Metal chelates of some simple diamides

J. inorg,nucl.Chem.,1968.Vol.30, pp. 2679 to 2687. PergamonPress. Printedin Great Britain

METAL

CHELATES

OF

SOME

SIMPLE

DIAMIDES*

MARY L. G O O D Louisiana State University in New Orleans, Lakefront, New Orleans, Louisiana 70122 and T. H. S1DDALL, Iil Savannah River Laboratory, E. 1. du Pont de Nemours and Co., Aiken, South Carolina 29801

(First received 23 October 1967: in revised form 25 March 1968) Abstract - A series of metal chelates of tetramethyloxalamide. H:3C

\

/

CH~:~)

N--C--C--N H:~C

CH3

and tet ramethylsuccinamide, H3C

C H ,;) /

\ N--C--CH2--CH2~--N H:jC

/

II

II

O

\

O

CH:~

have been prepared and characterized. The properties of these complexes are compared with those found previous for the tetramethylmalonamide, H3C

O \N~--C

/

O

CH,~,

H2-~--N /

\

HaC

CH2

ligand. The trends observed for the several compounds formed are rationalized in terms of steric interactions, metal ion size and size of chelate ring formed.

IN RECENT years a great deal of emphasis has been placed on the complexes formed between certain metal ions and a large variety of amides and other ligands containing oxygen bonded to sulfur or phosphorous [ 1]. Most of the work has been confined to the determination of the position of monodentate prototypes in the spectrochemical and nephelauxetic series and to a rationalization of the positions in terms of steric and inductive effects [2, 3]. Very recently the complexes formed by a bidentate ligand, tetramethylmalonamide (TMM) have been studied[4]. *The information contained in this article was developed during the course of work under Contract AT(07-2)-I with the U.S. Atomic Energy Commission. I. J. H. Bright, R. S. Drago, D. M. Hart and S. K. Madam Inorg. Chem. 4, 18 (1965) and references cited therein. 2. R. L. Middaugh, R. S. Drago and R. S. Niedzielski,J.Am. Chem. Soc. 86, 388 (1964). 3. R.S. Drago, D. W. Meek, M. D. Joesten and L. LaRoche, Inorg. Chem. 2, 124 (1963). 4. W. E. Bull and R. G. Ziegler, Inorg. Chem. 4, 689 (1966). 2679

2680

MARY L. G O O D and T. H. S I D D A L L , I l l

Infra-red spectra of the complexes indicated that coordination occurred through the two carbonyl oxygens of the diamide. From the visible spectra of the compounds in nitromethane solutions, the ligand field parameters, Dq and fl, were calculated for the Cr(III), Co(II) and Ni(II) complexes. The values obtained would place T M M in the spectrochemical series near water. The Dq values were found to be somewhat larger than those for monamides and the increase is attributed to the "chelate" effect, by which extra stability is gained from the formation of the six-membered chelate ring. Since the usual "rule of thumb" is that the stability of similar chelate systems (where resonance does not produce pseudoaromatic systems) decrease in the order five > six > seven-membered ring system, some basic information should be obtained by preparing and characterizing the corresponding tetramethyl oxalamide (TMO) and tetramethyl succinamide (TMS) compounds. EXPERIMENTAL The diamides were prepared by methods reported previously[5]. All other chemicals were commercial, reagent grade materials and were used without subsequent purification. The complexes were prepared by the general method outlined by Bull and Ziegler [4]. All attempts to prepare the corresponding Cr(III) compounds failed. Both the method used by Bull and Ziegler[4] and the method suggested by Madan and Denk [6] were tried. A dark green, intractable oil was obtained for both the T M O and TMS preparations. The compounds prepared and their physical characteristics are given in Table 1, with pertinent remarks about their preparation and isolation. Metal analyses were done by E D T A titrations and C, H, and N analyses were performed either by Galbraith Laboratories or by the Analytical Chemistry Division of the Savannah River Laboratory. Visible and near i.r. spectra were obtained for both nitromethane solutions and Nujol mull dispersions (on filter paper) on a Cary Model 14 recording spectrophotometer. Infra-red spectra were made on KBr pellets using a Perkin-Elmer Model 521 recording spectrophotometer, Estimates of molecular weights were made with nitromethane solutions of the compounds in an Arthur H. Thomas Company Model 12 isothermal osmometer. Calibration against standard, known compounds dissolved in nitromethane was carried out before each run. Magnetic susceptibilities were obtained with solid samples in Pyrex tubes (precalibrated with mercury(II) tetrathiocyanatocobaltate(II)[7]) on a Gouy balance. Measurements were made at three field strengths and molar susceptibilities were corrected using tabulated diamagnetic corrections for the ligand components [8]. The effective magnetic moment,/zaf., was calculated for ambient temperature using the Curie law. RESULTS

AND

DISCUSSION

Visible and near-i.r, spectra The spectra obtained for both the nitromethane solutions and the Nujol mulls are given in Table 2. A comparison of the spectra of the TMO and TMS complexes with that reported for TMM and for octahedral aquo complexes indicates that apparently an octahedral field is present in all of the compounds studied. The spectra in all four cases is remarkably similar. This implies weak co-ordination of the perchlorate ion in the Cu(II) complexes of TMM and TMS (possibly in the TMO case as well) and in the Ni(II) complexes of TMS. Complexes of 5. 6. 7. 8.

T. H. Sidall, III and M. L. Good,J. inorg, nucl. Chem. 29, 149 (1967). S. K. Madan and H. H. Denk,J. inorg, nucl. Chem. 27, 1049 (1965). B. N. Figgis and R. S. Nyholm, J. chem. Soc.4120 (1958). B. N. Figgis and J. Lewis, in Modern Coordination Chemistry, (Edited by J. Lewis and R. G. Wilkins), pp. 400--454. Interscience, N e w York (1960).

Metal chelates of diamides

2681

this type have been reported previously for Cu(II) and Ni(II) complexes of bipyridyl and 1,10-phenanthroline [9]. A direct comparison of the Dq values for the various Ni(II) complexes is not completely valid, nor are the comparisons for the Co(II) and Cu(II) T M O complexes, since mixed ligands are involved in some cases, but the Dq values do show some interesting trends. A decrease from 860 for the T M M complex to 820 for the mixed TMS complex is as expected going from a six to a seven-membered chelate ring. The value for T M O complex is also lower than that for TMM, indicating some destabilizing factor in the five-membered ring system. All of the values, however, are higher than the value of 769 for N,N-dimethylacetamide, indicating a "'chelate effect" in all three systems. From the spectra reported in Table 2 it seems that the diamides are weak ligands with Dq values about the same as water.

Infra-red spectra The values for the stretching frequency of the C-~---O bond in the isolated ligands and in the metal complexes are given in Table 3. The data indicate that bonding is occurring through the C ~ O group in all cases, although it appears to be weak in the case of the oxalamide chelates. In these molecules the major spectral change is in the character of the bands (from a symmetrical doublet for the free ligand to distorted singlets for the metal complexes) rather than in hmax. shifts. As reported by Bull and Ziegler, the band position for the metal malonamide complexes is shifted to significantly lower frequencies in all cases, implying a relatively strong metal ---C----O interaction. The shifts reported here for the succinamide complexes are of the same type, but somewhat smaller than those for the malonamide compounds. These data lend further support to the idea of a stability order: oxalamide < malonamide > succinamide. The region of the spectrum from 1100 to 1200 cm -1 was investigated for the N i - T M S and C u - T M S complexes in hopes of discerning evidence of perchlorate bonding in these species. Some changes in the spectrum on complex formation were observed, however, no conclusions could be drawn because of overlapping amide bands in this region.

Molecular weights and magnetic susceptibilities The approximate molecular weights of several of the diamide complexes are listed in Table 4 along with calculated monomer molecular weights based on the stiochiometry indicated in Table 1. For those compounds that were easily prepared as slightly hygroscopic crystalline solids and which formed normal tris complexes, the experimental molecular weights indicate a normal, monomeric species. However, the compounds that were crystallized with difficulty from a viscous oil to produce a very hygroscopic solid and which analyze as bis complexes give apparent molecular weights that imply polymerization. For the N i - T M S and C u - T M S complexes the polymerization may occur through perchlorate bridges or through the sharing of one or more diamide molecule(s) by two metal ions. Our data does not allow a definite description of the structure of these species. 9. N. T. Baker, C. M. Harris and E. D. McKenzie, Proc. chem. Soc. 335 (1961 ).

[Pb(TMO)2](CIO4)2

TMO

TMS

TMO

TMS

TMO

TMS

Pb(II)

Pb(II)

Cd(II)

Cd(lI)

Co(II)

Co(II)

[Co(TMS)2(H20)2](C104)~

{[Co(TMO)2~0-I~O)~I(C10,)~}

[Co2(TMO)5(H20)4](C104)4 or

[Cd(TMS)3](C104)2

[Cd(TMO)a](CIO4)2

[Pb(TMS)2](CIO4)2

[Mn(TMS)3](CIO4)~

Mn(II) TMS

Compound formed

[Mn(TMO)3](CIO4)2

Ligand

Mn(II) T M O

Metal

132-134

163-168

88-90

271

150-155

284

160

325-330t

M.P. *(°C)

Calc: Found:

Calc: Found:

Calc: Found:

Calc: Found:

Calc: Found:

Calc: Found:

Calc: Found:

Calc: Found:

30-10 30.80

27.53 27.59

34-81 33.85

31-08 29-00

25-60 25.86

20.75 20.99

37-41 37.21

31.50 31-14

C

5.64 5.67

5.23 5-84

5.80 5.60

4-88 4.86

4-27 4.24

7.48 3.51

6-23 6-56

5.29 5.51

Analysis H

8-78 8.82

10-70 9.00

10-15 9.80

10.49 10.53

7.46 6-94

8.07 7.76

10.91 10.10

12-25 11.63

N

Comments

9.23 light pink crystalline solid, hygroscopic. 8-94 Isolated from reaction mixture on addition of ether with stirring. See discussion of this compound in text.

9.01 lavender solid, very hygroscopic. Isolated 9-24 from reaction mixture as an oil on ether addition. Solidified only after several hours of stirring in many successive ether washes. Product may be mixture.

13.58 white crystalline, slightly hygroscopic solid, 14.08 easily isolated.

15.11 white crystalline slightly hygroscopic solid, 15.09 isolated by ether addition to reaction mixture.

27.61 white crystalline, hygroscopic solid which is 27.26 easily isolated.

29.84 white crystalline, hygroscopic solid which is 29-27 easily isolated.

7-13 white crystalline solid with faint ink color: 7.04 slightly hygroscopic, easily isolated.

8.00 white crystalline solid; slightly hygroscopic, 8-04 easily isolated.

Metal

Table I. Physical characteristics of diamide complexes

,r.

0 0

K 2, .<

G~

TMS

TMO

TMS

Cu(ll)

Ni(II)

Ni(I1)

[ZndYM O)~(H20)2](C100.~

TMO

TMS

Zn(iI)

Zn(ll)

Calc: Found:

Calc: Found:

-50

50-55~c

Calc: Found:

Calc: Found:

175-178

214

Calc: Found:

Calc: Found:

Calc: Found:

Calc: Found:

129

102

45-50

125-126

*TMS melts sharply at 70°C and TMO melts at -55°C. CApparent decomposition. :~Appears to dissolve in own water of hydration rather than melt.

[Zn(TMS):~](C104)..,

{[Zn(TMO) 2~H2OI(CIO4)e}

or

[Fe(TMS)3](C104):~

{[Fe(TMO):½(H20)zI(CIO0s}

or

[Fe2(TMO)5(HzO)4I(CIO4)s

[Ni(TMS)2](CIO4)z

[Ni(TMO)z(H20)2]CIO4'2 H20

or

[Ni(TMO)2(H20)4](CIO02

[Cu(TMS)2](CIO4)z

[Cu(TMO).,](CIO,)=,.2H.,O

or

[Cu(TMO)2(H20)2](C104)2

Fe(lll) TMS

Fe(lll) TMO

TMO

Cu(II)

36.91 35.15

33-10 32.55

23.99 23.16

31.92 32.25

22.99 22.38

31.66 31-70

24.56 24.70

6.15 6.24

5.52 5.57

4.56 5.73

5.32 5-59

5.18 5.06

5.28 5.40

4-81 5.21

10-76 9.63

9.67 10-70

7.99 8-50

slightly

8-37 white crystalline solid, very hygroscopic. 8.55 Isolation was similar to N i - T M S and C u TMS compounds.

10-20 white crystalline solid, very hygroscopic. 10.70 Isolation problems were similar to those for Co-TMO. Compound is probably an impure mixture of product and starting materials.

6.41 bright yellow crystalline solid, 6.11 hygroscopic, easil} isolated.

7.44 orange brown solid, very hygroscopic. 8-26 Isolated with great difficulty, problems of isolation similar to that of C o - T M O compound. Product may be mixture.

9-75 pale yellow-green solid, very hygroscopic. 8.32 Compound ppt. from reaction mixture but was difficult to filter. Similar to C u - T M S case above.

9.50 pale green solid, very hygroscopic. Isolation 9.24 problems similar to those for C o - T M O above.

9-07 8-44

9.31 8.41

1 0 . 4 7 pale blue solid, very hygroscopic. Compound 1 0 . 2 8 ppt. from reaction mixture but was difficult to filter.

10-83 light pale blue solid, very hygroscopic1 0 - 5 9 Isolation problems same as for C o - T M O case above.

9.23 8.80

9.55 9.36

o¢

3 r,'

r~

2684

M A R Y L. G O O D and T. H. S I D D A L L , III

Table 2. Near i.r. and visible spectra of some diamide complexes Compound

[ Co(TMO)2½(H~O)~]

Nujol mull ~". . . . (cm -~) 7620 16500(sh) 19600 21700(sh)

Nitromethane solution, h . . . . (cm -~)

Emax.

1. mole-lcm -~)

7790 15300(sh) 19100 20800(sh)

4.3 6.0 23.5 18.3

[Co(TMM)3](CIO4)2*

8330 16900(sh) 18900 20400 21400

7860 17000-16000(sh) 18700 20140(sh)

8.7 3.0 29"1

[Co(TMS)~(H20)~](CIO4)~

7250 14300(sh) 18500 19600(sh)

7200 15300(sh) 18700 20600(sh)

7.9 4.1 28.2 17.0

8000 19600 21600

Co(H20)6s+'J"

[Cu(TMO)2(H20)z](CIO4)2 [Cu(TMM)~](CIO4)2 * [Cu(TMSh] (CIO4)2 [Cu(H20)6] 2+ t [Ni(TMO)2(H~Oh](CIO4)2

[Ni(TMM)3](CIO4)z*

[Ni(TMS)2](CIO4)~

[Ni(H~O)6] 2+t

12100 15100 12650 8450 13700 14500(sh) 25600(sh) 8850 13700(sh) 15100 26600 8060 13200 14500(sh) 24400

12100 14300 12500 12000 8400 13700 14900(sh) 25700(sh) 8600 13600(sh) 14800 25200 8200 13500 14700 24500 8700 14500 25300

1.3 4-8 2.1 36.3 26.6 48.5 l 1.0 2.4 2.2 2.1 7.0 12.1 7.4 8.6 22.3 7-7 9.3 20.3 1-6 2-0 4.6

*Data from Ref. [4]. t D a t a from B. N. Figgis, Introduction to Ligand Fields, pp. 217-228. Interscience, New York (1966). Spectra is for ion in aqueous solution.

Metal chelates of diamides

2685

Table 3. C------OFrequencies for diamides and their metal complexes System free T M O

Stretch, cm -~ ~ 1655 11635 1645 1627

Av, cm ~

Remarks

--

Sym. doublet

--

Sym. broad singlet Sym. broad singlet

/

free TM M free T M S Mn(II) T M O Mn(lI) T M M Mn(ll) T M S

1635 1616 1605 1640(sh)

-0 -29 -22 + 13

Pb(l I) T M O

1635 1660(sh)

-0

Pb(ll) T M M

1645 1610 1582

0 -35 --63

F r o m Ref. [4]

Pb(II)TMS

1625 1590

0 --37

Band is split into two definite bands

Cd(ll) T M O

1635 1660(sh)

-0

Cd(II) T M M CD(ll) TMS Co(ll) T M O

1610 1625 1620(sh) 1630 1650(sh)

--35 -2 --5

Band is essentially sym. with disappearance of 1653 cm -~ band

Co(ll) T M M Co(l I) T M S

1631 1600 1640

-14 --27 --13

From Ref. [4] Peak is unsym, with extra intensity at lower wave number

Cu(lI) T M O

1640 1650(sh)

Cu(II) T M M Cu(lI) T M S Ni(II) T M O

~0

Sym. singlet, 1655 cm -~ band disappears From Ref. [4] Unsym. band with shoulder

Unsym. band with shoulder

Band is unsym, but almost resolved into two components. From Ref. [4] Essentially sym. band

Band is unsym, with shoulder at 1650

1603 1620 1635 1660(sh)

-42 -- 7

1621 1605 1640(sh)

--24

From Ref. [4]

--22

Broad, almost sym. band

Fe(lll) T M M Fe(lll) T M S

1600 1600 1640

--45 --27 +13

From Ref. [4] Band is split into two nonsym, peaks

Zn(ll) T M O

1635

-0

Zn(ll) T M O Zn(II)TMS

1621 ~1600(sh) 11630

Ni(lI) T M O Ni(ll) T M S

~0

--24 -27

From Ref. [4] Broad sym. band Band is very similar to the Cu(II) T M O case

Sym. band is obtained with disappearance of 1655 cm -~ band From Ref. [4] Very unsym, band with shoulder on low freq. side

2686

MARY L. G O O D and T. H. S I D D A L L , l l I Table 4. Molecular weights of diamide complexes

Empirical formula

Experimental Calculated mole weight monomer weight

[Mn(TMO)a](CIO4)2 [Mn(TMS)a](CIO4)2

690 600

686 771

[Cd(TMO)3](CIO4)2 [Fe(TMS)z](CIO4)a [Co~(TMO)s(HzO)4](CIO4)2

720 860 600

744 871 1309

[Co(TMS)2(H~O)~](C104)2

860

578

[Ni(TMO)~(H~O)4](C104)2

500

618

[Ni(TMS)2](CIO4)2

1500

602

ICu(TMS)2](CIO4)z

1050

607

Remarks -Shows some tendency to dissociate in CH3NO2 --This value implies that the unit in solution is Co(TMO)2 (HzO)2 State of polymerization is inclusive Shows some evidence of dissociation in CH3NO2 This material appears to be polymeric This product is at least dimeric

The [Co(TMS)~(H20)2](C104)2 complex is particularly interesting in that if it is heated in a vacuum oven to 60°C, it can be dehydrated to a bright blue complex with a spectrum in nitromethane solution (and in the solid) x,ery similar to that for COC142-. Apparently the two water molecules can be lost with a change in configuration from octahedral to tetrahedral. The reaction is reversible and the resulting blue solid hydrates quite readily to give the pink octahedral complex back again. The mixed ligand complexes (with co-ordinated water) are similar to the compounds isolated by Gentile and Shankoff[10] from the reactions of diacetamide with several transition metal ions. The effective magnetic moments for several of the complexes in Table 1 are given in Table 5. The values of the corresponding T M M complexes from Bull and Ziegler[4] are given for comparison. Each value corresponds to the average /xeff. of data taken at three different field strengths. Magnetic data for the Cu(II) compounds or for the T M O complexes of Fe(IlI) and Ni(II) could not be obtained with our equipment because of the extremely hygroscopic nature of these materials. The magnetic moments observed are characteristic of spin-free, weak field octahedral complexes in all cases, including the compound, [Ni(TMS)2] (C104)2. This is further evidence that the perchlorate ion is weakly co-ordinated in this compound. CONCLUSIONS

The instability of the five-membered ring oxalamide chelates as compared to the six-membered ring systems requires explanation. If a scale model of tetramethyloxalamide is constructed, it can be seen that the methyl group on one of the nitrogen atoms interferes with the methyl group on the second nitrogen so long as the amide groups remain coplanar. To relieve the steric strain thus introduced, the 10. P.S. Gentile and T. A Shankoff, J. inorg, nucl. Chem. 28, 1283 (1966).

2687

Metal chelates of diamides

amide groups apparently rotate out of plane with respect to one another and thereby reduce the electron delocalization and the bonding ability of the carbonyl oxygen atoms. This conclusion is consistent with results obtained when high molecular weight diamides are used as extractants for actinide and lanthanide ions from aqueous solutions [5]. Evidence for the non-coplanar model of the oxalamides has been reported from N M R studies of these molecules at low temperatures [5, 11]. The apparent strong influence of steric effects is not unique to the diamides, Drago and coworkers[2, 3] have found steric interaction to be the overriding effect when ligand-field parameters of transition metal complexes are correlated with inductive effects, basicities and steric factors of certain amide ligands. Table 5. Magnetic moments for diamide complexes

Compound

Magnetic moment

Source

[Mn(TMO)s](CIO4)z [Mn(TMM)3](CIO4)2 [Mn(TMM)3](C104)z

6-02 6.03 5.66

This work Ref. [4] This work

IFe(TMM)3](C104)3 [Fe)TMS)3](CIO4)~

6-05 5-29

Ref. [4] This work

[Co(TMO)2½(H20)2](CIO4)z [Co(TMM)3](CIO4)2 [Co(TMS) (H20)z](C10~),,

5.54 5.03 4-98

This work Ref. [4] This work

[Ni(TMM)s](C104)2 [Ni(TMM)2](C104)2

3.31 3-21

Ref. [4] This work

The variations observed in thermal stabilities of the complexes reported here can be rationalized in terms of metal ion size and the ionic character of the the compounds. Melting points of over 250°C are reported for the oxalamide complexes of Mn 2÷, Pb 2÷ and Cd 2÷, the metal ions of larger ionic radii (from about 0.80 ,& for Mn 2+ to 1.21 ,h for pb2+). Apparently these large ions can accommodate the oxalamide ligand and form an essentially ionic complex, which is crystalline and easily isolated from the reaction mixture. The remainder of the compounds, particularly those which show evidence of polymerization, are partially covalent compounds with lower melting points and less well-defined crystalline structures. This is consistent with the lower ionic radii for the remainder of the metal ions investigated (all of which have radii of approximately 0.70 A). Similar correlations have been given by Pitts and coworkers[12] for transition metal complexes of bidentate phosphinate ligands. They reported the formation of four-membered chelate rings with the larger ions and found a direct relationship between the thermal stability and the ionicity of the complexes. 11. T. H. Siddall, 111 and M. L. Good, Bull. chem. Soc. Japan 39, 1619 (1966). 12. J. J. Pitts, M. A. Robinson and S. I. Trotz, The characteristics o f Several N e w Metal Phosphinate Complexes, paper No. 80, in the Inorganic Division of the 153rd National ACS Meeting in Miami, April 1967.