Edited by D A N KALLUS Midland Senior High School 906 W. Illinois Midland. TX 79705
revi/i ted
RUSSELL D. LARSEN
Texas Tech University Lubbock, TX 79409
'Principlesof Electronegativity Part I. General Nature R. T. Sanderson' Arizona State University, Tempe, AZ For more than 50 vears the concept of electronegativity has been modified, expanded, and debated. I t is of great theoretical interest today. However, the practical aspect has greatest potential for pr¬ing chemical understanding in the classroom, for it is now pwsihle not only to currelate the properties of hundreds of compounds with the electronegativities of their elements bur also to calculate accurately the energies of polar wvalent bonds. Thus students can be provided u,ith valuahle insights and understanding ofthe caudeand-effect relationship between atomic structure and the properties of compounds. They can begin to understand the uriein of bond merry -. and thus the causes and direction of ch&cal reactions. This article and one to follow are not intended as a comprehensive review of all that has been proposed. There is no deliberate slighting of any particular contribution, and references 1-7 are provided to suggest some of the more recent work that interested readers may wish to study. As in any growing body of knowledge, there is plenty of room for controversy, and no two workers will agree completely. Some new ideas are included herein that are abundantly supported by convincing evidence, although space is available for onlv a few illustrative exam~les.Sources of much additional information will he cited. The electronegativity of an atom can perhaps least controversially he defined as that property which determines how the bonding electrons will become distributed between i t and anothel atom if the two atoms become connected by a chemical bond. The atoms are recognized as initially equal in electroneeativitv if thev share the hondine electrons evenlv. If the elehronsare n i t evenly shared, &at atom that a;auires more than half share is considered to have had initially a greater electronegativity. From a practical viewpoint, this concept of electronegativity seems to satisfy unamhiguously the need to predict the direction and approximate extent of the polarity of a covalent bond, which has been the most common application. For a more thorough understanding, however, we should consider the relation of electronegativity to atomic structure and what happens to it when atoms combine. Furthermore, although various routine attempts to evaluate the electronegativities of transitional elements have been made over many years, the only values reliably applied have been those of the major group elements. Recent work has now produced reasonably reliable
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Journal of Chemical Education
estimates of some of the electronegativities for transitional elements in various oxidation states and for lower-thannormal valence states of major group elements. A summary of the general principles governing the significance and applications of electronegativity follows. The purpose of this paper is to present and discuss such principles in a general way. The second article of this pair will he devoted to greater emphasis on the practical applications. 1. Most atomic properties are a consequence of atomic structure, which in turn must be related to the inherent nature of the component electrons and nuclei. Therefore it is almost ineuitable that such properties be related to one another., if, onlv " because of their common oriein. I t should not be surprising that a particular property, here the electroneeatiuitv. " . can be deriued from or correlated with a wide uariety of other properties, with reasonable agreement among the seueral results. Since electronegativity is a measure of the effectiveness of thenuclear charge as sensed within an outer orbital vacancy, the relationship-between atomic structure and electronegativity is quite apparent. I t becomes more so when it is realized-thatthe oukrmost electrons are not very effective in shielding one another from the nuclear charge, and thus permit the effectiue nuclear charge to build up a larger and larger lead as the atomic numher increases and the numher of outer electrons changes from one to eight. This results in an enforced contraction of the electronic cloud, and the lareer effective nuclear charee over a shorter radi- o~erative . us corresponds to a steadily increasing electronegativity across the period of maior erour, . elements. The maior effects of increasing the number of outermost electron; across a period of major group elements are summarized in Figure 1. Electronegativity is obviously not merely a number but the logical and inevitahle consequence of atomic structure. 2. The best method of eualuating electronegatiuity remains to be discouered. One successful evaluation may be based on the relatiue comnactness of the electronic clouds that surround atomic nuciei. I t is loeical to suonose that anv atom unable to hold its own elecGon cloud Eompactly aroknd its nucleus (for example. an atom of the alkali metals) mav have ample outer space capable of accommodating extra electrons-hut very little attraction for them. That atom must therefore have low electronegativity. Similarly, any atom capable of holding its own electronic cloud closely and compactly despite the stronger interelectronic repulsions existing within denser clouds (for example, an atom of a halogen) should reason-
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Wherever the atomic structure changes by increasing the number of outermost electrons,
the fraction of the.tota1 nuclear charge that can be sensed within an outer orbital (called the effective number charge) is caused to increase:
In turn, this causes three important changes in atomic properties: (1)The increasing charge pulls the electrons into a more compact cloud of diminished radius:
(2) Greater charge acting over shorter distance corresponds to increasing electronegativity:
I
(3) It also corresponds to increasing (unweakened)homonuclear bond energy:
Since earhatomean form one roralent bond for each half-filled outer orbital it can supply, the changing atomic srrtrrture, by increasing the number of electrons in the four outer orhitals, changes thp availability of outer half-filled orbitals and thus the
Repulsions among bonding electrons and lone pair electrons largely determine bond angles:
All theseproperties ore the l ~ & and d incuirable eonrequencesof the changes i n otomir struclure, as now quantitatively exemplified by the aerier from Xa t o CI:
I
element no. outer e's
eff. nuclear charge ~o~alentradlus, pm electronegativlty B.E.. homonuclear (unweakened,kcal) ~ovslen~e bond angle
Na
Mg
A1
Si
P
S
CI
1 2.0 154 0.84 16.4
2 2.7 137 1.32 42.3
3 3.4 126 1.71 48.2
4 4.0 117 2.14 54.1
5 4.7 111 2.52 60.0
6 5.3 105 2.96 65.9
7 6.0 99 3.48 71.8
2 180'
3 120°
4 log0
3 90-109'
2 90-109'
none
1 none
Figwe 1. Cause (struchre)and effect (properties).
1
ably be expected to exert a strong attraction for an outside electron, provided it can also provide an outer orbital vacancy to accommodate it. Such an atom must therefore have high electronegativity. For example, the potassium atom holds 19 electrons within a radius of 196 pm, whereas an atom of chlorine holds nearly as many electrons, 17, within a radius of 99.4 pm, only about half that of potassium, and thus within only about one-eighth the volume. The electronegativity of chlorine, 3.475, is nearly eight times that of potassium, 0.445. I t may also be recognized that the natural tendency of any highly electronegative atom is to acquire partial negative charge, which causes expansion of the electronic sphere to a less compact condition. The natural result of partial electron loss by any atom low in electronegativity is contraction of the electronic sphere to a more compact condition. The "relative compactness" method recognizes that although ideally the relationship between compactness and electronegativity should be quantitative and direct, certain features of the evolution of electronic configurations with increasing atomic number appear to affect compactness without necessarily being related to electronegativity. In particular, clouds on atoms or ions near to argon, atomic number 18, appear significantly lower in average compactness than those near either neon, atomic number 10, or krypton, atomic number 36. For example, an atom of ~otassium.radius 196 nm. is closer to iubidium, iadius 216 p i , than to an atom of sodium. 154 Dm. althouah in number of electrois it is much closer to sodium (19-11) than to rubidium (37). Similarly, chlorine is closer to hromine than to fluorine. The average electronic density of an atom, defined as the average number of electrons per unit volume of the electronic sphere defined by the nonpolar covalent radius, therefore cannot be taken directly as a measure of electronegativity but must be corrected for the chanee in electronic density with atomic number observed for atoms which do not normally form bonds and is therefore independent of electroneaativity. The ratio bf the average electronic density of an atom to that of an interpolated hypothetical "noble gas" (Ma) atom of equal number is a good measure of relative electronegativity. Since this development in the early 1950's (a),individual values have been revised from time to time and refined on the basis of better data, newer knowledge, and improved understanding, .and finally changed . ( 9 ) to assian . to fluorine, a s in most conventional scales, the value 4.000. The justification for this last change was the practi-
Volume 65 Number 2
February 1988
113
calone of trying to minimize the confusion created by having different numerical evaluations for the same elements. Pauling scale values differ from intermediate values on the relative compactness scale. Future studies may suggest shifting the scale upward or downward for greater convenience or to corresoond to a auantitative theoretical standard. A complete list of electronegativities as of this date is providedin Table 1.Valuesfor major group elementsmay be seen to exhibit a steady upward trend across each period, as exoected on the basis of the steadily increasing effective nuclear charge. An old assumption is that the electronegativity decreases down a major group, hut this is not really true and should not be expected. Althoughaluminum in the third period logically follows magnesium, both having underlying eizht-electron shells. the next element of the aluminum ~-.--~ group, gallium, does not follow calcium hut zinc, owing to the insertion of the first series of transitional elements between calcium and zinc, and both zinc and gallium have underlying 18-electron shells. The significance of this change in electronic type from period 3 to period 4 in the major group elements is that the period 4 elements have smaller and more compact atoms than otherwise expected. Thus gallium atoms have iust about the same radius as aluminum atoms (instead of larger) and are more electronegative (instead of less). This trend continues across the period but decreases until no longer in evidence a t bromine. Thus the electronegativity tends to alternate down a group instead of steadily diminishing as otherwise expected. 3. Where valence electrons of a major group atom remain unused within a com~ound, to reduce the ini. thev . appear .. tial e1ertronegatiuit;of the atom. Electronegativities of subnormal oxidation states of major group elements (10) tend to be substantially lower than the normal electronegativity. These elements in their lower oxidation states have bonds to nonmetals that are more polar than expected and hence stronger than expected. This is particul&ly important in accounting for the properties of compounds exhibiting the "inert pair effect". This is the tendencv of M3 elements to exhibit an oxidation state of I and of ~4 elements to exhibit an oxidation state of 11. For examole. oartial charee on oxvgen in thalliumU) oxide, & .the . TlzO, is more than twiceas negative as i t would he if the ~~~
~
~
~
Table 1.
s
TI (11) TI (Ill) TI (IV) V (11) v (Ill) V (IV) V ('4 Cr (11) Cr (Ill) Cr (IV) Cr (V) Cr (VI) Mn (11) Mn (Ill) Mn (IV) Mn (v) M" (VI) Mn (VII) Fe (Ill Fe (Ill) Co (11) Co (Ill) co (IV) Ni (11) Nl (Ill)
114
Journal of Chemical Education
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Electronegatlvltles and Change with Unft Charge
s
AS,
electronegativity of thallium(1) were the same as for thallium(111). This causes the (I) oxide to be strongly basic instead of amphoteric or weakly basic as it would otherwise be. This is because the "inert pair" reduces the thallium electronegativity from 2.25 for the (111) state to 0.99 for the (I) state. An anomaly occurs with lead, in that its bonds to highly electroneeative elements aooear to favor the (11) state hut those to Lss electronegati;e elements favor the (IV) state. For example, PbCla tends to decompose readily to PbCb + Clz, whereas lead dialkyls are unstable with respect to the tetraalkyls and lead. I t seems quite certain that, if the electronegativity of lead were not appreciably lower in the (11) state, PbCL would he favored over PbCIz. On the other band, the &ectro~egativitydifference between lead (IV) and carbon is much smaller so that the polar component of the Ph-C bond strength " is much less sienificant. and four Pb-C bonds to the same lead atom, plus the formation of metallic lead. are favored over two Pb-C bonds to each of two lead atoms. The electronegativities of the lower oxidation states as given in Table 1 are not as reliable as those of the normal states because they are based on less accurate thermochemical data. Where all the normal valence electrons may he oresumed to be occuoied in the bonding, there is as yet no kvidence to suggest
