THECOORDINATEBOND ANDTHENATURE OF COMPLEX INORGANIC COMPOUNDS'
0
I. The Formation of Single Covalent Bonds DARYLE H. BUSCH The Ohio State University, Columbus, Ohio
pedagogy of complex inorganic compounds is all too often quite haphazard both a t the undergraduate level and in more advanced classes. At the student's first encounter with complexes, he is usually told about the coordinate covalent bond as an electron-pair bond and then given a few examples to illustrate its particular nature. He then hears little about these substances except perhaps in conjunction with qualitative analysis where they are treated mainly in discussions of equilibria. At the advanced level, the stereochemistry of complex inorganic compounds is usually presented effectively in terms of the atomic orbital theory (as that theory was summarized by Pauling several years ago). Several excellent reference sources exist to equip the teacher for his lectures on stereochemistry. However, the teacher does not have a t his disposal similarly adeqnate sources which provide an integrated picture of the bond types encountered in these substances and serve to account for the existence of complex compounds of greatly varied properties. I n consequence, such topics as bond type, stability, and relative reactivity are often inadeqnately explained or not explained a t all in the classroom. The author feels that a higher degree of order exists among these compounds than is apparent from the secondary literature references most often consulted by teacher and student. Even before the inception of modern valence theory, it was realized that complex inorganic compounds posed a special problem. The very nature of these compounds would prevent their being anticipated in terms of the most general theories of valence, for they are usually formed as a result of the combining of molecules or the ions of salts which are capable of independent existence, as is shown by the equation: THE
NiCL
+ 6NHa
-
Ni(NH3)~Clz
To the early investigators, such behavior presented troublesome exceptions to the otherwise consistent combining capacities of the atoms involved. Obviously, in the case just cited, either the nickel or chlorine atom has assumed a new valence (other than that encountered in simple salts) and also either the nitrogen or the hydrogen has assumed a new valence. 1 Presented before the Division of Chemical Education a t the 128th Meeting of the American Chemical Society, Minneapolis, September, 1955.
The first successful correlation of the early accumulation of data on complex inorganic compounds appears in the brilliant deductions of Alfred Werner. Werner pointed out the great tendency of most metal atoms or ions to combine with certain numbers of other atoms, molecules, or ions, and he called the number of such groups the coordination number. He suggested that the manner of binding between the entities in question was different from that in simple compounds and he spoke of the bonds as involving a "secondary" valence for the metals, in contrast to the primary valence, which is now recognized as the electrovalence in the case of salts or the covalence in the case of nonionic, molecular substances. During the fist three decades of this century, an electronic picture of the coordinate bond was developed. The nonmetallic parts of the complex compounds, so intently studied by Werner, contain u n s h e d electron pairs in their valence shells, and the metal atoms are in almost all instances in positive oxidation states. This led to the idea that the coordinate bond is a covalent bond in which the nonmetallic part (the donor group or ligand) donates a pair of electrons which is then shared by both the donor atom and the metal atom (the acceptor atom). This simple theory represented a major advance in accounting for the existence and general properties of complex compounds. However, the concept was much too primitive t o account for the variations in the behaviors of the many different complexes known. It has long been realized that on the basis of chemical behavior all complex compounds may be placed into two broad categories: those which are in rapid equilibrium with their component parts in solution or in heterogeneous systems, and those which are only very slightly dissociated or which undergo no appreciable dissociation. These two classes are called "normal complexes" and "penetration complexes," respectively. Concurrent with the development of the covalent theory of the coordinate bond, an electrostatic theory evolved. According to the electrostatic theory, the attraction between metal ions and polar molecules, such as water and ammonia, results from coulombic forces. Because of the rapid equilibria involved in reactions of electrosubstances, yalent substances as ~twas logical to suppose that the "normal complexes"
VOLUME 33, NO. 8, AUGUST, 1956
were more likely to be bound by electrostatic forces, while the "penetration complexes" were then believed to be covalently bound. One of the most intriguing questions remaining unanswered is whether the differences in these two major types of complexes are a function of relative stability, bond type, or rate a t which the complexes may react, or a combination of these effects. An attempt a t approximating an answer t o this question may be found in the discussion t,o follow. THE RESULTS OF THE ATOMIC ORBITAL THEORY
Although the relative success of the two major schools of modern valence theory (molecular orbital and atomic orbital) may well be a matter of opinion, the atomic orbital theory has been more widely applied to the problem of the coordinate bond. Therefore, the use of the atomic orbital theory is chosen here, except in the consideration of double-bonding (second paper of this series) where the use of molecular orbital terminology facilitates discussion. Extensive review of the literature is not intended and for the sake of comparison or for review of some of the concepts used, the reader is referred to a number of recent reviews (1-4) and to texts treating modern valence theory. In order that the necessary conceptual tools be a t hand, a number of familiar consequences of the atomic orbital theory must be mentioned. Hybridized Bond Orbitals. The study of the chemical and physical properties of compounds revealed that many atoms have four or six equivalent bonds in their structures despite the fact that the earlier study of atomic spectra indicated that the same atoms have only three equivalent orbitals in their valence electron shells. Pauling cited the example of carbon which is known to form four equivalent bonds directed toward the corners of a tetrahedron. The valence shell of carbon, on the other hand, is made up of one sand three p orbitals, the three p orbitals being of equivalent euergy. On the basis of these available orbitals, carbon would be expected to form three equivalent bonds and possibly a fourth weaker bond. However, quantum mechanical theory has shown that the combining of one s orbital with three p orbitals may result in four equivalent bond orbitals, each having greater potential bond-forming ability than either a single s or p orbital. One of such a reorganized set of bonding orbitals is called a hybrid orbital. Hybridization has been shown to result in many possible combinations of bond orbitals. The relative bond-forming potentialities of these hybridized orbitals have been estimated by means of a parameter called the "angular strength" (9). The angular strength is a measure of the concentration of the particular hybrid orbitalabout the bond axis, and a relatively great bond-forming ability is indicated by a large value for the angular strength. These and the common types of hybrid orbitals are summarized in Table 1. The Magnetic Criterion for Bond Type. I n many cases, an atom may make use of the necessary orbitals
TABLE 1 Angular Strengths of Bond Hybrids (3, 6) Hybrid'
Configuration
ns npa Tetrahedral ( n - 1)d ns npa Square planar ( n - 1)d'nsnpa Octahedral ns npa nd* Octahedral n denotes the major quantum number of shell.
Angular strength 2.000 2.694 2.923 ? the valence electron
to form a certain set of hybridized orbitals with no complicating rearrangement of its underlying electron shells. This is usually true when the bond hybrid is ns npa or ns np3 nd2, and occasionally true when the bond hybrid to be formed is (n - 1)d ns np2 or (n - 1)d2ns np8. The reasons why rearrangements of electrons are sometimes necessary in order to form the bond hybrids using d orbitals of the penultimate shell (major quantum number of n - 1) are: (I) the (n - l)d orbitals are, in the case of transition elements, partially filled; and (2) the lowest energy states of transition element atoms are attained when, as the Hund rule of maximum multiplicity states, each of the five d orbitals receives one electron before any of the d orbitals receives its second electron. The result is that no d orhital is varunt und rherrfore n\.nilnt~lefor the formation of hylxi~lizedt ~ u t l t lorhitals if a total of mulv as tive electrons are present in the d subshell, and only one d orbital is vacant if there are four electrons in the d orbitals. If the number of these d electrons is between four and six, inclusive, two d orbitals may be made available for the formation of (n - l)d2ns npa hybridized orbitals by pairing a sufficient number of the electrons to place all of them in three of the d orbitals. If the number of d electrons is between five and eight, inclusive, one d orbital may be made available for the formation of (n - l ) d ns npZhybridized orbitals by pairing a sufficient number of these electrons to place all of them in four of the d orbitals. (See "ground state" examples in Table 5.) The most readily measurable result of these rearrangements of electrons is the number of unpaired electrons (before and after rearrangement). As has been discussed in many places in the literature ( e . g., reference (Z), page 208), each unpaired electron acts as a tiny electromagnet and by suitable techniques the number of unpaired electrons in each atom may be determined from the magnetic behavior of the substance in question. Pauling recognized the significance of the alteration of the number of unpaired electrons in metal atoms which are parts of complex compounds as an indication of the type of hybridized orbitals involved. He also associated the formation of a complex utilizing (n - l)d2 ns np3 or (n - l)d ns np2 hybridized orbitals with the presence of a relatively high percentage of covalent character in the bonds. From these considerations he formulated the "magnetic criterion of bond type" in which he chose to refer to complexes containing metal atoms that have undergone changes in the num-
378
her of unpaired electrons as "essentially covalent," while referring to those complexes in which no such change in magnetic moment (if such a change would he necessary in order to empty some of the orbitals for use) occurs as "essentially ionic." It was admitted that the criterion would not allow a distinction in all cases, e. g., chromium(III), as discussed below. Ionic and Covalent Character in the Coordinale Bond. I n the atomic orbital picture of polar bonds, or bonds of partial ionic character, the actual bond is described as something intermediate between a purely covalent bond and a purely electrostatic, or ionic, bond. Thus, the polar bond in the molecule AB is intermediate hetween the structures:
JOURNAL OF CHEMICAL EDUCATION
only between structures having the same number of unpaired electrons (6). This point has led some authors to suggest that no ionic contribution is involved in the bonds of diamagnetic (and similar) complexes (2, 5). This is equivalent to saying that the bonds can have no polar character, a highly unlikely situation. Pauling (7) was explicit in suggesting that these bonds could be polar in character; however, he provided no explanation of how this might occur. I n spite of the statement by Pauling that "essentially ionic" was intended to tag those species of relatively high polar character in their bond type and not necessarily to denote that the species were bound by electrostatic means alone (6), some literal use of the terms has occurred ( I S ) . Recently the magnetic criterion has been used to distinguish between octahedral complexes having d2sp3 hybridized orbitals involving the d orbitals of the pennltimate electron shell and sp3dZhybridized orbitals involving d orbitals of the valence shell. For those metal ions where a distinction is possible, the octahedral complexes having the same number of unpaired electrons (the same magnetic moments) as the corresponding gaseous ions in their ground states have been termed "outer-orbital" complexes, while those having magnetic moments differing from those of the ground states of the gaseous ions are called "inner-orbital" complexes (8). The same language is also in use to distinguish between dsp2 and spa hybrids in four-coordinate complexes. This classification includes among the innerorbital complexes those species in which d orbitals of the penultimate shell are logically expected to participate in the formation of hybridized orbitals even though no change in magnetic.moment is likely. It is in this respect that the scheme enjoys its greatest generality. Perhaps the most notable remaining consequence of this classification, other than avoiding the confusion occurring as a result of Pauling's unfortunate terminology, is the increasing realization that d orbitals of the valence shell must come into play in the formation of many octahedral complexes. Recent calculations have confirmed the likelihood of this type of hybridization (9-11). It is certainly unlikely that complexes, such as Ni (dipy),++,Ni(0-phen)3++,Ga(CzO4)a-~, and Ge(C20&-(dipy represents a&-dipyridyl and o-phen represents o-phenanthroline) which have been resolved into optical isomers (I%),are bound by use of only four spahybridized orbitals as was originally suggested by Pauling. Nonetheless, it should be realized that in the most literal sense the classification of complexes as "inner orbital" and "outer orbital" is almost as phenomenological as, though certainly more general than, the terms "diamagnetic" and "paramagnetic" (which indicate whether or not any electrons remain unpaired), since i t denotes no distinction in bond type but is intended merely to identify the nature of the hybridized orbitals.
where the atom B is the more electronegative of the two. The relative polarity of such a bond depends on the relative importance of structures I1 and I in determining the true structure of the molecule. (The true structure of the molecule in quantum mechanical terms may he approximated by a total wave function, summing up the wave functions of the contributing structures, each of which is multiplied by a weighting factor.) This mixing of limiting types of structures to describe an intermediate structure is called resonance. If for a pair of atoms, such as AR above, certain parameters may he wried, specifically theelectronegativity of each atom, then a gradual change in the degree of polarity, or ionic character, of the bonds may be realized. I n reality such changes are accomplished by varying the nature of the atoms or by varying suhstitnent groups on the atoms in question. The realization that differences occur in the degree of polarity, or of ionic character, in the bonds in various complexes preceded and coexists with the atomic orbital theory. However, alterations in magnetic moments apparently are associated with a discontinuity in the gradual change described above, for the compounds involving metal ions having reduced magnetic moments (fewer unpaired electrons than the corresponding simple ions) are logically considered to have relatively more covalent character, as shown by their decreased susceptibility toward rapid dissociation or substitution reactions. The relatively smaller polar character in the bonds of these complexes can be related to a decrease in the stability of the structures which define the ionic contribution to the bonds. This is made clear in the diagram by Pitzer (5) which illustrates that the ionic contribution to a bond in diamagnetic nickel(I1) involves a relatively unstable, excited state for the gaseous ion, while the covalent contribution involves a favored electronic configuration of the species and a bond hybrid of maximum angular strength. The same situation occurs for any ion which behaves in accord with the magnetic THE RELATIONSHIP OF BOND TYPE TO STABILITY criterion. The assumption that ionic resonance strucA complete treatment of the stabilities of complex tures contributing to the bond types of complexes of this class must involve excited states of the metal ion is compounds would require analysis of the effects of all necessitated by the restriction that resonance can occur the thermodynamic terms involved in a complete forma-
VOLUME 33, NO. 8, AUGUST, 1956
379
tion cycle, such as that shown below, and analysis of the factors determining the magnitude of each of these terms. I n this cycle, the heat of formation of the solid complex of a metal, M, with n moles of a ligand, L, in the form of a halide salt, X, is denoted by the symbol Q. As is seen in the diagram, the over-all formation of the complex is envisioned in one step in the reaction written vertically a t the left. The formation of this same substance is also envisioned by a series of steps, proceeding in a clockwise manner from the upper left-hand corner of the diagram and ending a t the lower left. This second reaction path is utilized in order to show how the heat of formation, Q, might be calculated by an indirect method and also to emphasize the several energy terms which may be considered to effect the magnitude of the heat of formation. The longer path imagines the formation of the complex as proceeding through four major steps (each of which is identified with one of the arrows shown in the cycle): (1) conversion of the solid metal and ligand into gaseous atoms or molecules, and dissociation of the diatomic halogen molecule into single atoms; (2) removal of two electrons from the metal and addition of one electron to each halogen atom, thus producing gaseous ions; (3) union of the gaseous metal ion with n moles of the ligand to form the gaseous complex ion (this term is the total heat of bonding for the formation of n bonds): (4) of the com. . crystallization plex cation and the halide anion to form the solid complex. The heat of formation is the algebraic sum of the heat quantities associated with each of these steps. The anantities which are written with neeative siens in the formation cycle are exothermic, while those written with positive signs are endothermic. The algebraic sign of the heat ~Fformationis, of course, dependent upon the relative magnitudes of the sums of the endothermic and exothermic terms.
-
-
Tharmochemicd Cycle for the F ~ m a t i o nof s Solid Complex Involving a Divdmne Metal Ion and a Halid* Anion +Q
U - F-&
-A
=
Su + nSL + Dx. + I ,
+
-
- 2EX - A
= energy of crystallization = .lo,.tmn affini+rr nf ".*-Y = enFrgy of bonding of gaseous ligand t o
"."""."
"J
gaseous metal
,on = sublimation energyaf metal = sublimation energy of n moles of ligand
SM nSr.
Examples can be cited where the dominant effect of each of these terms causes deviations from exnected behavior. For this reason it is essential that the relationship of the term upon which an expectation or premise is based to the remaining terms of the cycle be fully ~
~
~
~
~~~
~~
realized. I n discussing bond type and the stability of complex compounds, the nature of the hybridized bond orbitals has come to occupy a position of primary concern to chemists (1, 8, I S , 14). Obviously, any predictions based on hybridized orbitals alone ignore all the terms in the formation cycle except the heat of bonding of the gaseous metal ion with the gaseous ligand ( - A in the cycle above). For a series of closely related complex compounds, the terms II,I*, nSL,Dx,, -2Ex, and - U may be considered to be constant or their individual effects may be anticipated. For example, the relationship of ionization potential to bond type is often considered indirectly, since it is of primary significance in determining the electronegativity of the metal ion, which in turn leads to inference with regard t o bond type. This is discussed below. 3d
I-[
4s
4p
0
rn
bonding orbitals
L '
Structure of solid complex
Electronic configuration of metal in solid complex
O =-/-rr
..
H II I I
I
n
bonding orbitals
-
Electronic configuration of metal in complex in nonpolar solvents
Structure in nonpolar solvents ~h. strust-
of ~i~(di~laldimi~.)~ick~l(~~) in solution Solid State a Infemd horn Mwnetic Maw-men*
in th.
The fact remains that failure t o consider the possible intervention of any single energy term may lead to confusion. An interesting example is found in the magnetic behavior of the complex compound bis(sa1icyladimine)nickel(II) (see stmcture above). This substance is diamagnetic (implying completely paired electrons) and planar in the solid state; however, it exhibits naramaenetism in solution. even in non~olar solvent's such as benzene, and is therefore tetrahkdral under those conditions. This behavior probably results from a greater heat of bonding for the tetrahedral (paramagnetic) form (the absolute value of A is greater for the tetrahedral form) which is counterbalanced in the solid state by the considerably greater lattice energy of the planar form (15,16). I n the majority of discussions concerned with the structure, stability, and reactivity of complex inorganic compounds, primary concern is placed on the factors which determine the maenitude of the term A. These factors are listed in Table 2. Other factors which affect the solution behavior of the complex, such as the effect of chelation, have also received considerable at-
-
-
380
tention in recent years. These will not be considered here, but may be found by the reader in other reviews (2). Most of these factors are discussed by Nyholm (1) or by Martell and Calvin (Z), and some of them are restricted to significant but definitely limited and easily recognized groups of complexes. The intent here is not to treat exhaustively all the factors influencing the b e havior of complexes but rather to consider the most general phenomena and to point out and discuss a few factors which are sometimes ignored or poorly understood. The first of the factors listed in Table 2, the bond hybrid of the metal, has already been discussed and will be considered further as the discussion proceeds, as will the second, third, and fifth factors (numbered as they appear in the table). The following comments will serve to indicate roughly the effects of the remaining factors on the heat of bonding. The Effect of the Electronic Structure of the Ligand on the Heat of Bonding. In some cases the formation of complex compounds contributes materially to the conjugation of unsaturated donor molecules. When this occurs, the stability of the complex is often considerably greater than would be anticipated on the basis of other contributing factors. This effect is sometimes associated with the occurrence of double bonds between the ligand and the metal atom and it is discussed in that connection in the second paper of this series. A more definitive treatment of this phenomenon may be found on page 161 in reference (2). The Steric Requirements of the Ligand and of the Metal Ion in a Given Bonding State. The stability of a complex may be affected greatly by the relative spatial requirements of the ligand and of the metal ion. This is illustrated by the complex of copper(I1) with the base &p',O1'-triamiuotriethylamine (structure 111). This copper complex is less stable than the corresponding copper(I1) complex with triethylenetetwmine (stmcture IV).
The logarithm of the formation constant in the former complex is 18.8 (17),while that of the latter is 20.5 (18). I n contrast to this behavior, nickel (11) forms complexes of very similar stability with both amines, i. e., log K with compound I11 is 14.0 (17) and with compound IV it is 14.1 (18). The unusual stability order with copper(I1) is attributed to its great tendency to form square planar complexes. The branch-chained polyamine shown in (111) cannot occupy the four corners of a square plane but is supposed to cause the four bonds to be forced into a tetrahedral configuration. On the other hand, the linear polyamine of (IV) can array its four nitrogen atoms about the square plane. I n con-
JOURNAL OF CHEMICAL EDUCATION
sequence, the planar complex is the more stable. This example illustrates the effect of incompatible steric requirements on the parts of the ligand and metal atom. Presumably, it is entirely possible that the ligand and metal atoms might have ideally similar steric requirements, leading to complexes of unusually great stability. Perhaps the complex, bis(acety1acetonal)ethylenediiminecopper (11) (structure V) provides such an example, for the high degree of conjugation in the organic ligand should force its atoms to be coplanar and the copper complex is reported to be stable toward decomposition even when heated to redness (19).
Charge on the Complex Ion and Net Charge on the Central Metal Ion (The Criterion of Essential Electroneutrality). Pauling (20) has suggested that it is unreasonable to suppose that any single atom in a complex ion or molecule is the seat of a large amount of charge, but that the charge on a complex ion will be distributed over all the atoms in the ion or molecule. Thus, in ferrocyanide, [Fe(CN)a]-4,on the assumption that the iron atom was originally dipositive, if the iron atom gains an equal share in six pairs of electrons i t will have a net charge of -4 (one-half of six pairs of electrons, or six electrons, minus the two positive charges originally present on the iron). The minus four charge is probably removed from the iron atom either by the displacement of the pair of electrons of the coordinate bond back toward the donor atom of the cyanide group, thus giving rise to a polarity in the bond, or by donation of some of the previously unused electrons of the iron atom to the ligand, thus giving rise to a double bond (this is discussed in greater detail under the heading of double bonding in the second paper of this series). Base Strength of the Donor Atom. All other factors being equal, the formation of a donor-acceptor bond in a complex compound would be expected to depend on the relative effectiveness of the ligand as an electronpair donor. This donor power is measured by the base strength of the ligand (1,s). According to Nyholm (I), the bond strength for the metal-ligand link is the mean value per ligand of the heat evolved by the reaction of the simple gaseous ion with the gaseous ligand to produce the gaseous complex. This value is useful in considering the fundamental properties contributing to bond type and stability; however, the total heat of bonding to form the number of bonds equal to the coordination number of the metal (or - A in the cycle) is used here.
VOLUME 33, NO. 8, AUGUST, 1956 THE RELATIONSHIP OF SINGLE COVALENT-BOND CHARACTER TO STABILITY
In general, the polarity of a covalent bond can be given qualitative consideration on the basis of the relative attractions for electrons of the two bonded atoms. This electron-attracting power of atoms has been estimated by several investigators in terms of a parameter called the electronegativity ($1). This parameter has been given three different definitions. These are summarized and explained in the reference cited. Qualitatively, it remains the intrinsic attraction of an atom or ion for electrons. According to Pauling (13) the ionic (or polar) character of a bond increases as the electronegativity difference between the two atoms increases. This concept is quite satisfactory as a qualitative basis for it is exactly what one would anticipate on the basis of the relative attractions of the two bonded atoms for electrons. Mulliken ($2) defined electronegativity as the average of the ionization potential and electron affinity, and his definition is theoretically sound. It suffers the disadvantage that electron affinity data are not abundant. However, in the case of metals it is reasonable to assume that the electron affinity may be ignored, whereupon the ionization potential becomes an estimate of the electronegativity (1). This expedient is followed here. The word stability is often misused to denote reluctance of a substance to undergo chemical reaction. Since this term is most significant when it refers tathermodynamic stability its use is so restricted here. The measures of thermodynamic stability available for complex inorganic compounds are not plentiful. They are the heats of formation of ions and compounds, and formation constants derived from equilibrium measurements. The formation constant, it should be realized, is not an absolute measure of stability, but a measure of the relative stability of some complex, [ML,]+"', over the corresponding hydrate, [M(H20).]+" (if the solvent is water), since it is related to the free energy of the reaction: [M(H,O),l+"'
+ pL s [ML,ltm + nH*O
TABLE 2 Factors Which Determine the Magnitude of the Heat of Bondino- .( A. (1) The bond hybrid of the metal atom. (2) The relative electronegativities of the donor and acceptor atoms. (3) The underlying electronic structure of the metal atom in its bonded state. (4) Tho electronic structure of the ligand (possible conjugation, etr
(5) ~ h e - & s i b i l i t ~of double-bonding between the metd atom and the lieand fa. eanseauenee of the four oreviouslv mentioned factorsj. (6) Sterio requirements of the ligand and of the metal ion in a given bonding state. (7) Charge an the complex ion and net charge an the central metal ion (Pauling's criterionof essential electroneutrality). (8) Base strength or proton affinity of the donor atom.
tronegativity difference between the oxygen donor atom and the metal acceptor atom is smaller with nickel(I1) ion. This smaller electronegativity difference leads to the inference that the nickel(I1)-oxygen bond has a higher per cent covalent character than the manganese(I1)-oxygen bond. I n the second column in Table 3, the heats of hydration for these ions are given. Here it is seen that the heat of hydration of the nickel(11) ion is greater than that of the manganese(I1) ion. Thus it is seen that the more stable complex ion probably has the greater amount of covalent character. Since the electronegativity of nitrogen in ammonia is less than that of oxygen in water (IS), it is also consistent with this view that the ammines are more stable than the hydrates in the case of the metals of relatively great electronegativity (Table 3). The increase in stability of complexes involving uncharged base molecules as ligands with an increase in covalent character of the bond may be explained on the basis that the limiting electrostatic type of bond would involve a relatively weak ion-dipole interaction. I n the case of charged ligands, so simple an argument probably is not adequate. This view is supported by the investigations of Cottrell and Sutton (SO). As shown above, when the coordinate covalent bond is simply a single electron-pair bond, the covalent nature and the stability of the bond should increase
.
Agreement of ionization potential with such measures of stability for complexes of the divalent metal ions of TABLE 3 the first transition series as AH of hydration and the Relationship Between Ionization Potentials of Fimt formation constants of other complexes, as measured Transition Series Elements and Heats of Hydration and in water (1, $8-$9) (Table 3), indicates that the sta- Formation Constants of Complexes of Divalent Ions Sum of f i d and bility of these species increases as their covalent charsecond ion Hhydrotian Log Kpv. acter increases. This conclusion is based on the parallel ~otenlials (keel./ Log K., M(salzcy1change of electronegativity with ionization potential Ion (keal./mole) mole) M(NHA++ aldehude), among the metallic elements as discussed above. For example, it is seen from Table 3 that the sum of the first and second ionization potentials for nickel(I1) ion is much greater than that for manganese(I1) ion. This means that the electronegativity of nickel(I1) is greater than tbat of manganese(I1). Since the oxygen atom in water has a constant electronegativity which is greater than tbat of either of these metals, the elec-
382
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TABLE 4 Favored Electronic Configurations Gaseous ion (1) Filled d shell (2) Half-filled d shell (3) Empty d shell Complex ion involving covalent bonds (1) d'spa hybrid with filled underlyingdorbitals (2) dspPhybrid with filled underlying darbitals (3) daspa hybrid with half-filled underlying d orbitals (4) dsp' hybrid with half-filled underlying d orbitals (1)
with the electronegativity of the central metal ion. This, of course, is consistent with the trend observed among the hydrates and ammines of the divalent ions of the transition metals. It is, however, even more apparent in those cases where still greater increases in the electronegativity of the central metal ion are encountered. The oxidation of cobalt t o the tripositive state produces an ion which has a much greater attraction for electrons than does cobalt(I1) and which, in consequence, should form still more stable covalent bonds with such dipolar molecules as ammonia (the electronegativity difference between the donor and acceptor itom is less) (31). The same argument would be equally valid for the platinum metals, as compared to the corresponding triad of first transition series elements in any given oxidation state. THE EFTECT OF ELECTRONIC CONFIGURATION
The two parameters most often utilized in predicting covalent character are electronegativity and angular strengths of hybridized orbitals (Table 1). The angular strength is a property of a particular hybridized orbital and can provide no information with regard to the relative stabilities of two species having the same hybrid type. Similarly, although the electronegativity has a specific value for each chemical form of every atom ( $ I ) , transition element ions of the same charge
type probably do not differ greatly enough in their electronegativities to account for many of the great differences in the stabilities and labilities of their compounds, nor to account for the formation of particular hybridized orbitals. Unlike the elements of the representative families of the periodic system (i. e., alkali, alkaline earth, zinc, and aluminum families), the atoms and ions of the transition elements often possess different electronic confi~urationsof maximum stability in their covalentlv bonded forms from those of the corresponding gaseois atoms or ions. This may introduce a formidableenergy barrier into the complexing reaction and give rise to difficulty in relating all observed behavior to current theories. (The full significance of this factor is not always recognized as is evidenced by such statements as "double-bonding is probably the main factor which upsets predicted orders of stability of complexes based on electronegativities alone" (I).) In order to consider this aspect of the problem in a straightforward manner, it is convenient to make use
of a qimnle . - . ...r .. thermo~hemirnl .... ....- -. .- ....- -. rvrle 2.-
(gmund state) MC"(g)
+ nL(
