A study in the complex formation of iminodiacetic acid and nitrilotriacetic acid with aluminium, chromium and beryllium ions

Notes The complexes isolated were high melting, diamagnetic solids, insoluble in benzene and chloroform but soluble in acetone and acetonitrile giving deep purple solutions. They appeared to be stable to atmospheric moisture in the solid state though sensitive when in solution. Tris(dithiolene) complexes of titanium (IV) COMPLEX ~

Colour

[Et~NH]z[Ti(S,,C~H~CHO] [Et:NH=]: [Ti(S~C,H~CH0]b

dark green black

Yield

%m.p.°C

95 93

228-229 259-260

(a) By analysis, (b) This compound has been reported previously[4] (yield 96°~; m.p. 261-263°C). The IR spectra of the two complexes isolated were found to be virtually identical except for differences due to the different cations. Both spectra showed the absence of thiol protons in the compounds and also the absence of the strong absorptions around t120 and 1005 cm ~ present in the spectrum of titanium tetraisopropoxide which have been assigned to C-O stretching modes[5]. Absorptions at 2463, 2360 and 800 cm ~ in the complex prepared using diethylamine indicate the presence of the diethylammonium cation in that compound[6]. Comparison of UV-visible spectral data for the two complexes with data for previously reported complexes containing the dianion shows good correlation[4]. Marked similarities were also noted

407

when far infrared data in the Ti-S stretching region were compared with literature data [4]. Thus, the evidence outlined above has led to the assignment of the formulae of the complexes isolated as [Et2NH.q2 [Ti(S2G,H~CH3h] and [EhNH].~ [Ti(S.,C6H3CH0~] respectively. This, therefore, represents a novel synthetic route of comparative simplicity to some tris(dithiolene)titanium (IV) complexes, a route which could perhaps also be applied to the synthesis of the anologous tris(benzenethiolato) complexes which have also been reported [4].

Department of Physical Sciences and Technology Polytechnic of the Sottth Bank Borough Road London S E 1 0 A A England

JOHN JONES JOSEPHINE DOUEK

REFERENCES 1. E. I. Steiffel, Z. Dori and H. B. Gray, J. Amer. Chem. Soc. 8t), 3353 [ 1967). 2. E. J. Wharton and J. A. McCleverty, J. Chem. Soc. IA) 2258 (1969). 3. T. A. James and J. A. McCleverty, J. Chem. Soc. (A) 331811970). 4. J. L. Martin and J. Takats, Inorg. Chem. 14, 73 [ 1975). 5. C.G. Barraclough, D. C. Bradley, J. Lewis and I. M. Thomas, J. Chem. Soc. 2601 (1961). 6. L. J. Bellamy, The IR Spectra of Complex MoleculesMethuen, London.

inorg,nucLChem.VoL43,pp.407-409 PergamonPressLtd.,1981. Printedin GreatBritain

0022-190218110201~4071502. 0010

J.

A study in the complex formation of iminodiacetic acid and nitrilotriacetic acid with alurainium, chromium and beryllium ions (Received 17 September 1979: received for publication 3 July 1980)

Amino compounds have a wide range of applicability including biological, pharmaceutical, industrial and other chemical uses, and they are well known for their complexation tendencies. There have been some potentiometric studies of the complexation of IDA and NTA with metal ions[I,2]. But no attempt, however, appears to have been made to determine the protonligand stability constants of these ligands and the formation constants of their chelates with AI(III), Cr(lll) and Be(lI).

EXPERIMENTAL Materials. Solutions of AI(NO3h'9H20 (E. Merck), CrlNO3h" 9H~0 (Orianal Schiapparelli Torino, Italy), BeSO4-4H,O (Reachim USSR), IDA (Sigma Chemical Company, USA), NTA (BDH) and NaCIO4- H20 (Riedel) were prepared by dissolving the requisite quantities in conductivity water. Perchloric acid solution was prepared by diluting a calculated amount of the acid with conductivity water and was standardized volumetrically. The stock solutions of AI(NO3h.9H20 Cr(NOah ' 9H20 and BeSOa ' 4H20 were standardized using appropriate methods [31. Apparatus. A Philips pH-meter model PR 9405 M equipped with glass and calomel electrodes was used for pH measurements after standardising with buffer solution of pH 4 and 7. Procedure. The titrations were carried out in an inert atmosphere of nitrogen in a specially designed double-walled beaker employing the Bjerrum-Calvin pH titration technique [4, 5] as used by Irving and Rossotti[6]. Three solutions (total volume 50ml in each caset were prepared as follows and were titrated against carbonate-free standard caustic soda solution.

For the IDA System: 6) 4× 10 3M perchloric acid, (ii) 4× 10-3M perchloric acid + 3× 10 3M ligand and tiii)4x 10-3M perchloric acid + 3 × 10 3 M ligand + 5 × 10-a M metal ion solution. For the NTA System: (i) 2x 10-2M perchloric acid, lii) 2× 10-2 M perchloric acid + 3 × 10-3 M ligand and (iii) 2 × 10 2 M perchloric acid + 3 × 10 3 M ligand + 5 x 10 4 M metal ion solution such that the concentration of the common ingredients were identical in the different cases. An appropriate quantity of sodium perchlorate (2.0 M) was added to maintain the desired ionic strength. Mixtures (i), (ii) and (iii) were individually titrated against standard NaOH solution (0.4ml. The ratio of the metal and ligand was always 1:6. Three titration curves obtained were referred to as /A) acid titration (Bt ligand titration and IC) complex titration curves.

RESULTSANDDISCUSSION The pH range investigated for the various systems is as under:System AI(III)-IDA Cr(I1D--IDA Be(II)-IDA AI(III)-NTA Cr(III)-NTA Be(II)-NTA

pH range 2.95-4.75 3.45-6.10 4.40--5.20 2. I0-4.55 2.60--5.25 3.55-4.85

The proton-ligand stability constants and stepwise formation constants were determined as follows:

408

Notes Table I. Proton-ligand stability constants of IDA and Nta ~emDer acur O

~ g a~d

a5°c

45%

O.iP~

0.2H

0.3h

0.4/4

0.i~I

O.ik

log K1H

9.73

9.66

9.~

9.57

9.25

9.07

log ~

2.70

2~

2~

2e

2.66

2.~

log/;~

~ . 43

12.19

12.08

n . 98

n . 9~

11. ~

log ~i

9.75

9.68

9.58

9.42

9.53

9.47

log r~

2.~5

2.39

~.34

2.~a

2.,~

2.56

log r~

1.73

1.~

L64

1.59

178

~.aa

zog #~

13.92

1377

13.,~

1~.30

~.79

13.85

85%

hTA

H

Table 2. Formation constants of IDA and NTA complexes with Al(lll), Cr(lll) and Be(ll) stems

~emDer ature

a5°c

as°c

46°0

O,IM

0,2M

0,3M

0.424

O,IM

O,IM

log K 2 log K2 log P2

8.82 7.50 16.12

8.55 7,45 16.00

8.4,5 7.41 15.86

8.23 7.41 15.64

8.4S 7.22 16.64

8.28 7.21 15.40

log K1 log K9 log f12

10..53 8.55

19.08

10.46 8.89 18.85

10..31 8.18 18.49

I0.07 8.01 18.08

10.30 8.27 18.57

10.18 8,14 18.32

log K I log ~

8.88 6.82

8.17 6.25

8.08 6.24

7.61 6..17

8.71 6.67

8,56 6.68

Zog/~2

15.70

14.43

14.83

9.78

15.88

15.m

zog r 1

9.74

9.44

9.~

9.01

9.51

9.sa

log ~

8.34

8.19

8.16

7.88

8.04

8.01

log tip

18,11

17.~3

17.43

16.89

17.5.5

17.33

log ~

7.70

7.62

7.33

7.1S

7.59

7.43

log K1

8.44

8.18

8.04

7.82

8.16

7.94

Cr (III)-IDA

Cr(Zn)-hT~

Be(II)-IEg

The values of t/A were evaluated at various pH values from acid and ligand titration curves using the formula of Irving and Rossotti[6]. The values of rIA were plotted against the pH. From this graph, the values of proton-ligand stability constants, log K~ and log K~ were computed using different computation techniques[7]. The values are given in Table 1. In case of NTA, the values of log K3" were also determined.

The stepwise stability constants of the metal complexes were determined with the help of formation curves (Ji versus pL, Figs. 1-6) using various computational methods[7]. The values of Ji, the average number of ligands, attached per metal ion, and the free-ligand exponent pL were calculated from ligand and complex titration curves using the formula of Irving and Rossotti[6]. The values are given in Table 2.

409

Notes Table 3. Thermodynamic stability constants of AI(III), Cr(lll) and Be(II) chelates with IDA and NTA keSals use(: log

Kn D=O

L'L~i'*~Sl_USed

I]~,~ =i(zI:)

r,':'r(iIi-)

~ A

Io~3

I<1 p,=C

8.84

10,66

1~

~ ~=o

7.82

8.73

log //~2 ~=0

16.35

19.49

log I<1 D=O

9.30

9.96

I "ll

I

u

~

I

lc~ ~ / u ~

7.14

8.~

io~ /~2 ~=0

16.SI

18.47

I I

I

!

z Be{ :!)

ice K1 ~=0

7.89

I

8.62

I I

I

i%t j

o

o>01

,

,o

o

1~11o Thermodynamic stability constants, given in Table 3, were obtained by extrapolating the experimentally determined formation constants to zero ionic strength. The values of the overall changes in free-energy (AG°) enthalpy (AH °) and entropy (AS°) for the complexes, given in Table 4, were calculated using the temperature coetficient and Gibbs-Helmholtz equation [8].

L I I I

I

Acknowledgements--Thanks of the authors are due to the University Grants Commission, New Delhi, for providing a Junior Research Fellowship to one of them (A.S.) and also to Prof. H. L. Bhatnagar, Head of the Chemistry Department for providing laboratory facilities for the work. Department of Chemistry Kurukshetra University Kurukshetra-132119 India

S.N. DUBEY* A. SINGH D.M. PURl

o

Y~ I

"S

i

I%

I I I 0"~ 0

REFERENCES

I. S. Chaberek, Jr. and A. E. Martell, J. Am. Chem. Soc. 74, 5052 (1952). 2. P. L. Edelin De La Praudiere and L. A. K. Staveley, J. Inorg. Nucl. Chem. 26, 1713 (1964). ~. A. I. Vogel, Quantitative Inorganic Analysis, 3rd Edn., p. 516, 518, 522. Longman-Green, London (1961), 4. J. Bjerrum, Metal Ammine Formation in Aqueous Solution. Hasse, Copenhagen (1941). 5. M Calvin and K. W. Wilson, J. Am. Chem. Soc. 67, 2003 [ 1945I. 6. H. Irving and H. S. Rossotti, .k Chem. Soc. 2904 (1954). 7. H. Irving and H. S. Rossotti, J. Chem. Soc. 3397 (1953). 8. K. B. Yatsimirskii and V. P. Vasil'Ev', Instability Constants of Complex Compounds, p. 59-65. Pergamon Press, Oxford (1960),

*Author to whom correspondence should be addressed.

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