to make the point unambiguous, his own version of the Periodic Table (Figure 3.6) shows hydrogen bridging over boron and carbon.
Sanderson’s choice of placement of hydrogen in Group 14 was supported, and expanded upon, by Cronyn [2]. Cronyn pointed out the similarity in the preference for covalent bond formation by both hydrogen and carbon: for example, the H−H bond has a strength of 436 kJ·mol−1 while that of the C−H bond is 439 kJ·mol−1. He also commented upon the similarities of the chemistry of hydrogen to that of silicon, cementing the link of hydrogen with Group 14. In his own Periodic Table design, Cronyn reinforced his argument by displaying trends in ionization energy and electron affinity (in eV), showing that the values for hydrogen fitted perfectly in the sequence (Figure 3.7). Electronegativity values were also inserted in Cronyn’s Periodic Table, but the value for hydrogen better fitted a pattern for the Group 13 elements.
Figure 3.5 The center of Rang’s 1893 Periodic Table design showing the location for hydrogen (from Ref. [18]).
Figure 3.6 Part of Sanderson’s 1964 Periodic Table, showing the placement of hydrogen (see Ref. [19]).
However, as with the assignment of hydrogen to Group 1 or 17 (or both), pædagogic confusion is caused by a student perception that hydrogen is indeed a formal member of Group 14.
Hydrogen as a Member of Group 1 and Group 14 and Group 17
In all of Laing’s Periodic Table proposals, he believed that the Periodic Table was a means of visually displaying chemical linkages and that two or more locations of a single element were educationally beneficial. Rich and Laing suggested that the solution to showing the similarities of hydrogen to each of Group 1, Group 14, and Group 17, was to show hydrogen as a member of each of the three groups (Figure 3.8) [20].
Figure 3.7 Part of the Periodic Table by Cronyn showing the values of electronegativity, upper left; ionization energy, lower left; electron affinity, lower right (see Ref. [2]).
Hydrogen on Its Own
Why should we continue to try to fit hydrogen into a table that is simply a human construct? If hydrogen does not “fit in” perhaps it is because it indeed does not fit in and is best regarded as a unique element. The proposal by Kaesz and Atkins utilized the empty space above the transition metals to place the hydrogen “box” [21]. In doing so, Kaesz and Atkins rejected the two-location model of heading Group 1 and Group 17, stating it had to have a single location. Believing that the chemistry of hydrogen was totally unique, they placed hydrogen central but clearly level with the other 1st Period element, helium (Figure 3.9).
This idea prompted a rapid reply from Scerri who questioned this whole direction of involving observable chemical properties as a
