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A NEW PERIODIC CHART, WITH ELECTRONEGATIVITIES' R. T. SANDERSON State University of Iowa, Iowa City,Iowa
T H E R E are many qualities of its atoms which may influence the physical and chemical properties of an element. It is possible, however, to choose only three which alone, when considered within the structure of a conventional periodic table, may serve quite satisfactorily as a basis for reasonable explanations of the nature arid behavior of the elements. These are the relative atomic radius, the general type of electronic configuration, and the relative electronegativity. This paper describes a new chart which has been designed to portray clearly and vividly these three qualities. Each element is represented by a colored disc against a background of gray or black. The size of the disc indicates approximately the relative nonpolar covalent atomic radius. The background shade denotes the type of electronic configuration. The color of the disc represents the relative electronegativity. The arrangement of the elements is that of the conventional "short form" of the periodic table, with vertical alignment according to valence, and no physical separation of the subgroups. Such separation is nnnecessary because the electronic distinction between subgroups is clearly shown. A black background is provided for the inert elements and those resembling the inert elements except for the outermost shell. A dark gray background denotes the "18-shell" element,^, those having an outermost shell of a t least one and less than eight electrons over a shell of 18. A medium gray background indicates those elements with partially filled d and f orbitals, the transition and inner-transition metals. Copper, silver, and gold are represented as having both 18-shell and transition type features. The numerical electronegativity range is based on stability ratio values from 0.49 to 5.75.2 The qnantitative relationship between stability ratios and Panling electronegativity values has been r e p ~ r t e d . ~The corresponding color range from low to high is from pale yellow through yellow, orange, red, and violet to bright blue. Limitations in the color sensitivity of human visual perception made it pointless to attempt to represent in this way differences smaller than about tmo-tenths of an SR unit. Consequently any perceptible color difference on the chart represents a truly significant electronegativity difference. Colors assigned t o transition elements are less certain but -
Presented before the Division of Chemical Education a t the 128th Meeting of the American Chemical Society, Minneapolis, September, 1955. * SANDERSON, R. T., J. CHEM.EDUC.,29, 539 (1952). * SANDERSON, J. C h m . Phqs., 23,2467 (1955).
probably approximately correct. A green color means no useful estimate has been made. ADDITIONAL FEATURES
Electronically, or by valence, hydrogen belongs in both Groups I and VII. Its electronegativity, however, corresponds to a chemistry intermediate between the two extremes. Hydrogen is therefore placed above the other elements and just to the left of carbon. The general locations of the rare earth metals, or "lanthanides," and of the "actinides" are indicated without devoting space to all the individual members. However, the first and last lanthanides are represented in order to demonstrate the "lanthanide" contraction, with its corresponding increase in electronegativity (and therefore diminished basicity). The method of estimating the electronegativity values of the reactive elements is that of comparing the average electronic densities of their atoms with those of the inert atoms.2 Consequently the inert elements are represented by yellow discs to show that the average compactness of their atoms is very low and comparable to that of the reactive elements whose electronegativity is very low. There is no intention of implying that the inert elements can be assigned a significant electronegativity value. The ions of sodium and fluorine are included to demonstrate the extreme effects of electron gain or loss. Partial or complete loss of an electron produces shrinkage of the remaining electron cloud and a corresponding increase in electronegativity. Gain of negative charge has the opposite effects. APPLICATIONS TO INSTRUCTION IN CHEMISTRY
The relationship between electronic configuration and atomic radius is of course readily apparent. I t may be pointed out that in general the addition of s and p electrons within a principal quantum level (simultaneous with the addition of protons to the nucleus) over an inert type configuration results in diminishing atomic radius until the inert element is reached. The selection of reasonable atomic radii for the inert e l e ments must he somewhat arbitrary, but for reasons previously discussed2 it is believed that the radius increases with the completion of the outer shell of two or eight. A further increase corresponds to the beginning of a new principal quantum shell. A similar but less extensive shrinkage occurs when s and p electrons are being added but the shell immediately underlying is one of 18. A similar but still less extensive shrinkage seems to occur with the filling of d
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orhitals, although the radius data here are much less certain. There may possibly be a slight increase when the d orbitals reach the half-full mark and again when the shell of 18 becomes complete. Finally, there is the lanthanide contraction-a smaller but still significant decrease in radius wit,h the filling of the 4f orhit,als. The relationships between atomic radius and electronegativity are rlearly visihle throughout the chart. Perhaps the most vivid representation is that of the differences between the most reactive metals and the most reactive nonmetals. The former are displayed as large yellow discs which offer the greatest possible contrast with the small blue discs of the latter. It may be pointed out, however, that most of the elements are intermediate in elertronegativity, and that for these, atomic radius as well as elertronegativity helps to determine the physical character of the element. Thus, boron, hydrogen, carbon, silicon, and phosphorus have average or lower than average electronegativities, but mith their small atoms, exhibit some nonmetallir qualities. Gold, mercury, thallium, lead, and hismut,h, with a very similar range of electronegativit,ies,&have larger atoms and are all metals.
than two elements, the bond polarities are influcnccd by all the atoms present in a molecule. The principle of electronegativity equalization states that when atoms initially different in electronegativity combine, they change in electronegativity to an intermediate value. (This value is postulated to be the geometric mean of the electronegativities of all the atoms before combination.) Part,ial charges resulting from elertronegativity equalization can be e ~ t ~ i m a t e dI. t~ may be pointed out that bond polarity is not the only fartor determining the ease of separation of a molecule into ions. Assuming equal polarity, ease of ionizatio~iis expected to be greater, the greater the internuclear distance (the lover the lattice energy) and the greater the solvation energy for the ions. Other factors may also have influence. Such generalizations as that of increasing basicity of oxides in descending a major group are therefore not necessarily in conflict with alternating elect,ronegativities sometimes observed. Differences between the subgroups show so rlearly on the chart that extension to the "long form" is unnecessary. For example, copper, silver, and gold are seen not only to be electronically very different from the alkali metals, but also to have much smaller, more compact atoms of murh higher electronegativity. In a similar manner, zinc, cadmium, and mercury are seen to bc auite distinct in several resoects from hervllium, magnesium, and the alkaline earth metals, although there are interesting similarities in the chemistry. Further, the arbitrariness of grouping (as in most "long form" charts) either gallium, indium, and thallium, or scandium, yttrium, lanthanum, and t,hc rare earth metals under horon and aluminum is ~iuit,e ohvious. Students can see a t once that neit,hei of these subgroups closely resembles boron and aluminuin electronically. It can he pointed out that in rampounds with the active nonmetals, scandium, etc., appear better to cont,inue the trend begun by boron and aluminum, as expected from the lower electronegativities and larger radii. On the other hand, gallium, etc., bear a rloser resemblance from the standpoint of kinds of valence orbitals, resulting a t least in a more similar rovalent chemistry. For example, orgallo-metallic rompounds occur mith horon, aluminum, gallium, indium, and thallium but not with scandium, yttrium, lanthanum, or t,he rare earth metals. Similar relationships are observable in subsequent groups. Alternations in elertronegativity within a "major" group, and especially an increase in the transition from 8-shell to 18-shell type element, are shown in the The dirert,io~rand cst,errt of bond polarity ill hillary chart. These alternations are useful in explaining many chemical "anomalies" of the periodic table.' compoullds is readily evidellt: the initially more electronegative element always becomes more negative, CONSTRUCTION AND AVAILABILITY and the greater the initial difference in electronegativity, The chart pict,ured is 30 X 36 inches in size. A concompounds of more the greater the bond polarity, venient scale of disc size then ranges from s / ~inch ~ ' Recent studies, ss yet unpublished, have permitted a reviaion diameter for the H atoms(0.37 A. radius) to 2 ~ / *inch
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of earlier reported stability ratioa of thosc elements to the following: Au 2.88, Hg 2.93, TI 3.02, Pb 3.08, Bi 3.16. Vahtes for the lixhter elemenbs remain: B 2.84, H 3.55, C 3.79, Si 2.62, P 3.34.
~ANDDR J. CHEM. ~ X , EDUC., 32, 140 (1955). SANDERSON, J. Am. Chem. Soc.. 74, 4792 (1952).
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diameter for thc Fr atom (about 2.3 1.radius). The author will gladly supply details of the color-electronegativity code and size specifications to anyone wishing to construct surh a chart,. I t has been found possible t , make ~ quite satisfactory color slides (35 mm.) for projection on a screen for classroom use. (The State University of Iowa Photo Service has agreed to provide such slides, each an original photograph, not a ropy, taken from t,he original chart, a t moderate cost.' Two surh slides are recommended to be alternated every fev minutes of a prolonged showing, to minimize deterioration by heat aud light.) Address orders to S.U.I. Photo Service, State University of Iowa, Iowa City, Iowa, and enclose one dollar for each glassrnount,rd slid? wanted.
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For constant display, a useful chart can be assembled by preparing a set of discs in appropriate scale from plywood or pressed wood, and gluing the appropriate color of paper to each. These discs can then be hung or othenvise attached to a periodic wall chart. Pending the construction of a more permanent mall chart, we have placed such a set of discs in appropriate positions on a 15-foot periodic chart painted on the front wall of a general chemistry lecture room seating over 400. With the scale such that hydrogen is not quit,e two inches in diameter, cesium is ahout 12 inches in diameter and the entire chart can readily be seen throughout the room. I t makes an interesting and very colorful display and has been found very helpful in tearhing the relationships among the elements.
