Spectrophotometric and potentiometric determination of the protonation constants of dithiocarbazates and studies on some of their metallic chelates

Talanta, Vol. 32, No. 1, pp. 73-77, 1985 Printed in Great Britain. A11 rights reserved

0039-9140/85 $3.00 + 0.00 Copyright 0 1985 Pergamon Press Ltd

SPECTROPHOTOMETRIC AND POTENTIOMETRIC DETERMINATION OF THE PROTONATION CONSTANTS OF DITHIOCARBAZATES AND STUDIES ON SOME OF THEIR METALLIC CHELATES A. IZQUIERDO Department of Analytical Chemistry, University of Barcelona, Barcelona, Spain J. GUASCH and F. X. RIUS Department of Analytical Chemistry, University of Barcelona, Tarragona, Spain (Received 22 February 1984. Revised 17 April 1984. Accepted 27 July 1984)

Summary-The protonation constants of 2_methyldithiocarbazate, 3,3-dimethyldithiocarbazate and 3-methyl-3_phenyldithiocarbazate, have been determined potentiometrically (I = l.OOM) and spectrophotometrically (I = O.OlM) at 25” in aqueous solution. The analytical properties of the reagents have been studied, and also the compositions of some of their insoluble metallic chelates.

have been widely studied as potential substitutes for dithiocarbamates, which have found considerable application as surfactants, pharmaceuticals and very active fungicides, and in vulcanization processes. This has given rise to a number of papers dealing with their synthesisI and the study of their structural properties.s-” In contrast, their analytical properties have received little attention. 3-Phenyldithiocarbazate was proposed by Musil and Haas12 for the gravimetric determination of Ag(I), Cu(I1) and Pb(I1) and by T’ien and Wang3 for the calorimetric determination of Cu(II). Hydrazinium dithiocarbazate has been used in the photometric determination of V(V), Fe(I1) and Fe(II1) by Byr’ko et ~1.‘~and of Fe(II), Fe(III) and Co(I1) by Haas et aj 15916

Dithiocarbazates

Because the substances containing the dithiocarbazate group are of interest as multidentate ligands, a systematic study of these compounds as analytical reagents has been undertaken in this laboratory. In this paper, the analytical properties of 2-methyldithiocarbazate (I), 3,3-dimethyldithiocarbazate (II) and 3-methyl-3-phenyldithiocarbazate (III) are described. -

EXPERIMENTAL Insrruments A Beckman 5260 spectrophotometer (l-cm cells) was used for spectrophotometric measurements. A Beckman 4260 infrared spectrophotometer (KBr discs) and a Perkin-Elmer R-24B NMR spectrophotometer were used. A Radiometer

PHM-64 pH-meter along with a G-202B Radiometer glass electrode, an Ag/AgCl reference electrode prepared according to Brown” and a Wilhelm-type salt bridge” were used to measure pH values. Reagents

Reagents (I) and (II) were obtained by the method of Anthoni et aL4 by addition of CS, to the corresponding hydraxines in alkaline medium. Reagent (III) was synthesized by a modification of the method of Henriksen and Jensen.) A solution of 0.04 mole of CS, in 5 ml of dimethyl sulphoxide was added dropwise during 1 hr, at room temperature, to a suspension of 0.04 mole of sodium hydride (with 20% paralhn) in 15 ml of dry dimethyl sulphoxide containing 0.04 mole of I-methyl-1-phenylhydrazine. The resulting solution was added to a mixture of 300 ml of acetone and diethyl ether (2:3) and allowed to stand for 24 hr at 4”. The white crystals obtained were filtered off by suction, washed with diethyl ether and recrystallized from chloroform (yield 2430%). Melting points: (I) 178-179”, (II)128-129” and (III) 121-123”. Elemental analysis: (I) H,NN(CH,) CSSK .2H,O requires 12.2% C, 4.6% H, 14.3% N, 32.7% S, 19.9% K; found 12.1% C, 4.3% H, 14.6% N, 31.8% S, 19.7% K. (II) (CH,),NNHCSSK. 2H,O requires 17.1% C, 5.3% H, 13.3% N, 30.5% S, 18.6% K; found 17.6% C, 4.9% H, 13.4% N, 29.9% S, 18.7% K. (III) (CH,) (C,H,) NNHCSSNa.2H,O requires 37.5% C, 5.1% H, 10.9% N, 25.0% S, 9.0% Na; found 37.7% C, 5.0% H, 10.6% N, 25.8% S, 8.8% Na.

Because of the low stability of the solid dithiocarbazates in air (3-14x decomposition per day) they must be stored under nitrogen at 4”. They are then stable for at least I yr. Their purity was tested by the method of Clarke er al.,19 originally devised for the analysis of dithiocarbamates. The substance is decomposed in strongly acidic medium and the 73

14

ANALYTICAL

DATA

of the dithiocarbazates in 1:4 mole ratio. The precipitates were filtered off, washed with water, ethanol and acetone and dried under reduced pressure over calcium chloride.

CS, evolved is collected in methanolic potassium hydroxide solution. Iodine titration of the xanthate thus formed showed, in all cases, the purity of the dithiocarbazates to be higher than 99%.

RESULTS AND DISCUSSION

Protonation constants The protonation

constants were determined at 25 + 0.1” and I = O.OlM by using the spectrophotometric method of Stenstrom and Goldsmith,M which makes use of the relationship between absorbance and pH; 1 ml of the aqueous lo-‘M dithiocarbazate solution was diluted to 50 ml with O.OlM buffer prepared as described by Perrin and Dempsey?’ The pH and the absorbance at 273 nm (I), 267 nm (II) and 285 nm (III) were measured. The potentiometric determination involved titration substances of the test alkaline solutions Tlf.3x lo-‘-35 x 10m3M) with O.lOOM nitric acid at Z = l.OOM, with potassium nitrate as ionic medium, at 25 + 0.1”. Before and during the titration, a stream of pure : nitrogen was continuously bubbled through the solutions. The potential of the following cell was measured: Test solution: XM reagent yM KOH (1 --x -y)MKNO,

Glass electrode

(-)

ligands are given in Table 1. The IR absorption peaks are assigned in accordance with earlier works.‘*3.4,69 These bands confhm the structure of the dithiocarbazates; thus, the addition of CS2 to methylhydrazine takes place through the substituted nitrogen atom, giving rise to the appearance of the -NH2 bending peak (1580-1645 cm-‘) in the spectrum of (I). This band does not appear in the spectra of (II) and (III) since the CS, is added to the primary

I I I l.OOMKNO, ;

Beforehand, O.lOM nitric acid was titrated with O.lOM potassium hydroxide in the same cell. The Gran method” was used to determine E”’ and Ej which relate the measured potential E to the hydrogen-ion concentration [H+] by means of the equation: E(mV)=E”‘-59.157

The solubility values and the most significant infrared and NMR spectrophotometric data for the

log[H+]+Ej

A known amount of reagent was added to the alkaline solution, then the mixture was titrated with nitric acid. Chelation studies The pH range over which complexation occurred, as indicated by colour and colour formation, was determined for each combination of metal ion and test substance by the technique established by Benedetti-Pichler.23 The solubilities of the complexes in benzene, diethyl ether, chloroform, isoamyl alcohol and methyl isobutyl ketone were tested and absorption data (A,, , 6) for buffered solutions of complexes were measured. The insoluble metal chelates were prepared by reacting aqueous solutions of the metal ions with aqueous solutions

I I ;

“‘lM KC1 A&l,,) 0.99M KNO,

I and of 1,l -dimethylhydrazine atom nitrogen 1-methyl-1-phenylhydrazine. The NMR spectra recorded for D20 solutions of the reagents with DSS as internal reference, give peaks similar to those reported previously.‘“*l’ The large shift corresponding to the methyl group of (II) (6 = 2.55 ppm) relative to (III) (S = 3.66 ppm) is consistent with the presence of the neighbouring -CSS- group in the 2-methyldithiocarbazate. The ultraviolet spectra of the reagents in some common solvents were measured (Table 2). A bathochromic shift of J.,,,a,is observed for (I) and (II) as the polarity of the solvent decreases. This can be attributed to n-n* transitions. The spectrum of (III) is more complicated, with the presence of overlapping bands corresponding to K--IL* aromatic ring transitions.

Table 1. Solubilities and infrared and NMR absorption oeaks of the lieands Ligand

(I) Solubility in water, g/l. Solubility in ethanol, g/l. Solubility in acetone, g/l. Solubility in diethyl ether, g/l. Solubility in benzene, g/l. Solubility in hexane, g/l. N-H stretch, cm-’ NH, bending, cm -I N-C stretch, cm-’ CH, bending, cm-’ H,C-N stretch, cm-’ N-N stretch, cm-i CSS- stretch asymm, cm-’ CSS- stretch symm, cm-’ NZ-CH3 protons (NMR), ppm N3-CH3 protons (NMR), ppm Aromatic protons (NMR), ppm

Ag(+)

(II)

(III,

200 200 200 17.1 6.5 8.3 3.5 2.0 1.8 0.7 0.3 0.1 0.1 0.1 0.1 0.5 0.1 0.1 3222-2928m 3210-285Om 3240-3 1OOm 1635m 1494s 1458m 1463m 1434m 1438m 1443m 1310s 1268m 1306s 1012m 1005s 1000s 976s 974s 965s 694m 67Om 3.61 2.55 3.66 7.30

ANALYTICAL

Table 2. Ultraviolet

data

DATA

75

for the ligands

in common

(I)

glycol

85.1

alcohol

83.6

Ethylene Methyl

Ethyl alcohol

79.6

Isopropyl

76.3

alcohol

1.6 8.7 1.7 9.0 1.7 9.2

x x x x x x

lo4 10’ lo4 10’ 104 10)

The stability of the reagents in aqueous solutions is pH-dependent. In neutral or alkaline media, the absorbance at A,,,,,,decreases by about 15% in 24 hr, the sequence of stability being (II) > (III) > (I). In fairly acidic medium (pH < 2.5) all decompose rapidly with evolution of CS,. This breakdown hinders the experimental determination of the absorbance of the protonated species in acidic media, as is required for the calculation of the equilibrium constants by the Stenstriim and Goldsmith method. Therefore, the Sommer method” was used, since it allows graphical determination of the values by solution of the equation: (AR-A)

A,-A=AA,-A,,-----

W+l KY

where A, and A,, are the absorbances of the fully deprotonated and protonated reagents respectively, and KY is the protonation constant. The resulting data have been optimized by the graphical method described by Rossotti and Rossotti26 and McBryde.*’

RNH-C

p \

e

K1

RNH-CH

SH

SK2, \

TAL. 32,,--F

css-

nm

lo4 10) lo4 103 lo4 103 lo4 lo3 lo4 103

I.mole-‘.cm-’ 9.4 9.4 1.2 1.3 1.1 1.0 1.2 1.1 8.9 1.0

285 255 277 263 290 253 292 253 289 259

103 10’ 104 lo4 104 lo4 lo4 104 10’ 104

which can be calculated experimentally from the pH of the solution, the total concentration of titratable hydrogen ion HT and the total reagent concentration RT. The constants were determined by use of the least-squares program MINIPOT.3o The protonation constants are given in Table 3. The discrepancy between the spectrophotometric and potentiometric values is attributed to the difference in ionic strength. The results obtained in the present work are not in accord with the overall ionization mechanism proposed for the hydrazinium dithiocarbazate by Byr’ko et a1.,14 who suggested

2S_

+

R-N

_-C-S+

pK,

= 3.87

+ 0.10

pK2

= 8.90

+ 0.10

S2-

The fact that, in the present study, both potentiometrically and spectrophotometrically only a single protonation constant was observed, along with the absence of ionizable cations such the hydrazinium ion, provides strong evidence for a single acid-base equilibrium. This equilibrium, in a similar way to that proposed by Hulanick?” for the dithiocarbamates, could be represented by

R” R’\N-N/R’ R

x x x x x x x x x x

G’[H+l fi= & - W+l + &IW+l = 1 + K:[H+] &

\

\ R

x x x x x x x x x x

1.3 9.6 1.4 7.4 1.4 9.0 1.5 8.7 1.5 8.6

EmaX 9

AM,,,

sorbance of the I-methyl-1-phenylhydrazine released on decomposition of (III) in acid. However, this constant, along with those for (I) and (II), was determined potentiometrically. The E”’ and Ej values found in the Gran titration allowed determination of the protonation constant by use of the Bjerrum function,29 n,

RN=C\

S_-

The values of the protonation constants determined fully coincide with those obtained by means of the least-squares program MINISPEF28 (PDP 11/60 computer). Table 3 lists the values of the protonation constants of (I) and (II). Log KY for (III) could not be determined because of interference from the ab-

,H’ R’>N_N/R’

285 252 289.5 252 290 252 291 253 291.5 253.5

1.7 x 104 1.0 x 104 1.7 x lo4 -

273 228 279.5 279.5 231 281 231 282 231

94.6

Water

(III)

Gnax7 I.mole-‘.cm-’

Lx, nm

%8X. I.mole-‘.cm-’

Lx* nrn

Z values%

solvents

(II)

(,,-

R’

+ CS ‘H

2

76

ANALYTICAL

Table 3. Protonation constants of the reagents at 25 + O.l”C Potentiometrically (Z = 1.OOM)

Reagent

Spectrophotometrically (I= O.OlM) log KY

(1) (II) (III)

4.90 * 0.07 4.56 k 0.02 -

5.42 f 0.05 4.73 + 0.02 6.18 + 0.03

log KY

This mechanism is consistent with the reported zwitterion structure of dithiocarbazic acid’,’ and the indirect titration of the dithiocarbazates with iodine, where the presence of the CS2 evolved in strongly acidic medium is required. On the other hand, the differences between constants for the different compounds show the influence of the substituents bound to the protonated N-N skeleton (the protonation of the CSS- group would give rise to very similar values of the protonation constants, as occurs with xanthates and carboxydithioateP2). Table 4 gives the spectral characteristics of the soluble metal-reagent complexes, which have high absorptivities. The pH does not greatly influence the complex formation; this confirms that it is the N-N group rather than the CSS- group that is protonated.

DATA

All three substances react very similarly with metal ions, although it is interesting to note the closer resemblance between (II) and (III). This behaviour can be attributed to the difference in structure of the complexes (Table 5). Reagents (II) and (III) give MS4 type complexes, as previously reported by Battistoni et al.’ Chelation occurs through both sulphur atoms, according to the structure: I

R\ /N-N R

s( /I’

INSM\,

sL,/

R )N-N

R” MS4 type

<

R’

The infrared spectra of the complexes studied show that the C-N stretching absorption (1525-1490 cm-‘) has been shifted to higher frequencies than for the

Table 4. Spectrophotometric characteristics of the complexes 1, FlM t,I.mole-‘.cm-’ Optimum pH Ligand Metal ion Tl(II1) Cu(II) Fe(H) Ni(I1) Bi(II1) Tl(II1) Cu(I1) Fe(U) Fe(II1) Ni(II) Cu(II) Fe(I1) Fe(II1) Ni(I1) Co(H)

320 437 317 315 358 332 600 322 338 329 455 360 351 337 375

2-7 4-9 4-9 5-8 7-10 2-7 9-12 68 4-8 6-8 l-9 2-6 5-10 6-9 5-10

2.4 x 7.1 x 7.6 x 2.7 x 1.2 x 2.3 x 1.2 x 2.5 x 6.6 x 2.2 x 1.2 x 1.5 x 7.6 x 2.7 x 1.4 x

lo4 10’ lo3 lo4 104 lo4 104 lo4 lo3 104 104 104 lo3 104 104

Table 5. Infrared data for the dithiocarbazato metal chelates

Compound*

W-W(CiW,I

Decomp. point, Colour “C yellow 125 white 140 brown 132 white 154 red 185 orange 93 white 164 yellow 128 green 119 white 163 green 190 vellow 154 irown 235 green 245

Pb[NH,N(CHXS,I, Cu[NH,NCH,)C%J, CdN-W(CWCS,I, NW%WH3)CS212 Bi[NH2N(CI-WS213 &WW,NNHCS,I WCW,NNHCS,I, W(CW,NNHCS,I, CWJ-I,),NNHCS,I, NWI-U~HW, BiKCH,),NNHCS,l, C;i(C,Ii;)(CH,)Ni;jHCS,], Ni[(C,H,)(CH,)NNHCS,l~ *These formulae were confirmed by elemental analysis. toverlapping band. §Shoulders in brackets.

6(NH?)

1605 1600 1597 1583 1590 1585 -

"(C-N) t

1490 1495 1490 1512 1525 1502 1495 1492

“&s- asym) 955(945) 950(940) 970 960(940) 1010(1005) 950(930) 980 977 975 970 970 970 970 970

577 580 567 573 564 576 680 675 690 685 690 690 695 690 7

2,l

I,7

ANALYTICAL

alkali-metal dithiocarbazates [1494 cn-’ (II) and 1498 cm-’ (III)] because of the important double bond C=N contribution in this type of chelate.‘v* The antisymmetric and symmetric CSS stretching modes, absorbing at around 975 cm-’ and 685 cm-’ respectively, also confirm this kind of chelation as reported earlier for MS, chelates of nickel with 2-methyldithiocarbazate,‘,’ 3,3-dimethyldithiocarbazate’.’ and 3-methyl-3-phenyldithiocarbazate.’ In contrast reagent (I) forms MN,S2 type metal complexes, through the nitrogen and sulphur atoms:

MN,S,

type

The C-N stretching vibration, with a smaller double bond contribution, overlaps in this case with neighbouring lower-frequency bands such as C-H (around 1440 cm-‘).* The chelates formed with reagent (I) all display NH2 bonding peaks at lower frequencies (1605-1583 cm-‘), supporting the postulate of MN*& co-ordination. ‘*’ Finally, the presence of a shoulder in the absorption peak assigned to the asymmetric CSS- mode also confirms’~* the chelation through only one nitrogen and one sulphur atom.

REFERENCES 1. U. Anthoni, Acra Chem. Stand., 1966, 20, 2742, and references therein. 2. L. Cambi, E. D. Paglia, G. Bargigia and G. Crescentini, Chim. Ind., 1966, 48, 689. 3. L. Henriksen and 0. R. Jensen, Acta Chem. &and., 1968, 22, 3042.

DATA

71

4. U. Anthoni, B. M. Dahl, Ch. Larsen and P. H. Nielsen, ibid., 1969, 23, 1061. 5. A. Braibanti, L. Manotti, A. Tiripicchio and F. Lo Giudice, Acta Cryst. Sec. B, 1969, 25, 93. 6. C. Battistoni, A. Monaci, G. Mattogno and F. Tarli, Inorg. Nucl. Chem. Lett., 1971, 7, 1081. I. C. Battistoni, G. Mattogno, A. Monaci and F. Tarli, J. Inorg. Nucl. Chem., 1971, 33, 3815. 8. M. Akbar Ali. S. E. Livinastone and D. J. Phillins. Inorg. Chim. kcta, 1971, 5,; 19. 9. M. F. Iskander and L. El-Sayed, J. Inorg. Nuci. Chem., 1971, 33.4253. 10. D. Gattegno and A. M. Giuliani, Tetrahedron, 1974,30, 701. 11. Idem, J. Inorg. Nucl. Chem., 1974, 36, 1553. 12. A. Musil and W. Haas, Mikrochim. Acta, 1958, 156. 13. Ping-Shih T’ien and K’uei Wang, Sci. Sinica (Peking), 1957. 5. 651: Chem. Abstr.. 1957. 51. 17567. 14. V. M. ‘By&o, A. I. Busev, T. ‘I. Tikhonova, N. V. Baibakova and L. I. Shepel, Zh. Analit. Khim., 1975,30, 1885. 15. A. Musil and W. Haas, Mikrochim. Acta, 1959, 31. 16. W. Haas, ibid., 1963, 274. 17. A. S. Brown, J. Am. Chem. Sot., 1934, 56, 646. 18. W. Forsling, S. Hietanen and L. G. Sillen, Acta Chem. Stand., 1952, 6, 601. 19. D. G. Clarke, H. Baum, E. L. Stanley and W. F. Hester, Anal. Chem., 1951, 23, 1842. 20. W. Stenstrom and N. Goldsmith, J. Phys. Chem., 1926, 30, 1683. 21. D. D. Perrin and B. Dempsey, Buffers for pH and Metal Ion Control, 4th Ed., Table 3.8 (pp. 447). Longmans, London, 1978. 22. G. Gran, Analyst, 1952, 77, 661. 23. A. A. Benedetti-Pichler, Microtechnique of Inorganic Analysis, Wiley, New York, 1950. _ _ 24. E. M. Kosower. J. Am. Chem. Sot. 1958.80.3253: ,, , 1958. 80, 3261; 1958,‘80, 3267. 25. L. Sommer, Folia Fat. Scienc. Nat. Univ. Purkynianae Brno, 1964, 5, 1. 26. F. J. C. Rc .cotti and H. Rossotti, The Determination of Stability Co. ;tants, McGraw-Hill, New York, 1961. 27. W. A. E. McBryde, Talanta, 1974, 21, 919. 28. F. Gaizer and A. PuskLs, ibid., 1981, 28, 925. 29. J. Bjerrum, Metal Amine Formation in Aqueous Solution, 2nd Ed., Haase, Copenhagen, 1957. 30. F. Gaizer and A. Puskas, Talanta, 1981, 28, 565. 31. A. Hulanicki, ibid., 1967, 14, 1371. 32. L. G. Sill&n and A. E. Martell, Stability Constants of Metal-Ion Complexes, 2nd Ed., Special Publication No. 17, The Chemical Society, London, 1964.