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Chapter 3
First Period Problems
The placement of hydrogen in the Periodic Table is still a cause of dissention. In some versions of the Periodic Table, uniquely among the elements, the hydrogen symbol appears twice. Here, the different proposed locations will be described and contrasted. Though all common versions of the Periodic Table place helium in Group 18, some chemists propose its placement should be in Group 2, based upon electron configuration.
Rouvray has lauded the accomplishments of the Periodic Table [1]:
The periodic table is also deeply reassuring in that it accounts for and assigns a specific position to every element; even elements still waiting to be synthesized have their rightful place awaiting them.
This claim is not quite correct, as Cronyn described in 2003 [2]:
Quite remarkable, after 130 years of construction, the place of hydrogen in the periodic table is still the subject of doubt, confusion, and inadequate explanation that appears to be little more than numerology.
Despite Cronyn’s somewhat hyperbolic language, it is indeed true that the position of hydrogen remains a contentious issue. So much so, almost this whole chapter is devoted to the problem.
Hydrogen Location: An Overview
First, it should be noted that the hydrogen location problem is partially an artifact from using the rectangular short form or long form of the Periodic Table. Other designs, such as spiral or circular tables, enable the “H” symbol to be placed at the center, making the issue is moot. One such hydrogen-centered Periodic Table was proposed by Piutti in 1925 [3] while a more elegant spiral design was drawn by Griff in 1964 [4] (Figure 3.1).
Figure 3.1 Griff’s spiral Periodic Table of 1964 with “H” at the center (from Ref. [4]).
However, as the conventional Periodic Tables are almost ubiquitous, placement in these “standard” tables will be the focus of this chapter. While other elements are largely locked into place by the ordering of the atomic number, the fact that hydrogen and helium are above the next eight-member Period results in this dilemma.
Before discussing the pros and cons of each option, it is useful to the Reader to see the common alternatives listed. Most of these options have been reviewed by Petruševski and Cvetković [5]:
•Hydrogen as a member of Group 1
•Hydrogen as a member of Group 17
•Hydrogen as a member of Group 1 and Group 17
•Hydrogen as a member of Group 13 or 14
•Hydrogen as a member of Group 1 and Group 14 and Group 17
•Hydrogen on its own
•Hydrogen as a member of Group 1 and helium as a member of Group 2
Hydrogen as a Member of Group 1
Most commercial Periodic Tables show hydrogen as a member of Group 1. Cronyn has described the situation elegantly [2]:
So there is poor hydrogen, denied a chemical family to call its own, thrust like an unwanted orphan into a foster home where its chemistry cannot even be discussed in the same breath with the alkali metals where it now resides.
The strong argument for the placement of hydrogen in Group 1 is that of electron configuration, as championed by Scerri [6]. As the alkali metals have a valence electron configuration of ns1, then, as the ground state electron configuration of the hydrogen atom is 1s1, it follows that hydrogen should be placed at the head of the Group 1 column (Figure 3.2).
There are some supporting chemical arguments. In theory, the hydrogen atom ionizes to give a (positive) hydrogen ion, just as an alkali metal atom ionizes to give an alkali metal cation. Of course, hydrogen does not, in a chemical context, produce a free proton but is, instead, part of a positively charged network: [H(OH2)n]+. There is the similarity in that, in aqueous solution, the alkali metal ions, too, are all strongly aquated.
Figure 3.2 A traditional long form of the Periodic Table with hydrogen in Group 1 (from Ref. [6]).
Nevertheless, the alkali metal group form one of the best examples of systematic group trends. To place hydrogen at the top of it implies that hydrogen, itself, is an alkali metal. Yet every comprehensive inorganic book discusses the chemistry of hydrogen in a whole separate unit from that of the alkali metals [7].
Claims of the metallic nature of hydrogen under exceptionally high pressure are used to support the argument that hydrogen is actually an anomalous alkali metal. It was in 1935 that Wigner and Huntingdon were the first to propose on theoretical arguments that under extremes of compression, dihydrogen would undergo a transition to a metallic allotrope [8]. Much high-pressure research has been attempted in order to obtain physical evidence of the existence of metallic hydrogen. In part, such a search relates to the possibility of metallic hydrogen being in the core of the gas giant planets, Jupiter, Saturn, Neptune, and Uranus.
However, the goal of some researchers has been to provide evidence that hydrogen really should be considered as an alkali metal [9]. There has been a claim that solid, metallic hydrogen has indeed been produced [10]. However, does the formation of a metallic allotrope under such extreme conditions really provide evidence that hydrogen belongs to the alkali metal family? Moore has pointed out the fallacy of that argument [11]:
Boron, oxygen, sulfur, selenium, tellurium, and phosphorus all can be made conductive under pressure, but only in the case of hydrogen is metallization thought to vindicate its predicted properties.
The most intriguing evidence for hydrogen as a member of Group 1 has come from a different route. A compound has been synthesized containing a four-coordinate hydrogen atom that occupies the same site as lithium or sodium [12].
Hydrogen as a Member of Group 17
It can be argued that the 1s1 of hydrogen is one electron short of a complete electron shell, just as are the ns2np5 electron configurations of the halogens. However, the reasons for hydrogen placement as a member of Group 17 rely more on the chemical arguments. Supporting evidence for placement in Group 17 is that hydrogen, like the halogens, forms a stable diatomic molecule. By contrast with the halogens, dihydrogen is not as reactive as the dihalogens.
Also similarly, hydrogen can form a negative, hydride ion. Sacks argued vociferously for the placement of hydrogen unambiguously in Group 17 [13]:
A Coulombic model, in which all compounds of hydrogen are treated as hydrides, places hydrogen exclusively as the first member of the halogen family and forms the basis for reconsideration of fundamental concepts in bonding and structures.
But again, similarity in formula masks a completely different chemical behavior. In particular, the hydride ion decomposes violently in the presence of wa
ter, unlike the water stability (albeit with hydrolysis in the fluoride case) of the halide ions.
Hydrogen as a Member of Group 1 and Group 17
At this point, the discussion will be paused to summarize in Table 3.1 the pros and cons of hydrogen placement in Group 1 or in Group 17.
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 factor in the classification of the chemical elements [22]:
A very widely held belief, among chemists and others alike, is that the periodic system consists primarily of a classification of the elements as simple substances that can be isolated and whose properties can be examined experimentally. However, there is a long-standing tradition of also regarding the elements as unobservable bearers of properties, sometimes elements as basic substances.
Figure 3.8 The Rich and Laing proposal showing hydrogen as a member of Group 1, Group 14, and Group 17 (see Ref. [20]).
Figure 3.9 The placement of hydrogen according to Kaesz and Atkins (see Ref. [21]).
To summarize Scerri’s critique of Kaesz and Atkins, Scerri believed that the Periodic Table represented the order of atomic structure, not any bulk chemical behavior.
… And Then There Is Helium
Though most of the discussions have centered upon the location of hydrogen, the location of helium has also been contentious.
Hydrogen as a Member of Group 1 and Helium as a Member of Group 2
Up to this point, we have only considered the possible locations of hydrogen in isolation. The arguments were largely, but not entirely, on chemical grounds. If the Periodic Table arrangement is defined by the electron configuration, then logic demands that hydrogen, 1s1, and helium, 1s2, are placed as the top members of Groups 1 and 2, respectively. One of the Periodic Table designs by Janet in 1928 [23] followed this logic. Then in 1934, White authored a modern-style Periodic Table (reprinted by Laing [24]) displaying electron configurations to reinforce the reason for the design (Figure 3.10).
The first thorough modern discussion of this possibility was given by Katz [25]. In the article, Katz first described moving “He” to above “Be,” then shifting Group 1 and Group 2 to the right-hand side of the Periodic Table to generate the left-step (or right-justified) Periodic Table shown in Figure 3.11.
As Katz and, later, Scerri [26] have commented, the left-step table is more elegant than the conventional table. Also the orbitals are now in sequence as f>d>p>s instead of s>f>d>p. The complicating factor arises from electron configuration considerations.
Figure 3.10 The first three Periods of White’s 1934 sp
ectroscopic-based Periodic Table (reprinted in Ref. [24]).
Figure 3.11 The left-step Periodic Table (from Ref. [6]).
One may protest that helium is not a reactive metal. Bent was not swayed by such an argument. He believed that electron configuration — particularly using the left-step Periodic Table — was the essentiality: the Periodic Table is about atoms, nothing else [27]:
The answer given here to the Helium Question To Be or Not to Be? is, on both chemical and physical grounds, a resounding Yes! The most noble of the noble gases is not a Noble Gas. Helium’s natural position in Periodic Tables is in the s-block above beryllium …
Thus the debate becomes one of chemical properties versus spectroscopic energy levels. Novarro has cited a quote by Scerri that sums up the situation [28]:
Chemists may place He in the noble gas column, physicists however would rather place it above Be.
In the continuing debate, Ramíríez-Solís and Novarro used quantum-mechanical grounds to argue for helium’s place to be above neon [29]. Taking the contrary view, Grochala [30] noted that nothing is seen to be wrong in placing nonmetal hydrogen at the top of the Group 1 metals, so what objection can there be to placing helium also above a metal (beryllium)?
Hydrogen in Group 17 and Helium in Group 18 (Again)
More recently, Scerri has recanted his chemical heresy of placing helium in Group 2 [6]. By combining Period 1 and Period 2 in a single line, helium is safely returned to its chemical “home” atop the noble gases. Hydrogen, meanwhile, is given a home with the halogens, probably a more welcoming location for an element that exists in nature as a diatomic gas, not as a solid metal (Figure 3.12).
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