I. ADATO,A. H. I. BEN-BASSAT, AND S.SAREL
3828 for his helpful discussions and constant encouragement throughout this work. He also wishes to express his thanks to Professor Masaji Miura of Hiroshima
University for his kind advice and helpful suggestions and to Dr. Peggy Knecht for her assistance in the grammatical improvement of the article.
Metal-Ligand Bonding in Copper(I1) Chelates-an Electron Paramagnetic Resonance Study by I. Adato, A. H. I. Ben-Bassat, Department of Inorganic and Analytical Chemistry, Hebrew University of Jerusalem, Israel
and S. Sarel* Department of Pharmaceutical Chemistry, Hebrew University of Jerusalem, Israel (Received December 29, 1970) Publication costs borne completely by The Journal of Physical Chemistry
E p r spectra of some cupric p-ketoenolates, a-dioximates, p-dioximates, and Schiff bases in chloroform solution were recorded a t room temperature and a t - 160’ (frozen state). The bonding parameters were calculated from spectral data. The bonding properties of the odd electron were derived as a function of the nature of the ligand donor atoms, the size of the chelate ring, and substituent effects. The relationship between covalency of .rr bonding and complex stability is discussed.
This work was aimed at studying the electron spin resonance spectra of several copper(I1) complexes and to draw some conclusions concerning the bonding properties of the odd electron, depending on the nature of the donor atoms in the ligands, the size of the chelate ring, and effects of substituents on the chelate ring. The epr spectra of copper(I1)-acetylacetonate type complexes have been treated by several authors.’+ In all of the above complexes, it was assumed that the CuII ion is surrounded by the four donor atoms in a planar configuration (D2h symmetry) and that each donor atom has a 2s,2px,2p,,2p, atomic orbital available for the formation of molecular orbitals with the atomic d orbitals of the central ion. From group theory, Gersmann and Swalens obtained the proper linear combinations of ligand orbitals with the copper d orbitals which form the antibonding wave functions of the planar cupric chelate compound. The expressions for the in-plane antibonding wave functions are
Big = adxy -
a!’
- [-
2
+
gXY(l)
+
u ~ ~ ( ~~ x) ~ -( u~ ~ )Y
( ~ ) ]
+
A , = pdx,-,s - ‘/2(1 - pz)l”[-pxY(l)- pxY(2) pxY(3) pxy(4)l
+
The above wave functions are applicable also for The Journal of Physical Chemistry,
T701.
76, No. 26, 1971
tetradentate Schiff base chelates of Czv symmetry, wherc the corresponding symmetry designations are Bz and AI, in the sense used by Swett and Dudek.6 The bonding parameters, a!, p, and 6, are a measure of the covalency of the appropriate bonding. A value of 1 for the square of the parameter indicates a completely ionic character, while a value of 0.5 denotes essentially a purely covalent character. The odd electron is placed in the antibonding B1, orbital in the ground state. Overlap is included for the B,, orbital, where a! and a!’ are related a2
+
(Y’Z
- 2aa’S
= 1
For a ligand-to-metal distance of R = 1.9 8, the overlap integral values (8)have been assigned as 0.076 in &keto enolates, 0.093 in a!-dioximates and p-dioximates, and the value 0.084 in Schiff base c ~ m p l e x e s . ~ ~ ~ (1) A. H. Maki and B. R. McGarvey, J . Chem. Phys., 29, 31, 35 (1958). (2). D.Kivelson and R. Neiman, ibid., 35, 149 (1961). (3) H. R. Gersmann and J. D. Swalen, ibid., 36, 3221 (1962).
(4) H.A. Kuska, M . T. Rogers, and R. E. Drullinger, J . Phys. Chem., 71, 109 (1967). (5) V.S. Swett and E. P. Dudek, ibid., 72, 1244 (1968). (6) S. Antosik, M. M. Brown, A. A. McConnell, and A. L. Porte, J . Chem. SOC. A , 545 (1969).
METAL-LIGAND BONDING IN Cu(I1) CHELATES
3829
n
By applying the wave function for the BI, state to the Hamiltonian of Abragam and Pryce,’ the magnetic parameters are shown to bea 911 = 2.0023
- (8X/AE,n-,n)
[a2P2-
f(P)]
~ ~g(S)] ~ 9, = 2.0023 - ( ~ X / A E X , ) [ C YAll = P[-a2(4/7 K ) - 2Xa2(4P2/AE,,
+
+
362/7AE,,)
A, where
= P[a2(2/7
K)
-
- 22Xa262/14AE,,)
+ ffa’p(1 - /3”1’”(n)/2 = aff’62S + ffff’6(1 - 62)1’zT(n)/2
f(P) = y(6)
-
so c
aa’p2S
with the spin-orbit coupling constant for the Cu(I1) ion X = -828 cm-’, the free-ion dipole term P = 0.036 cm-l, and the Fermi contact term K = 0.43 as defined The corresponding values of T ( n ) are:2,5 T(n).itropen= 0.333, T ( n ) o x y g e n = 0.22, and T ( n ) = 0.276 for the Schiff base complexes. Values of go and 911 were obtained from the spectra recorded a t room temperature and a t - 160°, respectively. The g1 values werecalculated from the equation 90 = 2 9 J 3 19Il/3. The values of a2 were calculated from the copper hyperfine spectra, using thc approximate formula2
+
+
a2 = - ( A ~ l / p )
(si1
- 2)
+ 3(g,
-
2)/7
+ 0.04
and the values of P were calculated from the expression of 911 and the values of 6 were calculated from the expression of g,. Values of a’ were also calculated from the nitrogen hyperfine splitting (WL) exhibited by three complexes and then a2 values were recalculated from the normalization condition of the B1, state wave function and compared to those obtained from the approximate formula of Kivelson and Neiman.2 Adequate interpretation of the bonding parameters requires separation between the ground state term (u*,dxy) and the excited statc terms, (r*,d,I-yz) and (T*,dx.). The relative positions of the molecular d orbitals in planar bis(ketoenolato)copper(II) complexes were determined by Cotton and Wise8 from polarization studies of single crystal spectra of CU(DPRS)~. It was concluded that the four d-d transitions occur within a range of a few thousand wave number and that d,,, d,,, d,,.+, and d,, orbitals lie close together, some 20.000 cm--’ below thc d,, orbital. The possible d-d transitions within the planar cupric complex were arranged in the following order of increasing energy: d,, d,,, d,,, d,z-,,, dyz. In our estimations of the ligand field encrgies we counted on the above order of the d-d transitions. Due to the small difference between the coordinate systems of Gersmann3 and that of Cotton,*the order in ligand field transitions is AEyz < AE,,-,z < AE,,.
Figure 1. Epr spectrum of chloroformic solution of Cu(DMGL)t at room temperature.
Results The epr spectra recorded for P-ketoenolate complexes of copper(I1) are charactcristic of those of planar copper(I1) compounds in their anisotropies in g and hyperfine coupling constants. The epr spectra of bis(dimethylg1yoximato)copper(I1) “Cu(Dk1GL)” (Figure 1) and bis(dipheny1glyoximato)copper(II) “Cu(DPGL)2” in CHCla solution recorded a t room temperature consist of four almost equally spaced hyperfine lines characteristic to the Cu(I1) nuclear hyperfine interaction. Onto each copper hyperfine line are superimposed the lines which are derived from the extrahyperfine interaction with the nitrogen nuclei of the donor ligand atoms. If we account for the distortions that might occur in the nitrogen spectrum and consider the relative intensities of the signals (the amplitude ratios) in the epr spectra of Cu(DRIGL)2 and CU(DPGL)~,we should find that four nitrogen atoms are bonded to the central Cu(I1) ion. Tho spectrum of Cu(DRlGL)2at - 160” is shown in Figure 2. The distinction between the components of the 811 peak is very difficult, due to the overlap with the g, peak. In low-temperature spectra the nitrogen hyperfine structurc appears usually on the g, component of the spectra. We used this component for the evaluation of g, while gi 1 was calculated from the approximate relationship go = 1/3(g11 2gJ. The g values obtained this way approximate those obtained by Wiersemae for CU(DMGL)~in alcoholic solution, and those by Schubel’O in pyridine solution. Due to uncertainty in evaluations of A I 1 and A I,we calculated the in-plane u-bonding parameter using the splittings of the well-resolved nitrogen hyperfine lines of the spectra. The excellent agreement between the c y 2
+
+
(7) A. Abragam and H. M. L. Pryce, Proc. R o y . Soc., Ser. A , 205, 135 (1951). (8) F.A. Cotton and J. J. Wise, Inorg. Chem., 6, 909,915,917 (1967). (9) A. K. Wiersema and J. J. Windle, J. P h y s . Chem., 68, 2316 (1964). (10) W. Schubel and E. Lutze, 2. Angew. Phys., 17, 332 (1964). T h e Journal of Physical Chemistry, Vola76,No. 86,1971
3830
I. ADATO,A. H. I. BEN-BASSAT, AND S: SAREL
- 1 50 0
Figure 2. Epr spectrum of chloroformic solution of Cu(DMGL)a a t -160”.
V‘
I-------I 50 G
Figure 4. Epr spectrum of chloroformic solution of Cu(BAD0)z (at - 160”).
4 Figure 5. Epr spectrum of chloroformic solution of N,iV’-bis(salicyla1dehyde)ethylencdiiminocopper(11) (at
- 160’).
large number of lines is observed especially at the lowtemperature spectrum (Figure 4). Similar splitting in Figure 3. Epr spectrum of chloroformic solution of the epr spectrum of a frozen CHCI, solution of bisCu(AADO)* (at - 160’). (dibenzoylmethanato)copper(II) was noted by Kuska4 and attributed t o the phenyl group protons. value obtained from the epr spectra of C U ( D I ~ G L ) ~ The epr spectra of bis(N-pentylsalicyla1diminato)measured in this study and the a2value calculated by copper(I1) do not exhibit super hyperfine splittings and Ross” is noteworthy. are characteristic of planar Cu(I1) compounds. The CHCls solution of bis(acety1acctone dioximato)The epr spectra of N,N’-bis(salicyla1dehyde)ethylcopper(I1) “Cu(AAD0)” at - 160” exhibits a charenediiminocopper(I1) at - 160” (see Figure 5 ) suggest acteristic epr spectrum of a planar Cu(I1) complex that the Cu(I1) is placed in two different environments compound, especially with nine well-resolved nitrogen in the complex molecule. hyperfine lines at the g, component (Figure 3). The The magnetic and bonding parameters calculated agreement between a2 values obtained by the two from epr and electronic absorption data are given in different calculation methods (Kivelson’s approximate Tables I and II.1z-16 formula and extrahyperfine structure splitting formula) is remarkable (Table 11). In the spectrum of bis(11) B. Ross, Acta. Chem. Scand., 21, 1855 (1967). (benzoylacetone dioximato)copper (11) “CU(BADO)~” (12) L. L. Funck and T. R. Ortolano, Inorg. Chem., 7, 567 (1968). the resolution of the nitrogen hyperfine lines at the g, (13) D. P. Graddon and E , C. Watton, J . Inorg. NucZ. Chem., 21, component is indeed very poor. On the other hand, a 49 (1961). T h e Journal of Ph‘ysical Chemistry, Vol. 7 6 , N o . 86, 1972
METAL-LIGAND BONDING IN Cu(I1) CHELATES
3831
Table I : Magnetic Parameters of Cu(I1) Chelates in Chloroform All X 104,om-1
go
Compound
91I
Q1
1. Bis(hexafluoroacety1acetonato)copper(II ) 2. Bis(trifluoroacety1acetonato)copper (11) 3. Bis(benzoy1acetonato)copper(I1) 4. Bis (acety1acetonato)copper(I1)
2.318 (2.306) 2.312 (2.308) 2 281 (2.281) 2.283 (2.285) (2.266) (2.264)
2.046 (2.051) 2.042 (2.040) 2.045 (2.046) 2.047 (2.042) (2,053) (2.036)
2.171
2.088
@-Diketonates 2.137 174 ... (173) 2.132 165 ... (167) 2.124 178 *.. (176) 2.126 173 ... (175) ... (160) ... (145.5) p-Dioximates 2.116 189
2,211
2,073
2.119
2.167 (2.15) 2.136
2.054 (2.05) 2.048
9. Bis(N-pentylsalicyla1diminato)- 2.225 copper (11) 2.209 10. N, N '-Bis (salicylaldehyde)ethylenediimino-copper(I1) (2.19)
2.067
5. Bis(acety1acetondioximato)copper(11) 6. Bis(benzoy1acetondioximato)copper(I1)
7. Bis(dimethylg1yoximato)copper (11) 8. Bis(diphenylg1yoximato)copper(I1)
a
I n CHCla.
Single crystal.
c
I
2.046
...
A i X 104, om-1
Aa X 104, cm-1
24.6 (24.5) 25.2 (25.2) 26.5 (26.6) 28.1 (28.2) (19 * 0) (29.0)
74.4
...
4a 40
*..
4a
...
lb
3c
20.0
76.3
184
21.6
76.5
(144)
(14.6)
87.0
...
gd
84.6
Schiff bases 2.120 170
Toluene (60%)-CHCla (40%).
...
77.0 .*. 76.4
...
2.077
2.100 (2.11)
45
71.8
a-Dioximates 2,092
...
Reference
30.5
198 (208)
77.0
30.8
86.5
...
*..
Aqueous ethanol.
e
58
I n 80% CHCla-20% toluene.
Table 11: Ligand Field Energies and Bonding Parameters of Cupric Chelates B2
1. 2. 3. 4.
5. 6. 7. 8. 9. 10. a
14,000
15,800
14 ,800 15,600 16 ,100 16,600
16 ,800 18,000 18,000
14,814 18,000 18,000
18,000
18,600 18,000
18,500 27,000
0.861 0.830 0.835 0.824 0.824 0.774 0.793
I
Referenaea
0.836
0.777
12
0.913 1.0
0.795 0.866 0.870
13 8 14
0.777 0.714 0.740
0.766 0 818
62
0.65 0.843
15
0.80
The sources of the ligand field energies.
Discussion Calvin and Wilson'o were the first to suggest the possibility that two force components might be responsible for containing the copper ion in the chelate: one is of the same nature for both copper and hydrogen, and the other, of quite different nature for copper than for hydrogen. The first component should manifest the ionic or coulombic effects, i e . , the effects of charge and charge distribution of the ligand anion, and of the charge and radius of the cation. The second component, assumed t o reflect the stabilization due to the enolate resonance system, plays a different and far
greater role in bonding the copper atom than in bonding hydrogen atom. The in-plane a-bonding coefficients a* gives the expected trend in covalency. From the epr spectrum of P-ketoenolates it can be seen that electron-withdrawing substituents in the acetylacetonate skeleton decrease the covalency of the u bonding. The change in the (14) C. Dijkgraaf, Theoret. C'him. Acta, 3 , 38 (1965). (15) 8. Yamada and R. Tsuchida, Bull. Chem. SOC.Jap., 2 7 , 156 (1954). (16) M. Calvin and K. W. Wilson, J . Amer. Chem. SOC.,76, 2003 (1954). The Journal of Physical Chemistry, Vola76, N o . 86,1071
I. ADATO,A. H. I. BEN-BASSAT, AND S. SAREL
3832 covalency may result from the first force component between the copper(I1) and the ligand donors, i.e., a change in the effective negative charge on the oxygen donor atom, caused by the inductive effects of the substituents. On the other hand, electron-withdrawing substituents on the @-ketoenolate skeleton cause an increase in the covalent character of the n bonding, especially of the out-of-plane n bonding. Van Uitert, et al.,” also pointed out that @-diketones of comparable pK, values bearing one methyl and one aromatic substituent do not appear to chelate as strongly as those bearing two aromatic substituents. I n view of the above epr data, Van Uitert’s observation can be interpreted in terms of more covalent out-ofplane n bonding between the metal and the chelated ligand. The aromatic substituents induce a more effective condition toward formation of the enolate (benzenoid) resonance of the chelate ring. In fact, the polariBability of unsaturated molecules containing n bonds is greater than that of saturated molecules containing u bonds only. On conjugation, the n electrons are delocalized and present in the extended n orbitals. From Table I1 it can be seen that the greater covalent character in the in-plane u bonding is exhibited by the a-dioximates. The copper(I1) ion in the latter is surrounded by four nitrogen atoms which are part of two five-membered chelate rings. The increase in ring size in ,f3-dioximate series results in a decrease in the covalent character of the u bonding. This is due to the greater stability of the five- relative to the sixmembered ring structure. The observed greater covalent character of the inplane u bonding in the @-dioximates,as compared to similar bonding in 0-ketoenolates, can be attributed to the difference in the electronic characteristics between the nitrogen and the oxygen atoms of the donor groups. The intermediate case, where two nitrogen and two oxygen donor atoms are similarly bonded to the central Cu(I1) ion, is reflected in the bonding parameters of the Schiff base complexes. In these cases, one has to account for two different effects which operate in opposite directions on the stability of the complex: (a) the number of the chelate rings and (b) the decrease in benzenoid resonance of the salicylaldehyde derivatives. Calvin and Wilson’o have noted that the copper(I1) chelates derived from @-diketones exhibit greater stability than the analogous chelates from salicylaldehyde of comparable ligand basicity. This was attributed to decrease in the resonance contribution from the active grouping in salicylaldehyde than in p-diketone. From application of polarographic techniques, Calvin and Bailesl* noticed that the increase in number of the rings in chelated Schiff bases is accompanied by an increase in the stability of the complex. It appears from epr data that the increase in the stability of the complex molecule is reflected by the deThe Journal of Physical Chemistry, Vol. 76, No. 26,1972
crease in the p2 value, i.e., implying an increase in the covalency of the n bonding. Swett and Dudek6 recorded the epr spectra of a series of bidentate and tetradentate Schiff bases and found that the covalency of the in-plane n bonding is greater in tetradentate Schiff bases, while the covalency of the u bonding is in the opposite order. Table I11 gives the polarographic and epr data for bis(N-methylsa1icylaldiminato)copper(I1) “A” and N,N’-bis (salicylaldehyde) ethylenediiminocopper(I1) “B.”
Table I11 e11/2,a
A
B a
Reference 18.
v
+0.02 -0.75
a2b
P
0.79
0.77
0.80
0.64
Reference 5 .
The above data clearly iiidicatc the importance of the n bonding on the stability of the complex molecules. The large number of lines appearing in the epr spectra of bis(benzoylacetonedioximato)copper(II) (see Figure 4) suggest that additional hyperfine interactions occur with the phenyl protons. Kuska4 considered three possible mechanisms for these splittings. (1) In dilute solution the chelates constitute dimers where the phenyl group of the first complex molecule lies above the x axis of the second complex molecule. (2) The electron spin density a-delocalized, as predicted by extended Huckel molecular orbital calculations of Wise.* (3) The spin density is n-delocalized through a configuration interaction mechanism similar to that proposed by K i v e l ~ o n ’and ~ Fortman.20 Kuska has found that proton splittings do not appear when one methyl group is substituted by a phenyl group in the @-ketoenolate skeleton [bis(benzoylacetonato)copper(II) 1, and this phenomenon is dificult to explain in terms of u delocalization. In contrast, bis(benzoylacetonedioximato)Cu(II) does exhibit proton splittings. We think that the higher covalency of the n bonding is the main cause for these interactions. Similar interactions in bis(salicylaldiminato)Cu(II) have been recorded.’ The bonding parameters in the latter are very similar to those in bis(acety1acetonato)Cu(I1) which is devoid of proton splittings, except that the in-plane n bonding is more covalent in the case of the salicylaldehydimine than it is for the acetylacetonate.’ Substitution of the methyl groups by phenyl groups in 6-ketoenolates and @-dioximatesre(17) L. G. Van Uitert, W. Fernelius, and E. E. Douglas, J . A m e r . Chem. Soc., 75,2736 (1953). (18) M . Calvin and R. H. Bailes, ibid., 68,949 (1946). (19) D. Kivelson and S. K. Lee, J . Chem. Phys., 41, 1896 (1964). (20) J. J. Fortman and R. G. Hayes, ibid., 43, 15 (1965).
METAL-LIGAND BONDING IN Cu(I1) CHELATES sults in a decrease in the covalent character of the CT bonding and in an increase in the covalent character of the a bonding. Hence, the change in the donor atoms from oxygen to nitrogen on going from bis(benzoy1acetonato)coppcr(II) to bis(benzoy1acetone dioximato)coppcr(I1) results in an intensification of the covalent character of both u and a bonding; the interactions with the protons are therefore facilitated. In conclusion, the effects of substitution, ring size, and the nature of the heteroatom on metal-ligand bonding parameters are as follows. (1) Substitution in the six-membered chelate ring by electron-withdrawing groups cause both (i) a decrease of the covalent character of the Cu(I1)-ligand u bonding and (ii) an increase of the out-of-plane a bonding. (2) Enlargement of ring-size of Cu(I1) chelates from five (a-dioximates) to six (0-dioximates) entails a decrease of the u bonding. (3) Replacement of the oxygen donor atoms around the Cu(I1) ions in P-diketonates by nitrogen donor atoms (p-dioximates) causes an increase of the covalent character of the u bonding. (4) The a bonding has an important role on the stability of the complex compounds. (The increase in stability of the complex molecule is reflected by the decrease in the p2 value of the Schiff base complexes.) (5) Greater covalent character of the a bonding seems to facilitate the hyperfine interactions with the protons.
Experimental Section All reagents used in this work are of "analytical pure" grade. Acetylacetone dioxime (AADO), mp 149", was prepared according to the literature.21 Benzoylacetone dioxime (BADO), mp 86-87', was prepared following the procedure of Muller and Auv,wxZ2 The copper(I1) complexes of P - d i k e t ~ n e sand ~ ~ of adioxime~'~ used in this study were prepared by standard methods available in the literature. The copper(I1) complexes of 6-dioximes were prepared by a modification of Donaruma's procedure :24 a
3833 mixture of 0.01 mol of cupric acetate and 0.02 mol of 0-dioxime in 300 ,ma of toluene was refluxed for 3-4 hr; the toluene-acetic acid azeotropic mixture is distilled and the green oily residue dissolved in benzene and precipitated by addition of petroleum ether. The precipitate was recrystallized from CHCla, affording green crystals. Anal. Calcd for Cu(AAD0)2, C10H1s04N4Cu: Cu, 19.7; N, 17.4. Found: Cu, 19.2, N, 17.5. Calcd for Cu(BAD0)2, C2OH2404N4Cu: Cu, 14.4; N, 12.5. Found: Cu, 14.0, N, 12.1. Schiff base complexes, bis(N-pentylsalicylaldiminato) copper (11) and N ,N '-bis (salicylaldehyde) ethylenediiminocopper(I1) were kindly supplied by Dr. Hana Shechter of our department. The epr spectra of the cupric chelates in CHCla solution were measured with a Varian Model V-4502 X-band spectrometer equipped with a lOO-kc/sec field modulation unit and a Varian V-4257 variable temperature controller, allowing operation between $300 and -200". Chloroform was found to be the best common, nondonor solvent for all the complexes included in this study. No use of other solvents was attempted, because the effect of changing the solvent could be larger than the others due to changes in the ring size and the substituents, etc. The g values were compared with the free-radical standard, diphenyl picrylhydrazil (g = gDPPH = 2.0036). g1 I and A , I were obtained from frozen glass spectra at -160', and the isotropic go and A . values were obtained from solution spectra recorded at room temperature. Optical absorption measurements were taken on a Cary Model 14 spectrometer. (21) C. Harries and T. Haga, Rer., 31, 550 (1898); 32, 1192 (1899). (22) H. Muller and K. Auwers, J . I'rakt. Chem., 137, 81 (1933). (23) K. L. Belford, A. E. Martell, and M. Calvin, J . Inorg. Nucl. Chem., 2 , l l (1956). (24) L. G. Donaruma, Chem. Eng. Data, 9, 379 (1964).
The Journal of Physical Chemistry, Vol. 76,No. 26, 1971
