Werner Centennial - ACS Publications - American Chemical Society

40 Solvational Control in Spin-State Variations among Nickel(II) Complexes DARYLE

H.

BUSCH

Downloaded by UNIV OF PITTSBURGH on May 4, 2015 | http://pubs.acs.org Publication Date: January 1, 1967 | doi: 10.1021/ba-1967-0062.ch040

Chemistry Department, The Ohio State University, Columbus, Ohio

The various modes of interaction with molecules of sol vents that govern the spin states of nickel(II) ions are both subtle and potent. The best known examples in volve direct metal ion-solvent coordination. In certain cases, these processes represent poised equilibria with comparable amounts of the 6-coordinate triplet and lower coordinate singlet present. Such systems are particularly sensitive to minor variations in the character of the solvent. With bulky organic ligands, the complexes themselves probably act as solvent-structuring solutes. The effects of other solutes and other variables on solvent structure may be reflected in shifts in the equilibrium. The stoichiometric interaction of complex and solvent molecule via a hydrogen bonding process is also well documented and subject to similar variations. Apparently, in some nonprotonic solvents the dielectric properties support ionization equilibria that are associated with a spin-state change but, from the direction of the change, coordination of the solvent is not important.

Early references (2, 16, 17) to the variation of spin state for nickel(II) with the coordination environment consider two most plausible situa tions—the conversion of square-planar diamagnetic nickel (II) to octahedral paramagnetic nickel (II) by the addition of the fifth and sixth ligands above and below the plane of the first four ligating atoms, and the possi bility of the existence of either a ground state triplet or a singlet for nickel (II) in square-planar, 4-coordinate structures or tetragonal, 6-coor dinate structures. Indeed, the recurrent hypothesis has been offered that in certain poised systems the singlet and triplet states might have very 616 In Werner Centennial; Kauffman, G.; Advances in Chemistry; American Chemical Society: Washington, DC, 1967.


40.

BUSCH

617

Solvational Control

nearly the same energies and, therefore, might coexist. Extensive study has been dedicated to those systems that were suspected of illustrating this phenomenon (18). That the role of solvent in affecting spin-state variations deserves closer attention may be shown by summarizing a number of compelling phenomena. Though the nature of these effects has many aspects and might be variously described, three particular facets are explored here. In the case of certain nickel (II) systems poised equilibria occur with com parable amounts of 6-coordinate triplet and lower coordinate singlet present. Such systems are particularly sensitive to minor variations in the character of the solvent. W i t h bulky organic ligands, the complexes themselves probably act as solvent-structuring solutes, thereby lowering the activity of solvent molecules and making them less available for coor dination. The effects of other solutes and other variables on solvent struc ture may be reflected in shifts in the singlet-triplet equilibrium. The second set of phenomena involves stoichiometric interaction between nickel complex and solvent molecules without direct coordination by the solvent. In a number of instances it appears that hydrogen bonding in these definite solvates determines the spin state of the nickel(II) atom. Finally, in some nonprotonic solvents, the dielectric properties support ionization equilibria that are associated with a spin-state change but, from the direction of the change, coordination of solvent is not important. Solute Character and Solvent Activity It is instructive to begin with a consideration of systems that are not necessarily poised but may be shifted far toward a single, spin state. The macrocyclic ligand reported by Curtis (5), 2,4,4,9,11,11-hexamethyl1,5,8,12-tetraazacyclotetradecane ( N i complex, Structure I, hereafter 2 +

H / CH —C

CH

3

CH

CH 2

\

3

C—CH

NH

NH

NH

NH

C

C CH

3

3

3

I

In Werner Centennial; Kauffman, G.; Advances in Chemistry; American Chemical Society: Washington, DC, 1967.

618

WERNER CENTENNIAL

abbreviated as C T ) , is constrained by stereochemistry to coordination in a single plane. This constraint is most pronounced with the meso isomer, and discussion is confined to derivatives of that form. A series of salts and tetragonal complexes N i ( C T ) X has been reported (6). The anhydrous chloride and bromide, for example, are violet in color and exhibit the normal magnetic behavior expected of 6-coordinate nickel(II) complexes having triplet ground states (6, 11). When dissolved in nonpolar solvents, such as C H C 1 , the complexes remain 6-coordinate and paramagnetic (11). In contrast, when N i ( C T ) C l or N i ( C T ) B r is dissolved in water, it converts to a planar, diamagnetic complex which acts as a di-univalent electrolyte. Also, the nickel(II) ion in Ni(CT) + will not combine with N H when dissolved in aqueous solution (11). These observations present a number of disturbing features. The failure of H 0 and N H to coor dinate is inferred from the fact that the solution properties are typical of planar nickel(II). Both N H and H 0 occupy positions much higher in the spectrochemical series than Cl~~ or B r ~ . Consequently, if either of these neutral molecules were coordinated, the nickel atom would surely exist in its paramagnetic triplet state. The nagging question remains—why do these excellent ligands fail to coordinate? The most convenient molecular explanation would be of a steric nature; however, H 0 and N H molecules are substantially smaller than B r ~ so that spatial restrictions are not involved. It is tempting to seek some model based on specificity of the axial sites on nickel, perhaps for donors prefering B-type metals; however, quantitative studies (12) show that the spectrochemical series for the X groups in N i ( C T ) X is normal ( C N " ( > C T ) > N C S " > O N O > F " > N " > N O r ~ C l ~ > B r ~ > I~). It therefore becomes necessary to seek the cause in factors external to the complexes in question. That such factors are readily recognized will be evident shortly. A second example of failure of certain solvents to exhibit their expected coordinating abilities is found in the solution behavior of dibromo(*S,S -oxylyl-2,3-pentanedione-bis-mercaptoethylimine)nickel(II) (3, 4) (Struc ture I I , hereafter abbreviated as N i ( P E X ) B r ) . From the work of Imhof and Drago (9) we know that the ligand, field-splitting parameters for Ni(CH OH) in methanol and N i ( d m f ) in dimethylformamide are virtually identical. However, N i ( P E X ) B r shows sharply dissimilar be haviors in the two solvents (3, 4)- In dmf it dissolves to form a paramag netic, 6-coordinate, di-univalent electrolyte which is formulated as [ N i ( P E X ) ( d m f ) ] , 2 B r . In methanol solution, at equilibrium, it is a diamagnetic, di-univalent electrolyte. Also, it should be pointed out that N i ( P E X ) B r exists in simple paramagnetic, 6-coordinate form in nonpolar solvents, such as C H C 1 . As in the case of N i ( C T ) X , it must be inferred that C H O H does not coordinate with the nickel atom in N i ( P E X ) , although there is no obvious steric or electronic reason for this fact. 2

3

2

2


2

3

2

3

2

3

2

3

2

3

;

2

y

3

6

2 +

6

2+

2

2

2+

_

2

3

2

3


2 +

40.

BUSCH

619

Solvational Control

CH —CH 2

CH

C

CH,

S—CH

N

3

/

C

2

2

Ni \

. N \

Br

S—CH / CH —CH 2

2

2

2


II

Systems involving poised equilibria provide equally pertinent data. A n early and particularly revealing set of observations arises from J0rgensen's study (10) of the visible spectra of aqueous solutions of Ni(trien) +. In dilute solution in pure water, equal molar concentrations of nickel(II) salts and triethylenetetraamine exhibit properties typical of pseudo-octa hedral structure. In fact, the correspondence of the spectra to the octahedral prescription is so close that J0rgensen suggested the hydrated ion to exist in the cis form—i.e., cis-Ni(trien)(H 0) " . Now, as increas ing amounts of certain salts are dissolved in this solution, a new, intense absorption band appears in the visible region of the spectrum and increases as a function of the salt concentration. Perchlorates are particularly effective. The new absorption band corresponds to the appearance of planar nickel(II) and is accompanied by loss of octahedral nickel. The dissolving of, for example, N a C 1 0 shifts the equilibrium of Equation 1 to the right. 2

2

2

2

f

4

Ni(trien)(H 0) 2

(triplet)

2

;=±

Ni(trien) + 2

+

2H 0

0)

2

(aq)

(singlet)

Obviously, the added salt decreases the activity of the solvent in order to affect this poised equilibrium. K a r n (11) has recently found that the magnetic moment of aqueous (2,ll-dimethyl-3,7,ll,17-tetraazabicyclo(11.3.1)heptadeca-l(17), 13, 15triene)nickel(II) (Structure I I I , hereafter abbreviated as N i ( K N ) ) is indicative of a mixture of spin states. From the study of the temperature dependence of the magnetic susceptibility of the solution, the equilibrium of Equation 2 is deduced. 2 +

For the reaction as written, AH = +4.5 kcal./mole, and AS = 16 eu. These values are concentration dependent because of the effect of the com plex on the activity of water. Two facts are particularly notable. The large, positive entropy change associated with the formation of the diamag netic species is consistent with the desolvation model. Also, this is a rare


620

WERNER

H C

CENTENNIAL

CH,

3

v / \ / \

C

N

,

C H

H

H—N—Ni —N—H 2 +

CH

I

2

CH —CH Downloaded by UNIV OF PITTSBURGH on May 4, 2015 | http://pubs.acs.org Publication Date: January 1, 1967 | doi: 10.1021/ba-1967-0062.ch040

2

2

N

CH

H

CH —CH

2

I 2

2

III Ni(KN)(H 0) 2

2

2+

(aq)

—

(triplet)

Ni(KN) 2+

+

(aq)

2H 0

(2)

2

(singlet)

i nstance in which the singlet is at higher energy than the triplet in a system involving nickel(II) in a singlet-triplet equilibrium. The complex behavior of certain of the Lifschitz salts probably arises from similar causes. The work of Higginson, Nyburg, Wood (8), and N y b u r g and Wood (19) is sufficiently detailed to lead to interesting speculations. In acetone solution, N i ( m - s t e i n ) 2 ( C l C H C 0 ) 2 , where E a stern represents meso-stilbenediamine, exists as a blue, presumably para magnetic, material; however, portionwise addition of water leads to i n cremental formation of a diamagnetic yellow form. It would be most comfortable to conclude that the enhanced dielectric constant of water merely leads to a progressive shift of the equilibrium of Equation 3 to the right. The trouble is the blue crystals isolated from acetone contain the ion (x-ray crystal structure) N i ( m - s t e i n ) ( H 0 ) ' . This suggests the 2

2

Ni(m-stem) (Cl CHC0 ) (triplet) 2

2

2

2

2

^± Ni(m-stein) (singlet)

2

2

2+

2

2

f

+

2C1 CHC0 2

2

(3)

alternative and fascinating process given in Equation 4. Note that this requires the coordination of water to be most extensive when water is in low abundance. +H 0 2

Ni(m-stein) (H 0) 2

2

2

2+

^

-facetone

Ni(m-stein)

2

2 +

+

2H 0 2

(4)

The most readily apparent factor common to these several processes is their dependence on the availability of water (or other solvent) molecules


40.

BUSCH

621

Solvational Control

for coordination. A second, equally important factor is the large bulk of organic matter in the ligands. In fact, these complexes may be thought of as irregular doughnut shaped masses of hydrophobic material with positively charged Lewis acids occupying the "hole." A s a consequence of this structural character, a water molecule must be largely separated from the bulk of the solvent in order to coordinate. Thus, the fulcrum of the poised equilibrium has as its supporting bases the processes of Equations 5 and 6. H 0( ) Downloaded by UNIV OF PITTSBURGH on May 4, 2015 | http://pubs.acs.org Publication Date: January 1, 1967 | doi: 10.1021/ba-1967-0062.ch040

2

2H 0 2

(vap

)

aq

+

^ NiL

(5)

H20(vap) 2 +

-

NiL(H 0) 2

2

2+

(6)

The general hydrophobic character of the ligands has a major effect on the first equilibrium—i.e., on the availability of the water, and it is this factor that requires attention. It is well-recognized that hydrophobic solutes, such as hydrocarbons, dissolve with sizeable negative AH and AS of solutions. This and other data have led to theories explaining how such solutes cause the degree of structuring, or ice-like character, of water to increase (14). In the flickering cluster model of water, an adiabatic condition is marked by rapid, random fluctuation of partially structured regions throughout the volume, with the total fraction of the fluid that is struc tured a constant. The nature of the structuring is presumed to involve hydrogen bonding, as in ice, and a given water molecule at any instant might be characterized in terms of the number of hydrogen bonds it has formed. Nonpolar, structure-making solutes increase the degree of struc turing of the water by stabilizing "Frank-Evans icebergs" adjacent to themselves. The hydrophobic character of these solute particles may be considered to cause an interface of structured water to be formed that is not unlike a surface. The structuring presumably is transmitted through a number of layers of solvent molecules. In contrast to this behavior, ions of modest electrostatic fields reduce the degree of structuring of the solvent water. The solvent molecules in such cases can be sorted into three regions—the inner sphere, where they are strongly oriented; a structure-broken region; and the bulk of the sol vent where the molecules are normal. Apparently, strongly hydrated cations may behave quite differently. There is reason to believe that these species may not only alter but also increase the extent of structuring of water. Thus, an ion having a coordination sphere belonging to the cubic system might induce structure on the solvent through perpetuation of its own ligand array into the bulk of the solution by means of strong hydrogen bonds.


622

WERNER

CENTENNIAL

Returning to the case at hand, these complexes probably act as hydro phobic solutes and influence the structure of water so that the solvent molecules are less available than they otherwise would be. The presence of a positive charge is probably not effective in affecting solvent structure because the charge is buried in the large hydrophobic structure. One can envision a Frank-Evans iceberg enclosing one of the complex ions with the ice-like structure bridging over the "hole" containing the metal ion. Therefore, in the first example cited, N i ( C T ) + decreases the activity of water because it is a structure-making solute, thereby making the solvent unavailable for coordination. The example of N i ( K N ) + illustrates an essential aspect of this theory. Because it is the hydrophobic character of the complex itself that is causing the nickel atom to have to relinquish its coordinated water molecules, the relative concentration of singlet at equilibrium should be enhanced as the total concentration of complex is i n creased. This is indeed the case. Further, a critic must point out the fact that the concentration of water decreases as the concentration of complex increases. This cannot be the cause of the observed effect, for the range of concentrations is small and all solutions are dilute so that the concentration of water is essentially constant throughout the range. 2


2

Similarly, the role of acetone in the case of Ni(m-stein)2 " " might be to disrupt the structure of water, thereby counteracting the effect of the com plex which, again, is structure-making. The distinction between the be havior of C H 3 O H and dmf toward Ni + and N i ( P E X ) + suggests that C H 3 O H has a hydrogen-bonded partial structure analogous to that of water. In such a case, the cubic Ni + ion might readily be incorporated into the solvent structure, while coordination of C H O H to N i ( P E X ) requires removal of that solvent molecule from interaction with the rest of the solvent. 2

2

1

2

2

3

Specific Hydrogen Bonding

2 +

Effects

Despite the considerable coordinating ability of the water molecule (or CH3OH), other intermolecular associations sometimes take precedence under conditions where coordination to a metal ion would otherwise be expected. This is well-illustrated by the behavior of systems containing water in solvent quantities, as described in the preceding section. There also exist well-documented examples of related processes in which only stoichiometric amounts of water or related material are involved. Perhaps the most clearly defined example involves the macrocyclic complex of Curtis, N i ( C T ) X . The solid, violet, paramagnetic, anhydrous chloride and bromide must be carefully protected from moisture for they readily revert to yellow, diamagnetic dihydrates. The fact that two mole cules of water are taken up to produce a diamagnetic nickel (II) ion is especially interesting for displacement of the halide, and coordination by the water molecules would surely produce a violet, paramagnetic species 2


40.

BUSCH

623

Solvational Control

i n view of the relative positions of H 0 , B r ~ , and C I " in the spectrochemical series. The most reasonable supposition is that the water molecule inter acts with the complex in such a way that the ligand field of the axiallyoriented halides is diminished. Clearly, this may be accomplished by hydrogen bonding of the water molecule to the coordinated halide. Curtis first offered this explanation (6). Such a model had been offered earlier to explain the variation in properties with hydration of the iV,iV-dialkylethylenediamine complexes of nickel halides (7). The macrocyclic complex N i ( C T ) B r exhibits a similar effect i n solu tion in nonpolar solvents (11). If dissolved i n freshly purified chloroform, it has the expected molecular weight (osmometer) for a monomeric, 6coordinate structure and retains the color and spectrum of a typical pseudooctahedral nickel (II) complex having a triplet ground state. However, if a hydroxylic solute is added or, in fact, if the stabilizer ( C H O H ) is not removed from the chloroform, the nickel atom exhibits the spectrum of square-planar, diamagnetic nickel(II), but still shows a normal molecular weight. It is suggested that hydrogen bonding, as shown in Structure I V , weakens the bond between the metal atom and the halide ion, thereby producing the singlet state. 2


2

2

5

IV Recently, a closely related model has been invoked (13) to explain the role of a single water molecule in promoting 5-coordination. The diamagnetic, complex ion is bromo(2, 12-dimethyl-3, 7-11, 17-tetraazabicyclo(11.3.1)heptadeca-l, (17), 2, 11, 13, 15-pentaene)nickel(II) (hereafter abbreviated as Ni(CU)Br+). This ion forms only monohydrated salts, and the water of hydration is not removed at 100°C. in vacuo. The proposed function of the water molecule is shown i n Structure V . A similar role has been suggested for water molecules in systems ex hibiting more complex, magnetic behavior. In the case of the nickel (II)



624

WERNER

CENTENNIAL

V complexes of the tetradentate ligand tetrabenzo[b, f, j , n][l, 5, 9, 13]tetraazacyclohexadecine (Structure V I , hereafter abbreviated as T A A B ) , N i ( T A A B ) X , three classes of behavior were observed (18). For X ~ = C10 ~, B F ~ , or B D q ( B r - ) > D q ( I " ) . A s in other cases, the water cannot profitably be assumed to be coordinated for it is a stronger ligand than the halides, and it should surely produce the triplet ground state. The system can be satisfactorily explained as a result of detailed magnetic (18) and spectral (11, 12) studies. The temperature variation of the magnetic


40.

BUSCH

Solvational Control

625

susceptibility agrees nicely with a model based on an equilibrium between singlet and triplet states, and the thermal parameters are AH = +800 cal./mole, A S = +0.57 eu for N i ( T A A B ) C l H 0 and AH = +700 c a l . / mole, AS = - 0 . 4 5 eu for N i ( T A A B ) B r • H 0 . The strict, linear graphs of InK vs. 1/T reveal that AH is independent of temperature over the brief 200°C. range studied, thus sustaining the two-state model. The values of AS differ from that expected for a change in spin state alone. This demonstrates the superposition of at least one additional, structural change involving a negative entropy change. The first spectral band has been assigned to the transition Bi —> E , and the energy of this transition de pends in part on the axial ligands (1). Consequently, its position should reflect the position of the anion in the spectrochemical series. The ob served frequencies are C l ~ , 8800 c m . ; Br~, 8300 c m . ; and I~, 8100 c m . , as expected on the basis of the usual behavior of nickel(II). It is there fore concluded that the portions of the samples present as triplets are nor mal. It follows then that the occurrence of the singlet-triplet equilibrium probably should not be attributed to the destabilizing of the triplet but to stabilizing of the singlet. Such a situation can be envisioned i n terms of the equilibrium of Equation 7. The energy of hydrogen bonding is pre sumed to stabilize the singlet isomer. In view of the poor hydrogen bond ing ability of the iodide ion, the distinction between the iodide, on the one hand, and the chloride and bromide derivatives, on the other, is readily understood. 2

2


z

g

z

2

2

g

- 1

- 1

[Ni(TAAB)X ] + H 0 ^ [Ni(TAAB)X] + X - - . . H O H (hydrogen bonded (octahedral, (interstitial (5 coord., triplet) water) singlet) anion) 2

+

2

or

- 1

[Ni(TAAB)] (planar, singlet)

2+

(7)

+ X-.-.HOH.-.X(hydrogen bonded anion)

A more subtle example of the role of hydration is found in the proper ties of solid N i ( P E X ) B r (3, 4). For the related compounds, N i ( P E X ) Y , where Y ~ = C10 ~ or I~, the nickel atom is in its singlet state, while the triplet state occurs for Y ~ = C l ~ , N ~ , and N C S ~ . The bromide complex has a magnetic moment of 1.57 B M , and the magnetic susceptibility obeys the Curie-Weiss law with only a small Weiss constant. This result sug gests the possibility of nonequivalent nickel(II) atoms in the unit cell, some being diamagnetic and some paramagnetic. Precedents are found in the structure of bis(diphenylbenzyl phosphine)nickel(II) bromide (15), which contains two paramagnetic tetrahedral molecules and one isomeric diamagnetic planar molecule per unit cell and in the structure of a yellow form of bis(m-stilbenediamine)nickel(II) dichloroacetate which contains two paramagnetic octahedral nickel(II) complexes and one planar dia2

2

4

3


626

WERNER CENTENNIAL

magnetic complex per unit cell. paramagnetic ion is normal.

In such cases, the moment calculated per

If N i ( P E X ) B r contains both paramagnetic, pseudo-octahedral, and diamagnetic species in the unit cell, the ratio must be three diamagnetic for each paramagnetic ion because the squares of the moments average. This assumption leads to the assignment of /*eff = 3.14 B M for the paramagnetic species present. 2

The relationships just described apply to samples carefully protected from moisture. N i ( P E X ) B r slowly absorbs water, apparently approach ing a limit of 3^ mole of H 2 0 per mole of complex. This change in com position is accompanied by a parallel decrease in /i ff, approaching a small limiting residual value typical of spin-paired nickel (II). It is concluded that the one molecule i n four of N i ( P E X ) B r that was in the paramagnetic, triplet state has combined with two molecules of water and changed into a planar, diamagnetic form. Again, the hydrogen bonding role of water is apparent. 2


e

2

Spin-state Variations in Nonhydroxylic Solvents In CH3NO2, N i ( P E X ) C l behaves essentially as a nonelectrolyte, re taining coordination of the two chlorides and displaying triplet, magnetic and spectral properties, closely related to those observed for the solid (8, 4)- In contrast, N i ( P E X ) I is diamagnetic both i n the solid state and in CH3NO2 solution, having very similar spectra i n the two states. M o l a r conductance in nitromethane shows the compound to be a uni-univalent electrolyte (XJW = 75 ohm" at C = 10r M), thereby strongly suggesting the cation to be 5-coordinate, N i ( P E X ) I + . 2

2

z

-1

N i ( P E X ) B r is intermediate in properties and behavior. In nonpolar solvents, such as dichloroethane, the compound exists as a neutral, molec ular, 6-coordinate species having a normal, triplet, ground state, as inferred from molecular weight, magnetic moment, and electronic spectra. H o w ever, in nitromethane, at room temperature, the substance exhibits inter mediate values for molar conductance (58 ohm"" at 10~" M) and magnetic moment (2.65 B M at 9.44 x 10~ M). Dilution experiments yield values for the molar conductance that are consistent with an equilibrium constant of 1.66 x 10" (T = 25°C.) for the process given i n Equation 8. 2

1

3

3

3

Ni(PEX)B (triplet)r 2

Ni(PEX)Br+ + B r " (singlet)

(8)

N i ( P E X ) B r is assumed to be a diamagnetic, 5-coordinate complex in analogy to N i ( P E X ) I . The validity of the treatment is verified by the predicted moment of the solution showing 2.65 B M ; theory predicts 2.60 B M , which is in excellent agreement. This appears to be a satisfactory +

+



40.

BUSCH

Solvational Control

627

example of the variation of spin state in response to the dielectric properties of solvent, for there is no reason to assume coordination by the nitromethane. This example contrasts to the cases cited for water solutions, because the strong coordinating ability of water requires that solvent molecules complex or that there be a complicating reason this does not occur. It has always been the temptation of the coordination chemist to presume that the dissociation of a bound group in solution leaves a coordination site vacated, even in solvents that can readily provide ligands. Present orientation of the field would lead the investigator to presume exactly counter to that old difficulty and always assume that the ionization process involves displacement by solvent molecules rather than dissociation under the action of the dielectric property of the solvent. Clearly neither assumption is generally valid, and a substantial number of appropriate examples must be subjected to careful experimentation. Acknowledgment This investigation was supported by the United States Public Health Service Grant GM-10040 from the National Institute of General Medical Sciences. This financial assistance is sincerely appreciated. Literature

Cited

(1) Ballhausen, C. J . , "Ligand Field Theory," McGraw-Hill Book Co., Inc., New York 1960. (2) Ballhausen, C. J . , Liehr, A. D., J. Am. Chem. Soc. 81, 538 (1959). (3) Brubaker, G. R., thesis, The Ohio State University, 1965. (4) Brubaker, G. R., Busch, D . H . , 148th Meeting of the American Chemical Society, Detroit, 1965. (5) Curtis, N. F., J. Chem. Soc. 1964, 2644. (6) Curtis, N. F., private communication, 1965. (7) Goodgame, D. M. L., Venanzi, L .M.,J. Chem. Soc. 1963, 616. (8) Higginson, W. C. E., Nyburg, S. C., Wood, J . S., Inorg. Chem. 3, 463 (1964). (9) Imhof, V . , Drago, R. S., Inorg. Chem. 4, 427 (1965). (10) Jørgenson, C. K . , Acta Chem. Scand. 11, 399 (1957).

(11) Karn, J. L., thesis, The Ohio State University, 1966. (12) Karn, J . L., Busch, D. H., Abstracts of Papers, 151st Meeting of the American Chemical Society, Pittsburg, H94, 1966. (13) Karn, J . L., Busch, D. H., Nature 211, 160 (1966). (14) Kavanaugh, J . L., "Water and Solute-Water Interactions," Holden-Day, Inc., San Francisco, 1964. (15) Kilbourn, B. T., Powell, H . M., Proc. Chem. Soc. 1963, 207. (16) Maki, G., J. Chem. Phys. 29, 1129 (1958). (17) Matoush, W. R., Basolo, F., J. Am. Chem. Soc. 75, 5663 (1953). (18) Melson, G. A., Busch, D . H . , J. Am. Chem. Soc. 86, 5830 (1964). (19) Nyburg, S. C., Wood, J. S., Inorg. Chem. 3, 468 (1964). RECEIVED July 1, 1966.


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