Tetradentate ligands derived from β-diketones and 1,2-diaminoethane

352

Notes

frequencies and force constants for UF6: K= 3.85 -+0.05 md/,~ [7] and to an earlier determination of K for UF 6 : K = 2.44 md/,~ [8]. My own determination for UF6Cs is: K=2.65-+0.05md/,~[9], very close to the value 2.661md/,~ adopted by Ohwada. The uncertainty affecting )t thus depends much more on the assumptions involved in the model than on the uncertainties affecting the force constants themselves.

SerL,iee de Chinlie Physique Centre d'Etudes Nucldaires de Saclay B.P. No. 2-9119{) Gif-sur- Yeette France

EDGAR SOULIE

REFERENCES

I. M. Drifford and E. Souli6, Communication au Colloque International du C.N.R.S. No. 255. Spectroscopie des #lgments de

I inorg, nucl. Chem, 1978, Vol. 40, pp 352 354. Pergamon Press

transition et des #lPments Iourds dans les solides, Editions du C.N.R.S., Paris (19771 p. 95-104. 2. J. C. Eisenstein and M. H. L. Pryce, Proc. Roy. Soc. (London) A255, 181 (1960). 3. M. Boring, J. H. Wood and J. W. Moskowitz, J. Chem. Phys. 61, 3800 (1974); Comment by N. Edelstein and D. Karraker, J. Chem. Phys. 63, 2269 (1975). 4. M. Boring and J. W. Moskowitz Chem. Phys. Lett. 38, 185 (1976). 5. J. Shamir, A. Silberstein, J. R. Ferraro and M. Choca, J. lnorg. Nucl. Chem. 37, 1429 (1975). 6. T. Hiraishi, I. Nakagawa and T. Shimanouchi, Spectrochim. Acta 20, 819 (1962). 7. R. S. Mc Dowell, L. B. Asprey and R. T. Paine, J. Chem. Phys. 61, 3571 (1974). 8. W. A. Yeranos, Z. Physik. Chemie Neue Folge 45, 72 (1965). 9. E. Souli6, Th6se de Doctoral d'Etat, No. 1801, Orsay (1977L Rapport C.E.A. R-4849.

Printed in Great Britain

Tetradentate ligands derived from/]-diketones and 1,2-diaminoethane (Received 21 March 1977; receit'ed for publication 18 July 19771 #-Diketones have long been used as metal chelating agents for solvent extraction[l]. These bidentate ligands are readily convertible into tetradentate ligands by condensation with 1,2diaminoethane (ethylenediamine-"en")[2]. Although the ligands so formed give rise to stable, square planar complexes with several d-block metals[2], they have yet to be examined as potential solvent extractants. This note reports the preparation of four ligands of type I and their application to the extraction, into chloroform, of their copper(ll) complexes.

R'\ CHa--CH I 1 a/ /C=N

N=C

R2~'/.~,"

./~H~H CR 2

C=O

R3 /

R~

1

O=C

~R 3

I(a) R ] = R 3 = CH 3, R 2 = H (acacen) I(b) R l = CHt; R 2 = H; R 3 = CFI3CH 2 (pracen) l(c) R I = CH3; R 2 = H; R 3 = C6H5 (bzacen)

Rl~ R2C\'--R3 /

EXPERIMENTAL

,"H

C--O

RI

\N--=C~ /

H•..... "/ R 2 O--C~R

3

Solvent extraction studies involved equilibrium of equal volumes of aqueous copper sulphate (10 4moldm-3) with the ligand (0.1 mol dm 3) in chloroform. Aqueous pH, adjusted by addition of HCI or NaOH, was measured after equilibration. The copper(II) content of the aqueous layer was determined spectrophotometrically using the bis(acetaldehyde) oxalydihydrazide reagent[6] and the percentage extraction (E) thence calculated. A second series of extractions was carried out using buffered aqueous phases. Hydrolytic decomposition was followed spectrophotometrically in a Unicam SP 800 with 1 cm cells thermostatted at 25°C -0.1 °.

I(d) R I = R2 = R 3 = CH 3 (meacacen).

The ligands (I) were prepared by condensation of 1,2-diaminoethane with the appropriate diketone [2]. Hexane-2,4-dione [3] and 3-methylpentane-2,4-dione[4] were made by established methods. Products were crystallized from aqueous ethanol and characterised by C, H, N analysis. It was not possible to make the diphenyl derivative ( W = R 3= C6H5; R z= H). For l(b) and I(c), there are two possible products (R k and R 3 may be reversed). Using fragmentation patterns proposed for bzacen[5], the mass spectra of all four ligands were consistent with the structures presented above. In particular, the alternative structures for I(b) and I(c) were eliminated. NMR studies, in deuterochloroform, were also in agreement with structures proposed (MS and NMR carried out by PCMU, Harwell). The presence of a broad intense peak at low r-values (-0.9 to 2.05) suggests that these Schiff bases, in CDCI 3, exist predominantly in the hydrogen-bonded form II, and that the keto-imine structure, I, is effectively absent.

/CHa--CHa

C--N

RESULTS AND DISCUSSION

In the chelate extraction theory developed by Morrison and Freiser[7] and by Stary[l], the extraction (expressed as percentage extraction of metal, E, or as the distribution ratio, D, this being the ratio of the molar concentrations of metal in the organic and aqueous phases) is related to aqueous pH, reagent concentration [HR] o and extraction constant K : loglo D = logm E - logm (100 - E) = log K + npH + n log [HR] o where n=valency of metal (+2 for Cu(II)) and HR is the protonated ligand which forms the neutral extracted chelate MR..

The above is based on a bidentate monoprolic reagent. In these studies, the ligand is diprotic and tetradentate and there is evidence for the formation of l:l metal:ligand complexes[2]. The equilibrium equation thus becomes Cu 2÷ aq + (HzR)o~(CuR)o + 2H +

Notes for which the equilibrium constant (extraction constant) becomes K = [CuR]lille ~ [H]2 [(Tu][H2R],, L * ~ , , and IogK

IogD+21ogH-Iog[H:R] 0

where D = Distribution ratio [CuR]J[Cu] if these two are the main species present in the organic and aqueous phases, respectiveb. Rearranging, log D - log K + 2pH + log [H,RL log E

log!100

E).

[When organic and aqueous phase are of equal volume.] Plots of E vs pH and log D vs pH for I(a) (acacen) and l(b) (pracenl showed good agl:eement with the theoretical model, giving sigmoid curves and stlaight lines, respectively, although the slope of the latter (d(Iog O)ld(pH) was less than 2. the theoretical value. In view of this, the extraction equation above cannot be used to obtain log K and the extraction parameter given below is the pH at ~0% extraction by 0.1 M reagent. hktraction paramelers'

Reagent

Unbulfcred exu:~ction tPHv') I

I(a) acacen l(b) pracen +

2.4 32

I(c) bzacen'

~4

I(d) meacacen

5 56

d(log D) d(pH)

'43

indicates that (stability constant)x (acid dissociation constant) does not vary. Thus, above ~ K t K 2 = constant. Hence, although a series of ligands may change their individual equilibrium constants significantly, the combined extraction constant, K, may vary only slightly. In general, the presence of a buffering agent has a masking effect of the base component of the buffer forms a water-soluble complex with the extractable metal. This shifts the extraction curve towards higher pH. There is an increase of (I.35 for pHi/2 (pH at 50% extraction) for l(a), but a small decrease ((I.18) for l(b). The latter reagent also gives a low value of slope (d log D/dpH), at variance with the theoretical model. However. the decomposition of l(b) must cause uncertainty in these numerical results. The Stary extraction model does not include consideration of the possible interaction of extracted metal with the anions present in solution. [t is known that copper(l[) does react measurably with sulphate, chloride and hydroxide[g] so these ions have a masking effect of small mangnitude, while thal of the hydroxide will increase with pH and is known to be significanl above pH5. The main reason for non-conformance of the reagents l(c) and lld) to the extraction theory 'was their ease of hydrolysis under the experimental conditions. We have studied the hydrolysis behaviour of the four ligands. All hydrolyse at appreciable ra!:es in acidic solution, as do all Schiff bases. Hydrolysis is slower in neutral or alkaline solution meacacen being the most rapid. ]he dependence of initial rate of hydrobsis of meacacen on potassium hy&oxide concentration is shown in Fig. I.

Buffered extractions (PHI'e)°'

1.84 12 2.0 0.5 n.79

2.75 3.00 4.6 5.7 -

d(log D) d(pH) 1.50 1.7

O

08 ,

i

0.6

-

+Values for l(b) and l(c) can only be approximate in view of the hydrolytic decomposition of these reagents. $ln the buffered extractions the aqueous phase contained dichloracetic acid and its potassium salt (total concentration 0.10 m01 d'm ~).

x

%

0.4 J

o c~ o d

Extractions by l(a) and I(b) at low pH indicate the formation M a strong metal complex. The lengthening of the hydrocarbon chain, in l(b). lowers the extraction power of the ligand. This implies that the enhanced hydrophobic character of the metal chelate (increasing extractability) is more than compensated by !he lowered solubility of the ligand in the aqueous layer which decreases the formation of metal chelate (lowering extraction). Fhe extraction of copper(ll) by 0.1 M fl-diketones in benzene is known to show this compensating effect. Values of pilL/z for pentane-2,4-dione (acetylacetone), I-phenyl-butane-l.3-dione (benzoylacetone) and 1,3-diphenyl-propane-l,3-dione (dibenzoylmethane) are 2.9. 3.0 and 2.9[I], despite the increasing hydrophobic nature of the three ligands. The possible effect of change of dissociation constant of ligand and stability constant of metal [igand complex can be examined by applying the method of Morrison and Freiser[7] to this ,,ystem. Thus the extraction constant,

0.2

0D

I

-25

I

I

-2.0

-I 5

-4

- $ C,

Log [ KOH]

Fig. 1. Variation of initial rate of decomposition of meacacen as a function of concentration (mol dm 3~ of KOH. The initial decomposition rates of the chelating agents in aqueous solution of neutral, alkaline (0.1 moldm ~ KOH) and acidic (0.1 moldm 3 HCI) composition are given below. Initial ligand concentratton was 5 × 10 s mol dm

Inilialdecompositionrate (d~/dtl/moldm 'see K = K,,I~K~K2 .

KR K p K R a r e distribution ratios of the protonated ligand, H,R and metal complex, MR; /3 is the stability constant of the complex: K, and K z are the first and second acid dissociation constants H2R. This equation indicates the compensating effect of changes in hydrophibic character of ligand and complex, K~ and K R. For ~tructurally similar ligands, the linear relationship between pK, and log (stability constant) found first by Calvin and Wilson[8],

Reagent

Alkaline

Neutral

Acidic

l(a) l(b) lie)+ l(d)

0.044× 10 ~ Stable Stable . g x I0 s

0.136× 10 ~ 0.21 × 10 ~ Stable l a x 10 s

Rapid Rapid 9. I × I0 ~ Rapid

fBecause of Io.~ aqueous solubilitb 50% aqueous ethanol x~/as used.

354

Notes

The Fig. 1 closely resembles those determined for hydrolytic decomposition of 2-substituted-l,3-diketones[10, 11], suggesting C-C bond breaking rather than C-N. Copper(II) catalyses the decomposition of meacacen, but no simple dependence of first order rate constants on concentration of copper(II) was apparent. This may be due to complications such as the presence of varying amounts of hydroxo-copper(II) complexes in these alkaline solutions. In conclusion, we may state that acacen and pracen are effective extractants for copper(II), but they do not possess marked advantages over/3-diketones. On the other hand, bzacen and meacacen are not satisfactory due to their ease of hydrolysis. It is interesting, in connection with mecacen, to note that analogous a-substituted ~-diketones have also been shown to be unstable with respect to hydrolysis[10]. Thus, while pentane-2,4dione, (acetylacetone) the parent compound of the /Ldiketone extractants, is a stable and inert ligand, many of its derivatives are markedly less inert with respect to hydrolysis and are thus of limited (or zero) value as complexing agents or extractants.

Division of Chemistry and Materials Science H. J. HARRIES G. MOORCROFT Derby College of Art and Technology Kedleston Road Derby DE3 1GB England

J. inorg, nucl. Chem., 1978, VoL 40, pp. 354-355. Pergamon Press,

Chemistry Department University o.[ Leicester Leicester LEI 7RH England

J. BURGESS

REFERENCES 1. J. Stary, Solvent Extraction o.[ Metal Chelates. Pergamon Press, Oxford (1964). 2. P. J. McCarthy, R. J. Harvey, K. Ueno and A. E. Martell, J. Am. Chem. Soc. 77, 5820 (1955). 3. Organic Syntheses 20, 7 (1940). 4. Organic Syntheses 42, 75 (1%2). 5. S. H. H. Chaston, S. E. Livingstone, T. N. Lockyer and J. S. Shannon, Aust. J. Chem. 18, 1539 (1%5). 6. H. Irving and N. S. AI-Niami, J. lnorg, Nucl. Chem. 27, 419 (1965). 7. G. H. Morrison and H. Freiser, Solvent Extraction in Analytical Chemistry. J. Wiley, New York (1957). 8. M. Calvin and K. W. Wilson, J. Am. Chem. Soc. 67, 2003 (1945). 9. Stability Constants. Chemical Society (1%4). 10. J. L. Ault, Ph.D. Thesis, University of Leicester (1973). 11. J. P. Calmon and P. Maroni, Bull. Soc. Chim. (France) 9, 3761 (1968).

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The preparation of some N,N-dialkyldithiocarbamates of titanium(Ill) (Received 2 May 1977; received for publication 15 July 1977) Although, many N,N-dialkyldithiocarbamato derivatives of transition metals are known, little work has been done on the lower oxidation states of the early transition metals. Bradley and Gitlitz[1] have reported the preparation of the tetrakis N,Ndialkyldithiocarbamates of titanium(IV) and zirconium(IV). Mixed complexes, which contain both the dithiocarbamato and chloride ligands are also known[2,3]. Courts and Wailes[2,4] reported the preparation of some dicyclopentadienyl, dithiocarbamato complexes of titanium(Ill): to date these are the only reported titanium(Ill) complexes which contained the dithiocarbamate ligand. In this publication we wish to report the synthesis of some titanium(Ill) tris-N,N-dialkyldithiocarbamates; produced by the reaction of the N,N-dialkyldithicarbamate sodium salts with titanium trichloride, in ethanol, in a three to one -atio. 3R2NCSzNa + TiCI3

EtOH

~R2NCS203Ti.

All materials were handled under dry, oxygen free argon, using Schlenkware for all synthetic work. Dithiocarbamates were purchased from Eastman Organic Chemicals and R&K Chemicals of Hartville, Ohio. Titanium trichloride was purchased from AlfaVentron and used without further purification. The following example of the preparation of Ti(S2CNBu2)3 is typical of the method used for the synthesis of these N,N-

dialkyldithiocarbamato complexes. Titanium trichloride (0.0112 mole) in anhydrous ethanol (15 ml), was added dropwise, with constant stirring, to a solution of sodium N,N-dibutyldithiocarbamate (15.23 g, 0.0670 mole) in 75 ml of ethanol. The resulting solution/was filtered to remove sodium chloride and then stored at - 4°C for 72 hr. The mother liquor was removed and the crystals were dissolved in benzene and filtered to remove excess ligand, which is not soluble in benzene. The solution was then evaporated to dryness under vacuum. The analytical data and melting ranges of the compounds are summarized in Table 1. The compounds are yellow-green and quite air sensitive. They are considerably soluble in ethanol, benzene, chloroform and dichloromethane; decomposition occurs upon standing in dichloromethane. The solubility of the complexes increased with increasing chain length, the methyl complex being almost totally insoluble in all solvents. The IR data are listed in Table 2 and the presence of bands which may be attributed to the Ti-S stretching mode, indicate that a complex is formed. A single C-S band is observed at 995 cm t intermediate to the C-S single and C-S double bond stretching frequencies, this may be considered to be indicative of a chelated ligand. Perhaps the most characteristic band in the spectra is the so-called "thioureide ion" band, which falls in the 1600--1400 cm ~ region. Chatt et al.[5] reported, after extensive

Table 1. Analytical data and melting ranges

Compound Me2NCS2)3Ti Et2NCSz)3Ti Pr2NCS2)3Ti Bu2NCS2)3Ti

(% Theoretical) Ti C H 11.73 9.73 8.31 7.25

26.47 36.58 43.75 49.09

4.41 6.09 7.28 8.17

N

S

10.28 4 7 . 1 0 8.53 3 9 . 0 6 7.28 33.37 6.36 29.12

Ti 11.69 12.51 8.75 7.73

C 27.41 38.57 44.01 48.47

H

(% Found) N S

4 . 4 3 10.29 6.84 6.75 7.71 8.93 8.53 6.34

46.97 34.42 28.83 28.20

M.P. 227°C decomp. 168-170°C 125-127°C 64-67°C