THE USE OF ELECTRONIC STRUCTURE IN INTERPRETING CHEMICAL REACTIONS WILLIAM F. KIEFFER The College of Wooster, Wooster, Ohio
T H E TOPIC of electronic configuration or electronic structure of atoms is invariably included in the early stages of the freshman's course in college chemistry. He is led more or less in detail past the milestones that science has passed on its quest for an understanding of why matter behaves as it does. He learns that the first signpost is the discovery that all matter contains the same negative units of electricity (cathode rays). This points the way to the eventual concepts of the nuclear atom. The way then leads to the Bohr theory which gives the chemist a woiking model of the atom. The student begins to feel the more familiar.ground of chernist,ry under his feet as he undertakes to apply the generalizations inherent in Lewis' concepts. Atoms exhibit predictable behavior in establishing eight electrans in their outermost energy levels. . Almost no college text and few high school texts allow the &dent to miss the opportunity of drawing miniature solar systems to represent the common elements. T h e adequate course will take care not to let the student. think of the atom as a.series of concentric railroad tracks with magically regulated numbers of electron trains. Rather it will present a t least a qualitative picture ,of the difficulty of locating any particular electron a t any particular spot, perhaps by introducing the concept of electron clouds or electron densities on a three-dimensional model. A good summarizing discussion of this problem has been presented by DeVault ( I ) . The detailed presentat,ion of atomic structure often demands too strict an adherence to some specific picture of the atom. Spectroscopy supplies the data which merely indicate the discreteness of the energy changes which accompany the shifting of electrona within atoms or molecules. The simple Bohr theory had to adopt elipses. Even thus modified, it yielded to the more inclusive vector model with its space quantization. Quantum mechanics discarded all conventional analogy and a t the same time employed a vocabulary of mathematical manipulations. Any complete treatment is virtually unintelligible and of questionable value to the beginning student. Regardless of the treatment, he does get a t least some qualitative concepts. He should recognize, for instance, t,hat the description of an electron's status within the atom demands several quantum numbers to adequately account for its behavior. Some texts supply him with a complete table of electron distributions, and a t least set the stage for the Pauli Exclusion Principle. He may be introduced to the peculiar jargon borrowed
from the spectroscopist ("8, p, d, f') used to designate varieties of electron habitation within any main energy level. "Valency" is probably already in the student's working chemical vocabulary and is the topic which almost always follows the section on atomic structure. The student immediately recognizes the logic of electrovalency in the light of electronic distributions. He sees consistency in the inertness of the helium family and the invariable monovalency of the alkali metals. Covalency mars somewhat the complacent feeling established by finding how simple it is to explain ionic reactions. Along with covalence comes a new rule for electron behavior, an almost invariable preference for pairing off. In fact, a bond now appears to be a shared pair of electrons. Coordinate covalence is a refinement introduced to explain the ability of such electronically satisfied molecules as ammonia to accommodate a wandering electronless proton which contributes no electrons, yet is bound into the ammoniumion. Here the whole problem of s t ~ c t u r eand valency rests for the student until an advanced course (usually physical chemistry) is reached or even until graduate school. He does, however, immediately begin to learn a lot of descriptive chemistry. Qualitative analysis, for instance, presents a baffiing array of selective reactions. Their interpretation and prediction are a challenge that the beginning student'can hardly accept, armed only with the electrovalency and covalency of the elementary course. The "rule of eight" of the Lewis concepts acts as a fairly adequate basis for the applications of electrovalency and many cases of simple covalency which he may be expected to recognize. It leads, however, to inconsisteiicies and contradictions as soon as the student encounters complex ions and other illustrations of what Sidgwick (8) calls the "coordination number" of an atom (the number of groups joined to a central atom by nonionized links). It can rightfully be argued that much of the discussion which follows has no place in the elementary textbook. There is, however, much to be said for making available to teachers and interested students a consolidated, if qralitative, presentation of some present structural concepts and to focus attention on certain features of atomic structure which offer explanation of chemical behavior, especially the phenomenon of complex ion formation. Chemistry teachers, need a reservoir of information greatly in excess of what they impart even to their superior students.
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It is just a step beyond the electron pair bond, invariably familiar to the elementary student, to think of the energy levels within the atom as being subdivided into "orbitals" which can be occupied by only two electrons. If. the idea of electron spin has been introduced, it can be seen that the restriction of allowing only electrons of opposite spin in ahy one orbital puts a consistent limit of two on the population of that orbital. Luder (8)has suggested a "Rule of Two" to emphasize the structural implications of this. Figure 1 represents the ordered arrangement of the electrons in atoms of the first four periods according to their orbital designation. It should be recognized that any atom will be expected to have its electrons arranged in such a way that the potential energy of the nucleus-electron system is a minimum. In other words, electrons will consistently fall into the available orbital nearest in energy to those already occupied. As a consequence of this, the following significant generalizations can be made, based on an observation of the table. (1) In any main energylevel, s orbitals are lowest in energy followed by p, d, etc., in order. (2) The d orbitals seem to be higher in energy than the s orbital of the following main level until that s sin b a r g y Leva1
Type of UrDltll tl
He
Li 0s
B C Y 0
F
no
F i g " . .
1.
Orbital Designation of El=trons
scandium, after the 4s is filled a t calcium.) However, the last d orbital to be entered may be even lower than the next s orbital. (Note the irregularities in the series a t chromium and copper.) (3) AU of the orbitals of any one type are of nearly the same energy and the normal tendency is partially to fl1 all before completely filling any. (The 3p orbitals a t phosphorous or the 3d orbitals a t chromium and manganese.) The above table can be extended by direct analogy to the following long period The pattern for families of elements is essentially the same with some minar irregularities. The period starting with cesium requires further extension. The rare earths fill seven 4f orbitals starting with cerium after the 6s has been filled a t barium and half of one 5d has been occupied a t lanthanum. Figure 2 (after Herzberg (4)) is useful in indicating the relative positions on an energy scale of the various orbitals as the atomic number increases. This is essentially the same as the detailed charts of DeVault ( I ) which show the actual energy positions plotted from spectral data. COMBINATIONS OF ORBITALS IN BONDING
It is a probable assumption that whatever order is apparent in the tendency to fill the orbitals of normal atoms may have a considerable selective influence on the filling of an atom's orbi6als by sharing electrons. with another atom in.. ckmical combination. It can also be argued that the availability of certain orbitals or combmations of orbitals may allow a sensible pr,ediction of the type and tenaency of chemical reactivity. (By "available" is meant the likelihood of an orbital being close on an energy level scheme to those already filled.) The fractional distillation of a mass of quantum-mechanical calculation yields certain very useful structural predictions. Pauling (5) has been able to shov that when certain orbitals are occupied by electrons shared between two atoms they establish themselves in a spatially symmetrical pattern. The "model" of an atom so involved would showthegreatestdeni n Atom.
OCTOBER, 1948
with information quite familiar to him. The example of chemical geometry invariably presented to the freshman is the tetrahedral carbon atom. In order for a carbon to combine covalently with four other atoms, it must share the population of four of its orbitals. Since the normal atom appears to have vacancies only in its p orbitals (see Figure I), the mechanism of its reaction must include the "promotion" of one of its s electrons into the third p orbital. This step is obviously endo-energetic, hut probably of relatively low requirement, since that p orbital is the next one occupied (in the normal nitrogen atom). The carbon atom so activated then is able to offer half an orbital to an electron from each of four other atoms to establish the hybrid tetrahedral bonding orbitals predicted by the quantum mechanics. The existence of the similar spatial symmetry of the NHa+ ion can also be seen as a direct consequence of the ability of the central nitrogen atom to hybridize its full s orbital with its three p orbitals. The latter are filled by ordinary convalencies with hydrogen atoms, the former is involved as the donor of an electron pair in a coordinate covalency with a proton. The identical nature of the four tetrahedral N-H bonds thus formed is emphasized. There is no Figure 2. Repleaentation of En.rgi.s of &bitel. with Change of Atomic Numb.. opportunity for the student to visualize one bond differ(4. attar H...b*r. (41 ing in strength from the others. Illustrative chemical reactions can be chosen particusities of valence electrons appearing as humps or horns larly to prove with apparent conclusiveness the sticking out from the sphere of the atom in specific accuracy of the theoretical concepts here outlined. directions. When one s and three p orbitals are simul- Many of these would undoubtedly be unfamiliar to the taneously involved, they scramble together ("hybridize" beginning student; nor would he encounter them unless according to Pauling) to produce four likely regions pursuing the subject in a speciali~edcourse or possibly for finding electrons which are pointed toward the in reference reading such as Chapter IV of Emeleus corners of a tetrahedron. Since the pairs of electrons and Anderson (7). This paper is intended to deal in such orbitals are involved in two atoms, the spatially with the elementary course. Accordingly, the follo\Ying oriented orbitals are essentially valence bonds. In a discussion is based on the application of these ideas to similar fashion, when one s, three p, and two d orbitals the reaction scheme used in the usual freshman qualitaare used, six octahedrally oriented hybrid orbital bonds tive analysis. It is hoped that it may make clear are produced. One s, two p, and one d flatten out some of the limitations imposed on 'chemical behavior into a square coplanar pattern. Rice 16) also lists a by what is understood to be the structure of the atoms linear combination for bonds formed by the hybridiza- involved. tion of an s orbital with one p. This last type is enSCHEME OF CATION QUALITATIVE ANALYSIS countered most frequently among those atoms such as In the scheme for qualitative analysis of cations the silver and mercury which have only an s orbital occupied on top of filled d orbitals. A second possible phenomenon of complex ion formation is repeatedly orbital combination can result in tetrahedrally oriented coordiition. According to Rice (6) one s, three p, and three or more d orbitals can hybridize to produce four bonding orbitals so arranged in space about the central atom. Examples of this type are rarely encountered except among the transitional elements in their higher oxidation states. Apparently not all of the hybrid orbitals are involved as bonds, yet the resulting complexes are very stable, for example, CrOaE. There are various other possible orbital combinations, but the above list, summarized in Figure 3, represent those most commonly encountered. Before making predictive application of these structural generalizations, it may be well to strengthen a student's belief in their existence by correlating them Figure 3. Spatially Odented Orbitd Combinatiolu
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up, it has not been necessary for the coordinating units to overcome any considerable electrical repulsion in order to enter the complex. The last S- attempting to enter the hypothetical Cd$-' complex would be kept from doing so by this repulsion. A second limitation on the coordinating tendency is the effect of atomic or ionic size and its relation to the nuclear charge. The structural significance of this has been pointed out by Campbell (8). Bismuth is a case in point. Although it has the same outer electronic arrangement as arsenic and antimony, it exhibits a preference for oxidation numbers +3, and generally electrovalent basic tendencies. The behavior of lead also offers illustration. The larger atoms have less sharply defined regions of electron density and consequently less tendency to retain shared bonding pairs. Of the cations in the f i s t group of the conventional analysis scheme the only one a t all likely to react by forming coordination complexes is the silver ion. Figure 5 indicates the electronic distributions. The silver ion is very willing to fill its open s a n d one of the open p orhitals which hybridize into the linear configuration. A great variety of electron-rich donors are acceptable to it. AgCl mill even dissolve in the presence of high chloride ion concentration to form AgCI2- in apparent violation of the solubility product principle. Obviously, AgCl is the only member of this group precipitate susceptible to the action of NH3 as a coordinating agent, hence aqueous NH, is a logical choice for the separating reagent. The peculiar ability of the mercurous ion to form the double ion appears to have an kxplanation consistent with the orbital point of view. TWOmercurous ions can fill their s orbitals covalently by each sharing the one remaining s electron in a common filled s orbital. This can be seen to be structurally similar to the familiar covalent bond of the, hydrogen molecule. The question may rightfully be-asked v h y zinc or cadmium make no such a n a l o g ou s monovalent double ion. The answer probably lies in the much great~r amount of shiekling afforded by the seven filled 4f orbitals in the mercurous ion. Ionization potential information supports this point of view (9). Even though zinc's and cadmiunl's electronic arrangemwt,~might allov the structural prediction of an s orbital bond, t,here would not be sufficient "insulation" between the positively charged nuclei to keep t,he r~sulting repulsion from Cations of tho Usual Qualitetive overcoming t,he bonding
used to advantage. In nearly every case the complex is formed by the mechanism of coordinate covalency. The cation loses its chemical identity because it becomes surrounded by other molecules or anions which are able to act as the donors in the linkage and supply it with pairs of electrons. If a cation is found to have available vacant orbitals which can hybridize to form one of the possible structures, it is logical to predict that a complex ion will be formed if the circumstances are right. Figure 4 locates the usual elements of the elementary qualitative analysis scheme and indicates the types of orbital hybridization most frequently encountered. This figure can be interpreted to advantage in terms of the electron distribution of Figure 1. For example, arsenic of family 5A, when displaying the oxidation number +5, can be considered as having emptied out its outer s and p orbitals. This gives it the same residual electron structure as Zn++. As the figure indicates, they can then be expected to exhibit the same coordination characteristics, as indeed they do in AsSF, and Zn(NH&++. Thus it can be seen that there is some predictable consistency in the number of units to be involved in complex formation. This is a most satisfactory discovery for the student who naturally looks for some regularity in chemical behavior to supplement the inadequate "rule of eight." Caution must be used, however, in attempting broad generalizations. The mere availability of an open orbital does not imply its invariable use for coordination. Several factors have limiting influence. One seems to he that very seldom will the charge of the resulting aggregate exceed -3. This is illustrated by the behavior of the Cd++ cation. CdS is not susceptible to solution in excess sulfide ion. Under these conditions S b S r is formed, but the hypothetical CdS4-6 is not, although it would employ identical orbitals. The structurally similar complexes Cd(CN)4= and Cd(l\THa)a++ are readily established in solution. For these last two to be huilt E
Figure 4.
Types of Covalent Orbitd Combinations Illustrated by th. Analysis Scheme
OCTOBER, 1948
., ! I
Orbitals
d
I
]
Exsllplsa P
I
Ccmplex forration ullllkdy
Ha (+I). Pb
Hgz
(+2)
Figure 5.
(+z)
m mmm full m m m m fun
Avdlable Orbit& In Group
KybridIcIng orbitnls not e-*llmble ror conp1e. forr~tlon
.
on.
C.tion.
force. The Hgz++aggregate is not very stable and is easily disrupted. As-might be expected, the electrons 6ll the orbital of one atom making metallic mercury and leave a vacant orbital in the other as the mercuric ion. The latter is then in a position to establish the same type of linear s-p hybrid as the silver ion. HgClz is a linear covalent molecule of extremely slight ionic character as might be expected from its electronic similarity to the very stable silver complexes. Group two of the cation analysis scheme employs a variety of reactions for separation and identification. Some of t,hese can be interpreted directly by structural considerations, others not so aptly. Figure 6 lists the cations (or more properly the atoms in their oxidation states), according to the usual coordination complexes which they may form. It should be noted that in addition to the indicated tetrahedral bonding, occasioually Sn (+4) is able to utilize two of its d orbitals for octahedral bonding, as in SnClam. In this case the d orbitals involved are of higher energy, being those occupied in normal atoms a t the beginning of the following transitional series. An analogous SbClr ion has likewise been suggested. The separation into subgroups by the action of excess suEde ion or the oxidizing polysulfide can be interpreted as the establishing of soluble tetrahedral complexes. The coordination of donor sulfides to the central As (+5), Sb (+5), and Sn.(+4) supply the open orbitals with the electrons needed. Apenic and antimony in the 4-3 oxidation stage seem to be able to accomplish the same type of reaction, probably forming the incomplete configurations represented by Ass3= and SbSae. Possibly the hybridization is almost the same except that the pair of electrons belonging to the s orbital of the As (+3) are utilized instead of a donated pair, as in the case of the completely coordinated AsS4'. If soluble bisulfide complexes are considered to be formed, the expected symmetry can be realized with As(SH)~-. This situation is not paralleled by the behavior of the tin ions. Although tin ($4) can form the tetrahedral complexes, Ens4-' or the incomplete SnSaq no such soluble complex exists for tin (+2). SnS is soluble only in the presence of the oxidizing polysulfide, whereas SnSz can be dissolved merely in excess of sulfide ion. Bismuth and lead likewise might be expected to form soluble sulfide complexes if they could be oxidized to their higher states in a similar fashion. This is virtually impossible to do in basic solution. (Latimer (10) calculates both to have
an oxidation potential of about - 1.6 volts.) Accordingly, bismuth, lead, and stannous sulfides are unaffected by treatment which dissolves the sulfides of arsenic, antimony, and tin. The amphoterism of stannous tin (and similarly lead (+2) ) cannot be explained if an analogy to sulfide complexes is to hold. Probably the relatively smaller and more electronegative oxygen is bonded by other than a simple C00~diIK%te covalence mechanism. Mercuric sulfide will dissolve by complex formation in a solution of very high sulfide ion concentration, i, e., NalS but not (NH4)2S. The linear HgSs- is the expected product of such reaction. Copper and cadmium do not share this reaction. As pointed out above, the requisite number of sulfides would not likely coordinate to build up such a great negative charge as would be involved in CdS4c8 or even the incomplete CdS3-'. Students frequently encounter trouble when using (NH4)2Sby having some of the CuS appear with the soluble arsenic group sulfides. Possibly some form of a mixed complex is the cause of the trouble. The readily formed amrqonia complexes of copper and cadmium are utilized to effect the separation of these cations from bismuth which cannot escape precipitation as the hydroxide or basic oxide by aqueous ammonia solution. The cupric complexes deserve a word of special attention. It might be expected that although the Cu++ has one half-filled d orbital behind its open s and p orbitals, it might use the latter four for tetrahedral hybridization as does Cd++. Direct structural observations on copper compounds have indicated their coplanar nature (11). Psuling (6) assigns a greater bond strength to tbe'dsp2 hybrids than to the spa type. He also points out that the energy to promote the electron into the nearest p orbital would not be a great requirement. It would not be expected to be as great as that required to unpair the two s electrons and promote one as was postulated for the tetrahedral carbon atom. The d orbital thus opened out could then be involved in the stronger square hybrid orbitals. In the presence of the reducing action of CN- ions, the copper retains its two d electrons in the last orbital.
Fi-
6. A l a i k h l * Orbital. in Group Two Catiolu
JOURNAL OF CHEMICAL EDUCATION
The resulting Cu+ has only s and p orbitals available, which it hybridizes into the Cu(CN)F tetrahedral complex. As long as no oxidizing agent is present, this is a very stable complex. It is specifically more stable than the corresponding Cd(CN)r, since the latter surrenders to H2S as CdS making the identification of cadmium in the presence of copper possible. The successful manipulation of Group 3 of the cation analysis presents a problem which prompts nearly every textbook author to suggest his own particular scheme of reactions. It is beyond both the scope and the time allowed this discussion to treat all of these in detail. The following suggestions are some of the pertinent generalizations that can be made.
orbitals and the 4 s are of about the same energy, there may be a reshuffling as indicated by the dotted lines in Figure 7. If it is borne in mind that the resulting complex is essentially covalent, filling hybridized d and s orbitals anyway, it is of less consequence definitely to assign the + 2 and +3 oxidation state an electron distribution. The question of course arises as to why cobalt ($3) or nickel (+2) mill accept the neutral yet essentially electronegative NH, molecule to form complexes while ferrous and ferric apparently are unwilling to do so. The latter will form ccmplexes with the less electronegative CN-. The 'answer is not at all clear. It must be remembered that in all probability these ions are "aquated," (CO(H20)6+++,etc.). The forming of any other complex implies an exchange of coordinating units. The situation probably involves much more than the simple structures here represented. Double bonds utilizing other d orbitals of the iron as mell as electrostatic ionic bonds probably have to be considered as contributing to the stability of the ferric and ferrocyanide structures (5). A second difference introduced by consideration of these ions is that the completion of a coordinatin~ complex may have a stabilizing influence for one oxidation state over another (12,.15): Cobalt is the most frequently cited example. of this, yet it is often overlooked by the authors of analysis textbooks. Cobaltous (+2) ion contains one more electron than cobakic (+3) ion as diagrammed in Figure 7. In order for the two requisite d orbitals to be open for coordination, this electron mould have to be promoted into a d orbital belonging to the next higher main energy level. (The nearby s and p orbitals must Be left vacant for the coordination process.) This promotion would require Figure 7 shows that all of the cations here included much more energy than either of the previously noted have the ability to coordinate complexes of some type. promotions in carbon or copper. Such a situation Thisis to be expected from their positionsin the periodic makes any so-formed Co(NH2)6c4complex very willing table (Figure 4). Their electronic structure allows to lose the single promoted electron:, Evidence of this is for many vacant orbitals which canebe hybridized the fact that C O ( N H ~ )is~ a- ~good reducing agent and into bonding possibilities. Inspection of the figure is itself readily oxidized. Latimer (10) tabulates along with any outline of analytical procedure mill -0.1 volt as the potential for the formation of the cobaltic complex from the corresponding cobaltous. suggest possible interpretations. One new consideration which must be allowed from CO(NH&~ could hardly be expected to ehist in the an examination of these examples is the way in which presence of air or other oxidizing agents. Another the electrons of incompletely filled d orbitals are appar- interesting valence stabilization is that afforded the ently able to shift and pair up so that two d orbitals will higher states of chromium (+6) and manganese (+7) be available for octahedral hybridization. Iron, cobalt, by the highly negative 0- coordinating unit. Mention should be made of the possible interpretaand occasionally manganese follow this procedure. Nickel performs the same type of internal rearrange- tion of amphoterism. The Bronsted concepts of acid ment to open the way for its very stable coplanar and base are becoming more and more videly accepted in the elementary course. These definitions can be square complexes. Some question may arise as to how to indicate cor- utilized in a manner consistent with predictions based rectly the electron distribution in the +2 and +3 oxida- on structure. See, for example, Hogness and Johnson tion states for the elements chromium, manganese, (14). The following equations illustrate how the iron, cobalt, and nickel. It may be argued that since amphoterism of zinc hydroxide can be considered as the 4s orbital was filled a t calcium, and the subsequently proton exchange between the coordinated units (either added electrons are entering d orbitals across thetransi- OH- or HzO), of the zinc complex and water, hydroxyl, tion series, these d electrons would be lost in ionization or hydronium ions in the solution. rather than the 4 s electrons. If, however, all the 3d Zn(H*O)4++ 2H20 2 R O + + Zn(H,0)2(0H). solid
+
-
OCTOBER, 1948
pair of electrons from each of these nitrogens is available to fill the hybrid spd2orbitals of the Ni++. The behavior of aluminum hydroxide can be interpreted by a similar set of equations involving the tetrahedral complexes, AI(OH)*', AI(OH)I(H,O), and Al(H,O)a+++. Some authors assign a coordination number of six to the aluminum for analogous equations (15,16). This is possible, but would require the aluminum to involve its d orbitals of the same principal quantum number as its available s and p orbitals. These d orbitals would be expected to be of somewhat higher energy. This situation is analogous to that mentioned ahove for the octahedral complexes of Sn (+4). The hydrated ion AI(H20)s+++ is recognized as an octahedral complex in crystals and aqueous solutions. It is uncertain, however, whether it persists in such orientation when exchanging H2O for OH- as postulated ahove. It is also likely that its size and general electropositive tendencies would favor establishing of bonds of ionic character rather than pure covalencies. The dissolving of chromic hydroxide in strongly alkaline solution might be explained by a similar argument if it is truly evidence of amphoteri~mand not colloidal peptisation. The interference of various anions with the successful analysis of this group of cations is another illustration of the ease with which they form coordinated stable complexes. Ferric, aluminum, andchromium especially are capable of forming fluoride, oxalate, and often borate complexes which completely block the ordinary analytical reactions. The action of the oxalate ion is of special interest since it will arrange itself to accept two of the coordination locations about the central cation. This is an example of the widely encountered phenomenon of "chelation" (17). The term implies the cyclic nature of the coordinating unit, like the hook of a lobster's-claw. Trivalent cations with available octahedral orbitals become surroundkd by three oxalates so that each end of the latter is involved with one of the hybrid orbitals. The resulting complex, e. q., F e (C?01)3-, is thus electronically and structurally the equivalent of any other octahedral arrangement of six single coordinating units, e. q., the interfering FeFsZ. In the case of the oxalate, both of the anionic ends (-COO-) of the molecule are involved as donors of electron pairs in the coordination with the central atom. As Diehl (17) points out, this is actually a special case of the chelation phenomenon. Often a pair of electrons in a "neutral" portion of the attaching molecule is involved. This usually occurs with molecules having a conveniently placed nitrogen atom. It is illustrated by the action of the familiar and invariably used specific reagent for nickel, dimethyl glyoxime, which assumes a cyclic arrangement about the central Ni++. The diagram indicates that each nitrogen is located a t the corners of a square with the nickel in the center. A
H,C-$----C-CH,
0-h \
fi-0
,,'
,'
Ni
H'
\
O-N
:'
11
\
N=O'
Many other specific reagents have been discovered among similar organic molecules capable of forming resonating complexes by chelation. A-nitroso-p-naphtho1 for cobalt and cupferron for iron are frequently employed in this way. CONCLUSION
It is hoped that the foregoing discussion has indicated one more useful angle of approach to our problems of visualizing and recognizing consistencies in chemical behavior. The "orbital" is not intended to be a startling structural implication, but merely a convenient vocabulary into which to translate electronic structure for interpretative purposes. Nor is it to be used as a magical work intended to make possible a correlation of all known reactions with the spectroscopist's electron designations. The fact that it, does allow new consistencies to become apparent, though, does suggest its value as a means of analyzing many chemical phenomena left a puzzle by other representations of atomic and molecular structure. LITERATURE CITED (1) DEVAULT, D., J. CHEM.EDUC.,21,626,575 (1944). K, "The Electronic Theory of Valency," (2) S ~ G W I C N.,V., Oxford University Press, London, 1927, p. 110 ff. (3) LUDER, W. F., J . &EM. EDUC., 22, 221 (1945). (4) HERZBERG, G., "Atomic Spectra and Atomic Structure," Prentice-Hall, Inc., New York, 1937, p. 148. L., "The Nature of the Chemical Bond," Cornell (5) PAULING, University P r e q Ithaca, New York, 1940, pp. 81-106. (6) RICE,0. K., "Electronic Structure and Chemical Binding,'' McGraw-Hill Book Company, Inc., Nev York, 1940, p. 267 ff. (7) EMELEUS, H. S., AND J. S. ANDERSON, "Modern &peck of Inorganic Chemistry," D. Van Nostrand Co., Ine., New York, 1938. (8) CAMPBELL, J. A., J. CAEM.EDUC..23. 525 (19461: 25. 200 ~
(1948). H. H., AND C. A. VANDERWERF, ibid.,22,390 (1945). (9) SISLER, (10) LATIMER,W. M., "Oxidation Potentials," Prentice-Hall, Inc., New York, 1938. (11) Cox, E. G., AND K. C. WEBSTER, J. Chem. Soc., 731, (1935). (12) COPLEY, M. J., L. S. FOSTER,AND J. C. BNLAR,JR.,Chem. Rev., 30, 227 (1942). (13) BLANCHARD, A. A,, J. CHEM.EDUC.,20,454 (1943). (14) HOGNESS, T. R., AND W. C. JOIINSON, "Quditative Analysis and Chemical Equilibrium," 2d ed., Henry Holt & Co., New York, 1940, p. 231 ff. (15) MIDDLETON, A. R., AND J. W. WILLARD, "Semimicro Qualittttive Analysk," Prentioe-Hall, Inc., New York, 1940, n r.
227 .
(16) WILDMAN, E. A., J . CEEM.EDUC.,12, 11 (1935). (17) DIEAL,H., Chem. Rev., 21, 39 (1937).
