Electronic absorption spectra and electron paramagnetic resonance of


Publication Date: October 1971. ACS Legacy Archive. Cite this:J. Phys. Chem. 1971, 75, 21, 3355-3359. Note: In lieu of an abstract, this is the articl...
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3355

EPRSPECTRA OF COPPER(II)-AMINE COMPLEXES

Electronic Absorption Spectra and Electron Paramagnetic Resonance of Copper(I1)-Amine Complexes

by M. Dale Alexander,* Patricia C. Harrington,' Chemistry Department, New Mexico State University, Las Cruces, New Mexico 88001

and Alan Van Heuvelen Physics Department, New Mexico State University, Las Cruces, New Mexico

88001

(Received February 26, 1971)

Publication costs assisted by the National Science Foundation and the National Institutes of Health

The synthesis and characterizationof two new copper(I1) complexes, Cu(3,2,3-tet)C12and Cu(3,2,3-tet)(NOs)2, where 3,2,3-tet is NH2(CH2)sNH(CH2)2NH(CH2)3NH2, is reported. The visible absorption spectra of these complexes together with the spectra of the aiialogous complexes of ethylenediamine and 2,3,2-tet, where 2,3,2-tet is NH2(CH2)2NH(CH2)3NH(CH2)2WH2,give information pertinent to the ordering of electronic energy levels of copper(I1) in amine complexes. The absorption spectral data and epr data when used to calculate covalent bonding parameters for the copper(I1) complexes give results which are not satisfactory.

Introduction

(NO&. The complexes were prepared according to the method of Bosnich, et aL3 Complexes of the new tetradentate ligand 4,7-diazaPreparation of Cu(S,S,S-tet)Cl2 and Cu(S,d,S-tet)1,lO-decanediamine, NH~(CH~)BNH(CH~)~SH(CH~)S(N03)2. These complexes were prepared in the same NH2(3,2,3-tet), containing cobalt(II1) have been premanner as that used for the 2,3,2-tet complexes. A n a l . pared recently.2 This ligand has a much greater tenCalcd for Cu(3,2,3)C12: C, 26.56; H, 6.14. Found: dency to form trans complexes with cobalt(II1) and C, 27.17; H, 6.09. Calcd for cU(3,2,3)(N\To3)2: C, presumably other six coordinate metal ions than two 31.12, H, 7.19. Found: C, 31.45; H, 7.41. other linear tetramines which have been under investiPreparation of Cu(en)ZCl and C'u(en)z(NO~)z. The gation: triethylenetetramine (2,2,2-tet) and 3,7-diazacomplexes were prepared by the method of Caglioti, 1,9-nonanediamine (2,3,2-tet). One would predict et aL4 that 3,2,3-tet would be an ideal ligand to coordinate to Visible Spectra. Visible spectra were recorded on a copper(I1) because of this metal ion's well-known Cary Model 14 spectrophotometer. Solid samples ability to bond strongly to four coplanar donor amine were finely ground in Nujol and supported on filter functions. I n our laboratories copper(I1) complexes paper. Visible spectra were also recorded for aqueous of 3,2,3-tet have been prepared and characterized. M ) of each complex. solutions (0.01 Epr and visible absorption spectral data for these 3,2,3Carbon and hydrogen analyses Elemental Analysis. tet complexes and analogous complexes containing mere performed by Galbraith Laboratories, Knoxville, 2,3,2-tet and ethylenediamine (en) have been obtained Tenn. (Table I). The visible absorption spectra are interE p r Measurements. The epr spectra were taken with preted using a ligand field model, and the epr data are a Varian 4502 spectrometer. DPPH was used as a used in theories relating covalent bonding with the reference line. The magnetic field was measured with a experimentally measured results. The results in apVarian Fieldial magnetic field regulator. A Varian plying these theories to the epr data are not satisvariable temperature accessory was used for the lowfactory. temperature studies. Epr spectra were taken of the various copper complexes as chlorides at - 160" (0.004 Experimental Section M copper in 60% glycerin-40% water frozen solution) Reagents. Reagent grade copper(I1) chloride dihydrate, copper(I1) nitrate trihydrate and Eastman White Label N,N'-bis(2-aminoethyl)-l,3-propanediamine (2,3,2-tet) were used without further purification. 4,7-Diaza-l,lO-decanediamine(3,2,3-tet) was prepared by the method of Alexander and Hamilton.2" Eastman Practical ethylenediamine was distilled before use. Preparation of Cu(d,S,%tet)CZZ and Cu(d,S,d-tet)-

(1) NDEA Predoctoral Fellow. (2) (a) M. D . Alexander and H. G. Hamilton, Jr., Inorg. Chem., 8, 2131 (1969); (b) G.R. Brubaker and D. P. Schaefer, ibid., 9, 2373 (1970). (3) B. Bosnioh, R. D. Gillard, E. D. McKenzie, and G. A. Webb, J. Chem. AYOC., A , 1331 (1966). (4) V. Caglioti, C. Furland, G . Dessy, and C. Ibarra, G a m Chim. Ital., 92, 1276 (1962).

The Journal of Physical Chemistry, Vol. 76, No. 91, 1971

M. D. ALEXANDER, P. C. HARRINOTON, AND A. VANHEUVELEN

3356 Table I: Visible Spectra Data hm.x

Compd

Cu(en)z(NOa)Z Cu(en)zClz C~(2,3,2-tet)(NO$)~ Cu(2,3,2-tet)C12 Cu(3,2,3-tet)(NOa)2 Cu(3,2,3-tet)Cl~

hmax

(solid), nm

(HzO), nm

533 550 532 544

547 545 527 524 535 538

540

553

fmax

(Hz0)

57 66

86

I I I and at room temperature (0.01 M copper in water). 2690 2&0 3b80 In these solutions the copper complexes exist as Cu(amine)4(H20)22+. The g values, hyperfine constants, Figure 1. The epr spectra of Cu(amine)a(Hz0)z2+complexes taken at 90°K. The top curve (- - - -) is a calculated spectrum and line widths were evaluated for the three copper which best matched the Cu(en)z(HZO)z+epr spectrum. complexes by matching a computer calculated spectrum to the observed spectrum. The computer spectrum is for tetragonal symmetry and varies the g values (gl I and l O O , I I , I g)l, the hyperfine constants ( A and B ) and the line -C " ~ 3 , 2 , 3 I C l 2 _ _ _ C"12,3,21 CI2 widths (AH11 and A H I ) to get a least-squares fit to the cu Isn lP c1 observed spectrum. The program uses either Gaussian or Lorentzian line shapes. The epr spectra of the frozen solutions along with the calculated spectrum for the C ~ ( e n )complex ~ are shown in Figure 1. The computer-calculated parameters are given in Table 11. There was little difference in the epr spectra of the three different copper complexes, C ~ ( e n ) z ( H 2 0 ) ~ ~ + , Cu(3,2,3-tet)(H20)22+,and Cu(2,3,2-tet) (H20)2+ at 16 18 20 22 24 26 28 30 32 -160" and also at room temperature. A careful Kk visual inspection of the three spectra does indicate Figure 2. Absorption spectra of aqueous soIutions of that the value of g for the 2,3,2-tet complex is slightly Cu(amine)4(H20)22 complexes. less than for the other two complexes. This is also confirmed by the computer evaluation. ogous 3,2,3-tet complexes of copper(I1) using proIt is difficult to assign uncertainties to the values of cedures very similar to those employed in the synthesis the epr parameters given in Table 11. The complexes of the en and 2,3,2-tet complexes. The similarity undoubtedly have some rhombic distortions, whereas between the solid-state spectra of corresponding 3,2,3our program assumes tetragonal symmetry. Thus tet, 2,3,2-tet, and en complexes, as seen in Table I, for example the value of gl is probably an average of supports the notion that 3,2,3-tet also functions as a the values of gz and g,. We have estimated the abtetradentate ligand coordinating in the tetragonal solute uncertainties for any one of the following complane with the anionic ligands coordinating weakly cm-l), plexes: A ( i 6 X lov4 cm-l), B(*6 X perpendicular to the plane as is the case for the (en)z g11( =kO.Ol), and g,(*O.Ol). The uncertainties of the and 2,3,2-tet complexes. parameters for the three complexes relative to each Aqueous solution spectra of the complexes as chloother are somewhat better than this. The lowering ride salts are shown in Figure 2. In aqueous solution of the gli value for the 2,3,2-tet complex as compared the amines retain their "planar" coordination but the with the (en)2 and 3,2,3-tet complexes was observed axially coordinated anions are replaced by water moleconsistently for several separate runs with complexes cules. The spectra are similar especially in the shape in the water-glycerin solvent and also when the comof the bands. Before examining the spectra in detail plexes were dissolved in methanol. consider the ligand field splitting diagram for Cu(1I) in Figure 3, which shows the state splitting for various Results and Discussion symmetries. The complexes under consideration are Electronic Spectra. Copper(I1) complexes of the of pseudo-Dlh symmetry. Procter, Hathaway, and formulation C ~ ( e n ) ~ Xand z Cu(2,3,2-tet)Xzwhere X is Nicholls5 have examined the absorption spectra of an anionic ligand have been prepared previously. 3 , 4 I

,

,

,

I

,

,

,

,

,

,

,

+

I n these complexes all amine groups bond to the metal ion in the tetragonal plane. We have prepared analThe Journal of Physical Chemistry, Vol. 76, No. 21, 1971

( 5 ) I. M. Procter, B. J. Hathaway, and P. Kicholls, J . Chem. SOC.A , 1678 (1968).

3357

EPRSPECTRA OF COPPER(II)-AMINE COMPLEXES Table I1 : Epr Parameters for Cu-Amine Complexes

2;

Complex

Cu(en)dH20)22 C ~ ( 3 , 2 , 3 - t e t ) ( H ~+O ) ~ ~ C~(2,3,2-tet)(H~O)~~+ a

9

Ga

15.8 18.7

x 104, cm-1

- 197

26.0 25.4 25.5

14.1

+

A

HI

- 189

-202

B

x

104,

om-'

- 30

- 32 -31

QI I 2.201 2.204 2.195

Q1

2.040 2.038 2,045

The spectra were best fit using gaussian line shapes.

\

/

\

/

\

'

2 '\

E,

/

/ \

Figure 4. Probable configuration of 3,2,3-tet in copper(I1) complexes.

\ \ \\

*B,,

cluster the energies of the 'E,, 2A~g, and 2Bzgstates more closely while if the 2Bz, state is higher in energy an increase in separation of the energy of the 'AI, state from the 2E, and 2Bzgstates should occur. Thus in the former case a slight decrease in Dqxu should result in a Oh D4h narrowing of the band as well as a shift in the position Figure 3. Ligand field splitting of copper(I1) in of the maximum while in the latter case a broadening various symmetries. of the band is predicted. The half-width of the (en)z band (4900 cm-1) is indeed greater than that of the 2,3,2-tet band (4500 cm-l) and the half-width of the several Cu(en)zXz complexes and have concluded that 3,2,3-tet band (4530 cm-l) is nearly the same as that the energies of the 2Eg,2Bz,,and 2Algstates are similar of 2,3,2-tet. Therefore the results are compatible enough so that the single band obtained in the spectrum with the 2A1, state being lower in energy than the of each of these complexes is due to all three transitions 2Bz,state. with the 2B1, --t 2E, being highest in energy. However, The intensity of the band for the 3,2,3-tet complex is they could not reach any conclusions concerning the considerably greater than those for the (en)z and 2,3,2relative energies of the 2Bzgand 2Algstates. Recently, Smith6 has developed a theory which works well in tet complexes. This is likely a result of a greater deviation from Dansymmetry in the 3,2,3-tet case. Consider explaining the presence of the three d-d transitions in the one band. Furthermore, his theory indicates that the most likely configuration and conformation of 3,2,3the 'A, state is the lowest in energy of the three excited tet in the complex (Figure 4). The six-membered states. Comparison of the spectra of the (en)z, 2,3,2chelate rings are in the preferred chair conformation, tet, and 3,2,3-tet complexes supports this conclusion. the five-membered ring is in the preferred gauche The position of the band maximum decreases in the conformation, and the secondary amine functions are order 2,3,2-tet > 3,2,3-tet > (en)z indicating a slight in the RR(SS) configuration. One will note that the decrease in ligand field strength in the amine plane trimethylene linkage for one six-membered ring is (Dyxg)in the same order. The effect on the relative crowded above the amine plane and that of the other energies of the states is t o first of all decrease the sepis crowded below the plane. This crowding could be aration between the center of gravity of the 'E, alleviated by raising one primary amine function and *Bzg states and the center of gravity of the 2A1g and lowering the other resulting in a pseudo-tetrahedral 'B1g states, and secondly, to decrease the 2E,, 2Bzg distortion. A slight distortion of this type would splitting and to a much greater extent the 2A1,, 2B~g account for the increased band intensity.' splitting. Now, if the 2A1, state is higher in energy than the 2Bzgthe effect of a decreased DqXywould be to (6) D. W. Smith, J. Chem. Soe., A , 1708 (1969).

+

The Journal of Physical Chemistry, Vol. 7 6 , N o . 91, 1971

3358

M. D. ALEXANDER, P. C. HARRINGTON, AND A. VAN HEUVELEN

Epr Spectra. Epr and visible absorption spectra have been used many times to determine covalent bonding parameters for copper(I1) ions in various ligand field environments. The epr parameters gl 1, gL, and A and B and the separation of the d orbitals (Exu - Exa-yZ = AEx,; E,,,,, = AExz,ys) are used to calculate the covalent bonding parameters d ,p12, and p2 and the Fermi contact hyperfine interaction “constant” K (or K ) . a2 measures the covalency of the dxz-ya-a bonds, P I 2 of the d,,-r bonds and P 2 of the dxz,,,-n bonds. These parameters are one for ionic bonds and are 0.5 for pure covalent bonds. One should refer to the original paperssvg for a more detailed discussion of the theory and parameters used therein. There seems to be one fairly consistent trend resulting from the use of these theories, Le., a decrease in the value of gll implies an increase in the degree of covalent bonding. This behavior is predicted by the formali s m * used ~ ~ to evaluate the covalent bonding parameters from the epr spectra. Also the experimentally measured values of gll for copper(I1)-ligand complexes are reduced as the bonding becomes stronger. For example, gi 1 is about 2.3-2.4 for copper-oxygen bonds, 2.2-2.3 for copper-nitrogen bonds, and 2.1-2.2 for copper-sulfur bonds. Table I1 gives the epr parameters for Cu (en)z(H20)2+, Cu (3,2,3-tet)(Hz0)z2+, and Cu(2,3,2-tet)(H20)22+. The epr spectra imply that the covalent bonding is nearly the same in all three complexes but is slightly stronger in the 2,3,2-tet case. This is consistent with the visible spectral results. The general theory8g9used to determine the covalent bonding parameters from the epr and visible absorption spectra does not, however, work well for these complexes. The problem encountered here is a common one, i e . , the amount of T bonding predicted by the theory is greater than expected (see e.g., ref 9-12). Reasonable results can usually be obtained if one assigns the band energies AE,, and AE,,,,, “correctly.” For example, AE,,,,, is often assigned an energy of about 25,000 cm-1 such that the out-of-plane r bonding is negligible (pZ = 1). Usually this band cannot be confirmed because of a very intense charge-transfer band at an energy around 25,000 cm-l. Our experimentally measured and computer-analyzed g values and hyperfine constants for the copper complexes given in Table 11. These results agree quite well with those of Alei, Lewis, Denison, and Morgan,10t13 who have also measured the epr spectra of Cu(en)2(H20)z2+. The epr data for the Cu(en)z complex have been analyzed using the theory of Kivelson and Neimang and the results are reported in Table 111. The epr data were treated in two ways: (1) (line a of Table 111)values of a2,p,2, and 8 2 were calculated using .,~ the the d-d band resolution of Procter, et C C ~ with ordering of energies of the states being 2E, > 2B2, > The Journal of Physical Chemistru, Vol. 76, N o . $1, 1871

2Azg; and (2) (line b of Table 111) we have assumed no in plane n bonding ( i e . , P12 = 1) and have used the data to calculate cy2 and AE,,. Table 111: Covalent Bonding Parameters for Cu(en),( H Z O )+ ~ ~ K a2

(a) 0.80 (b) 0.79 a

= aaPKo,

01%

02

om-‘

0.76

0.73

0.01258

1

AEZv, om-1

AEzs. om-1

17600 19400 24500 (409 nm)

The data from Table I1 have been used along with the follow-

ing values for other parameters: S = 0.093, T = 0.333, h = -828 cm-1, and P = 0.0360 cm-1.

The first treatment gives the covalent bonding parameter /312 = 0.76 which, according to the theory, should indicate appreciable d,,-n bonding. However, there are no orbitals available on the amine donors for such bonding, and thus PI2 should be one. The other n-bonding parameter, p2, is calculated to be 0.73 thus indicating significant d,,, dyz-n bonding. This would have to involve the water molecule coordinating along the x axis perpendicular to the coordination plane of the amines. Bonding along this axis is usually considered to be weak, however; and furthermore, water is a poor n-bonding ligand. There is one other point in considering this treatment of the data. The a-bonding parameter a2 is 0.80 which would indicate more in and out-of-plane n bonding than in-plane Qbonding which seems unusual. I n the other treatment we have assumed p12 = 1.0 and have used eq 5d and e from Kivelson and Neimang (KN) to determine a* (these are the equations for the hyperfine constants A and B ) . Using this value of a2,then the KN equation (5b) (the equation for glI) was used to find a consistent value for the d-d transition corresponding to AE,, (24.500 cm-I). As can be seen in Figure 2, there is no band at 24,500 cm-l (24.5 kK). Using larger AE,, results in p12 > 1, a physically unrealistic value. Thus both treatments indicate that for the complex considered here the theory certainly does not work well. Attempts to use the theory (7) A. S. Brill and G. F. Bryce, J . Chem. Phys., 48, 4398 (1968). (8) A. H. Maki and B. R. McGarvey, ibid., 2 9 , 31 (1958). (9) D. Kivelson and R. Neiman, ibid., 35, 149 (1961). (10) W. B. Lewis, M. Alei, Jr., and L. 0. Morgan, ibid., 45, 4003 (1966). (11) W. Schneider and A. V. Zelewsky, Helv. Chim. Acta, 48, 1529 (1965); A. Van Heuvelen and L. Goldstein, J . Phus. Chem., 72, 481 (1968). (12) S. Antosik, N . M.D. Brown, A. A. McConnell, and A . L. Porte, J . Chem. SOC.A, 545 (1969); H . R. Gersmann and J. D. Swalen, J . Chem. Phys., 36, 3221 (1962); E. M. Roberts and W. 8 . Koski, J . Amer. Chem. Soc., 82, 3006 (1960). (13) M . Alei, Jr., W. B. Lewis, A. S. Denison, and L. 0. Morgan, J . Chem. Phgs., 47, 1062 (1967).

3359

GROWTH PATTERNS OF REACTION INTERMEDIATES by juggling the values of X (the spin-orbit coupling constant) and P , both of which have an ( y - 9 dependence, have also failed.

Acknowbdgment. We are indebted to the National Science Foundation and the Public Health Service for their support of this research.

Growth Patterns of Reaction Intermediates Produced by Self-Radiolysis

of Tritiated Ethyl Iodide at 77"Kla by Paul J. Ogrenlb and John E. Willard* Department of Chemistry, University of Wisconsin, Madison, Wisconsin 66706 (Receiued April 8 , 1971) Publication costs assisted b y the U.S. Atomic Energy Commisswn

The trapped intermediates produced by the self-radiolysis of tritiated glassy and polycrystalline ethyl iodide a t 77°K have been investigated with the aid of their esr and optical spectra. The initial growth rates are proportional to the specific activity but decrease with time a t rates which are different for the different species. I n glassy ethyl iodide, ethyl radicals grow to plateau concentrations which are proportional to less than the first power of the specific activity, the kinetics indicating that the radicals are removed by a process stimulated by the radiation, or that competing processes interfere with their production as the concentration of reaction intermediates increases. The average radical lifetime in the steady state a t 101"K, the highest temperature a t which radicals could be observed in competition with thermal decay, is ea. 15 sec. Between 77 and 95°K the radical decay rates increase with an Arrhenius factor of 2.5 kcal mol-1, but increase much more rapidly above 95°K. Illumination of ?-irradiated samples in the 400-500-nm range removes trapped CzHs radicals. Tritiated glassy samples emit a steady-state luminescence with an intensity approximately proportional to the C2H4, and 1 2 is described. specific activity. The synthesis of CzH4TI from Tz,

Introduction y Irradiation of glassy or polycrystalline ethyl iodide

a t 77°K produces a number of trapped reaction intermediates observable by their visible-uv2 or esr ~ p e c t r a . ~ Ethyl radicals have been clearly identified. Other tentative assignments include ions and charge transfer complexes. The stabilities of the species differ, and there is evidence that some are precursors of others. The present paper reports more extensive investigations of their growth and decay including evidence on photobleaching, temperature effects, and recombination luminescence* 0 Rays from tritium, present as C2H4T1, rather than y rays, have been used for irradiation. This method of irradiation allows convenient continuous Observation at a variety Of dose rates and tures, avoids the interfering spectra which y rays induce in Ohe walls of the sample container, allows an evaluation of LET effects by comparison with y-irradiation studies, and allows easy calculation of varied dose rates from a knowledge of the tritium concentration. Tritium has previously been employed as a source of radiation in solid state radiolyses by Spinks and coworkers using HTO dissolved in HzO, DzO, and organic materials such as CzH60H a t 77°K and by

Damerau and coworkers, who examined the decomposition Products of tritiated trYPtoPhane and dimethyldiethYltin at low

Experimental Section6 Synthesis of C2H4TI. C2H4TIwas synthesized from Tz (half-life 12 years, maximum 0 energy 18.6 keV) using the apparatus of Figure 1. The reaction sequence was Tz Iz+ 2T1, activated by a hot wire, followed

+

(1) (a) This work has been supported in part by the U. S. Atomic Energy Commission under Contract AT(11-1)-1715, by the W. F. Vila8 Trust of the University of Wisconsin, and by NSF and Danforth Foundation fellowships held by P. J . Ogren; (b) present address: Department of Chemistry, Maryville College, Maryville, Tenn. 37801. (2) R. F. C. Claridge and J. E. Willard, J . Amer. Chem. Soc., 88, 2404 (1966). (3) (a) H. Fenrick, S. V. Filseth, A. L. Hanson, and J. E. Willard, {bid., 85, 3731 (1963); (b) H. Fenrick and J . E. Willard, ibid., 88, 412 (1966) ; (0) G. L. Hermann and L. A. Harrah, Technical Report AFML-TR65-333 (1966); (d) R. J. Egland and J. E. Willard, J . Phys. Chem., 71, 4159 (1967). (4) (a) J. Kroh and J. W. T. Spinks, J. Chem. Phys., 34, 1853 (1961); (b) J. Kroh, B. C. Green, and J. W. T. Spinks, Cun. J . Chem., 40, 413 (1962); (c) W. Damerau, G . Lassman, and H. G. Thom, 2. Phys. Chem. (Leipzig), 223, 99 (1963). (5) Further details are available in the Ph.D. Thesis of Paul J. Ogren, University of Wisconsin, 1968. The Journal of Physical Chemistry, Vol. 76,N o . 91, 1971