Some Reflections on the Periodic Table and Its Use W. h a r d Femelius Kent State University, Kent, OH 44242 Recentlv the author had occasion t o nartici~atein a review of some aspects of the periodic tabie ( I ) . he point of this review was to note differences in the desianation of families (more often called groups today) in the periodic tables commonlv used and to call for suggestions to resolve the confusion caused by these differences. Considering the importance of the periodic table in the teaching and practice of chemistry, the number of suggestions was dkapp&tingly small and the individual suggestions had little in common. Further. manv of them related more to the form of the table than to the designation of families in it. he bodies charged with resolving the conflict had very little guidance from their repeated requests for suggestions. Unfortunately the recommendation (2) for the designation of groups involved simultaneously arecommendation for the form of the table. There have been a number of objections to the latter as well as the former. The total experience has caused the author to reflect in depth on the periodic tahle and its use. He would like to sharithese reflections with other users of the table. Early Perlodlc Tables The periodic tahle (3) was essentially an extension or svstematization of the existence of families of elements with closely similar behavior. In another sense, the periodic table was a renresentation of the periodic law: the ~ronertiesof the elements and their comparable compounds are periodic functions of their atomic w e i ~ h t s(later atomic numbers). Because the periodic law was-never quantitative, it might more correctly have been designated as a "principle" rather than a law, hut Newlands had already established a precedentwith his "Law of Octaves". The success of Mendeleevin predicting the properties of elements and their compounds, which at the time were unknown, and the ease with which the noble gases were fitted into the table completely overshadowed other aspects of the table where there was a good deal of uncertainty: reverse order of increasing atomic weights for Co and Ni, for Ar and K, and for Te and I; position of hydrogen; increasing length of period (periods 4 and 5 had more members than periods 2 and 3 and period 6 more than 4 and 5)l; three separate vertical columns in group 8; the number of rare earths to be expected; exactly where new elements were to be fitted in once they were found; etc. When one used the periodic law either to organize his knowledge or to predict the unknown, one had first to take knowledge to the periodic table before one could make a reasonableguess about what one did not know. The periodic table helned one to decrease knowledee entroov. . In order to aid in the use of the periodic table, users of i t addedvarious items to the formal arraneement of svmbols in columns and periods: atomic weights, numbers to designate groups, A and B to designate the subgroups (beginning with period 41, formulas for characteristic hydrides and oxides, etc. However, none of these altered the character of the table; i t remained essentially a series of filing cabinets, albeit a very useful set, into which one placed in orderly fashion
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'The number of elementsthen known was insufficient to permit any Statement about oeriod 7. 2As Jseo "ere "oxidation number" refers to species involvng a Single atom of !he element to which reference is made. Formal oxidation nmbers lor a c Aer of like elements or for a homoatomic chain are not considered.
one's own knowledge, and, havingdone this, one was enabled to glimpse the unknown. Inasmuch as each user had different uses in mind, he had his preference as to the form of the table and to the auxiliary information beyond the symbols of the elements. Effect on the Perlodlc Table of Atomlc Numbers The establishment of atomic numbers did much to clarify the periodic table: the justification for the apparent abnormal location of Co and Ni, of Ar and K, and of Te and I; on the basis of U being element 92, the number of expected elements was known. However, atomic numbers did not settle how many rare earths were to beexpected (whether or not element 72 was a rare earth ( 4 ) ) , nor did they answer the question of why there should be a periodic table.
The Perlodlc Table In Relatlon to Electron Dldrlbutlon in the Atoms of Elements It was the concept of the nuclear atom and the electron distributions in the-various kinds of atoms that gave greater significance to the periodic table. Now the real basis for the existence of families of elements, variations from element to element across a period, variable length of periods, the number of rare earths to be expected, etc., became evident. The power of the periodic table was thereby enhanced and thus its usefulness~However, the manner of using the table was not fundamentally altered. In order to utilize the periodic table maximally, one must continue to supply knowledge far beyond that found in the table itself. Yet the feeling that the table was the font of much chemical information gained credence. There were significant losses as well as gains from this altered point of view. Also the passage of time has witnessed the alteration of the usage of certain terms. Terms and Concepts Related to the Table Few, indeed, would challenge the statement that the properties of the elements and of their compounds are intimately related to the electronic distribution of the atoms. The periodic table, particularly in the long (or extended) form, readily reflects electronic structures. Thus there is an s block of elements where there are electrons in outermost ns orbitals (n is the major quantum number of the outermost electrons), a D block where there are electrons in the outermost n orbitals, a d block where the (n-ljd orbitals are present, a i d a n f block where (n-2jf orbitals are present (Fig. 1). Bohr (5)classified elements in a similar but slightly different manner (Fig. 2). He was concerned whether electronic shells were complete or incomplete. (Note that shell is used here rather than group toavoid any confusion with the use of group to designate a vertical column in the periodic table.) Type I. Type 11.
All electronic shells are complete (noble gases). All electronic shells hut the outermost are complete (main group elements). Type 111. The two outermost shells are imperfect (transition elements). Type IV. The three outermost shells are imperfect (innertransition elements). Associated with each of these types are certain characteristic properties. Types I and I1 require no comment. Type I11 contains those elements which exhibit paramagnetism in all compounds except those of maximum oxidation number2, Volume 63 Number 3 March 1986
263
have strong catalytic properties, and most of whose simple monoatomic ions are colored and have variation in oxidation states by single units in contrast with type I1 elements where the variation in oxidation states, when exhibited, is almost always by steps of two. Type IV, containing most of the rare earths, also exhibit paramagnetism, variable oxidation states by single units and except for a few elements have colored ions as do the type 111 elements. However, the magnitude of the paramagnetism and the nature of the absorption spectra of the ions are quite distinct from those of type I11 elements. Modern Basls of the Perlodlc Table I t is fair to say that most chemists agree that the periodic table today, no matter what form i t may have, must reflect electronic configurations. However, complete consistency is difficult to attain.3 The regular order of fillings, p, d , and f orbitals throughout the whole array of elements is followed so closely that it may be easy t o overlook the irregularities even though those irregularities do have chemical consequences. One must always keep in mind that the electronic configurations one finds in the usual tables apply to the
3Jorgensen (6)argues convincingly that there is "no straight-forward relation between the ground electron configurationof the neutral atom and the chemistry of a given element."
"normal" atoms-those separated in space and thus not influenced by surrounding atoms and a t such a temperature that there is no radiation of energy from the atom. This condition is seldom realized in chemical compounds and even in the free element in solid or liquid condition. Consider the following cases:
Note that for the "normal" chromium atom the tendency for one electron t o be present in each of the five 3d orbitals is such that accompanying the added electron (over V) is the withdrawal of another electron from a 4s orbital. Similarly, in copper the tendency to completely fill allof the 3d orbitals is such that once again only one electron is present in the 4s orbital:
For the second long period, the withdrawal of one electron from a 5s orbital begins a t Nb and continues through Rh: Zr
Nh Mo-Rh
S-block
d-block
p-block
I
SC
Ti
V
Cr
Mn
Fe Co
Ni
Cu
Zn
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
La-Lu HI
Ta
W
Re
0s
I
Pt
Au
Cd Hg
Figure 1. Various blocks of elements,
Type Ill
Cr Rb
Mn
Fe Co Ni
Sr
Figure 2. Types of elements according to Bob
264
Journal of Chemical Education
Is2;2@,2pB;3s2,3p6,3d1Q; 4@,4p8,4d2;51% la2;2s2,2p6;3s2,3pe,3d10;4~2,4~6,4d'; 5s' la2;2s2,2p6;3s2,3pfl,3d10; 4~2,4~$4d"s; 581 (except t h a t Tc mimics Mn ..;4s2,4p6,4d6;582)
.
P d is more irregular with no electrons in the 5s orbital
(One must remember that orbital distributions are for the "normal" atoms as defined above. These conditions do not hold for palladium metal because it is an electrical conductor and conduction takes place only when electrons have been promoted to a conduction band.) Among the elements with incomplete 3d orbitals (period 4) theie is a regular piogression of maximum oxidation numbers from Sc"' t o Mnv" and then a decrease to Cu. The ~roeressionin neriods 5 and 6 is-different in' that the oxidation number is a maximum a t RuV"' a n d OsV"'. Further, beginning with Hf the resemblance between the transition elements of period 6 and the corresponding element of period 5 is much closer than that between the corresponding elements of periods 4 and 5. This departure from strict regularity can be attributed to the lanthanide contraction (7). The contraction in the size of atoms eoine
that the atoms of all of the rare earths of greater atomic number than Gd (64) are slightly smaller than Y (39). The effect of this contraction is such that the chemical hehavior of the element Hf is a mimic of that of Zr (difference in atomic radii of 0.1% (4)). In fact, these varied anomalies make it impossible to construct a single table which representsfully the factsof the chemical hehavior of the elements. Comments
1) While, in general, chemical behavior is related to electronic structures for the "normal" atoms. there are disc re^. ancies which require explanations beyond just the steady and orderly progression in electronic structures. 2) The variation of oxidation states between elements of types I11 and IV,on the one hand, and type 11, on the other, results from different types of situations heing involved in the two cases. For types 111 and IV variahle valence results from different numbers of electrons heing recalled from d or f orbitals, whereas that for type I1 usually results from sharing of different numbers of pairs of electrons. 3) The chemistrv of the chemical elements and their compounds does not fit neatly into the classification of Bohr. The differences in energy between nd and (n l ) p states for Sc, Y, etc., is so slight that S C ~ +etc., , are never encountered, and thus the chemical transition character is almost nonexistent. As the filling of d orbitals nears completion (e.g., Ni, Pd, Pt), electron release requires relatively s e a t e r e n e r n and so M2+ (representing the loss only of i h i outer s electrons) are the characteristic ions. The compounds representing the release of some d electrons (e.g., M4+) are known largely in the form of coordination derivatives. In the case of Cu, Ag, and Au the last remaining empty (n - 1) d orbital may hold an electron more tightly than a n s orbital, and thus M+ is common. Similarly in Ce the two electrons (over the Ba structure) readily are removed to form Ce4+and for Yh the last electron over the previous atom is held so tightly that Yh2+can he formed readily. A word or two more about the innertransition elements. Most tahles smugly place the rare earths in a single position in the periodic tahle implying the marked similarity in properties of the individual elements. Yet there are variations in oxidation states among the f hlock elements which are similar to those found in the d hlock elements and which can ~ - he ~ related to the general principles previously encountered about the relative stabilitv of svstems in which each of the f orbitals is "half filled" and the tendency t o anticipate this situation (lower oxidation number) or revert to it lhieher oxidation number). The colors of 'the triply chargdd ;are earth ions reflects the numher of half-filled orhitals. However, these simple relationships do not hold for the actinides where the tendency t o display higher oxidation states per-
+
~
Max Period Cwrd NO. NO. 1. 2
~
~
~~
~~
~~~~~
~~
sists from thorium to berkelium. No wonder early workers regarded uranium as a congener of tungsten rather than an innertransition element. 4) Consider some vertical relationships. Everyone is familiar with the decrease of electronegativity with increase in atomic numher for the elements a t the far left and right of the long form of the periodic tahle. For the alkalis and alkaline earth elements, this means that CsOH and B ~ ( O H ) Z are the stroneest bases in their families. whereas LiOH and B ~ ( O H )are Z ;he weakest. For the e~emedtsto the right in the table one must he careful to distinenish between simnle hydrides as acids and the 0x0 acids. HI is the strongest HX acid, and NH3 is the strongest XH3 base. The reverse order holds for the HX04, H2SOa,etc., acids. While this order of electronegativities holds for the early transition elements (Sc, Y, La, Ac) i t does not hold for Cu, Ag, Au, nor for Zu, Cd, Hg. The use of lead dioxide and sodium bismuthate as oxidizing agents recalls another generality. On the right side of the periodic tahle, the stability of the highest valent oxide or salt decreases with increasing atomic numher in the various families of the main group elements. (This is associated with the "inert pair" of outer s electrons.) The reverse relationship holds for tvne 111 elements: CrOl is a strone oxidizine - aeent " while ~ 0 ; ; snot. 5) The maximum coordination number exhibited hv the elements increases in irregular stepwise fashion with increasing period numher. 6 ) Finally there is need to note some cautions in thematter of nomenclature. The terms "s, p , d, and f blocks" have definite and consistent meaning as do "types I, 11, 111,and IV". However, the meanings are not the same. Prior to Bohr's work, the term "transition element" referred t o group 8 elements. Today the term is frequently used as a synonym for "d hlock" elements and so includes Zn, Cd, and Hg even though these elements do not possess those properties of naramaanetism. colored ions. and variahle valence which ~ o hregarded r as the distinguishing characteristics of transition elements. The ions Zn'+. Cd" as well as CII-. Aet. and Au+ have an outer shell of 18klectrons instead of 8 &istrue of the so-called inert-gas ions. The former might he called "pseudoinert-gas ions".4 Sanderson (8)labeled the family ~ containing Zu, Cd, and Hg as M 2' (M = main group). The one characteristic Zn possesses in common with Co2+,Ni2+, Cu2+, etc., is approximate size. IS this sufficient to justify
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-
'The author encountered this term years ago in a discussion of the Fajans' rules governing the preference of two oppositely charged ions to form covalent rather than ionic compounds. However, neither he nor those of his colleagues who recall encountering this term can locate the exact reference.
Valence (or oxidation) numbers 1
2
H
(He)
3
11.
4
Li
Be
111. IV.
6
Na
6
K
Ss
V.
8
Rb
Mg Ca Sr
VI. ViI.
8
Cs
Ba
L
?
Fr
Ra
A
