Ionization equilibria of copper(II) chloride in dimethyl sulphoxide

Electrochimico Arta. Vol. 25, pp. 1033 Pergamon Press Ltd. 1%“. Rimed m

Briton

IONIZATION EQUILIBRIA OF COPPER CHLORIDE IN DIMETHYL SULPHOXIDE W. LIBUS, M. PILARCZKK and T. SZUCHNICKA Department of Physical Chemistry, Institute of Chemical Engineering, Technical University of Gdatisk, 80-952 Gdarisk, Poland (Received 20 April 1979; in revisedform 23 November

1979)

Visible absorption spectra of the CuCl+ and CuCl; complexes in dimethyl sulphoxide (DMSO) are derived from an auxiliary study of the Cu(CIO,), + Et,NCI mixed solutions of variable chloride concentration. The second step association equilibrium CuCI+ + Cl- = CuCl; is then shown, from the Abstract

visible absorption spectra, to predominate over a broad concentration range in the solutions of CuC1, in DMSO. Accordingly, the second step association constant is derived from the molar conductivity curve of CuCI, using the Shedlovsky’s method, while the first step association constant is estimated from the conductometric data of the most dilute solutions: KY = (16.5 k 1.5)10*, KY = (8 + 0.2)10’ at 25°C. The resultsarecompared with those earlier obtained for NiCl, and CoCI, in DMSO solution, non-conformation of the association constants to the Irving-Williams series being pointed out.

INTRODUCTION

Solutions of transition metal salts in donor solvents (including water) represent a broad class of electrolytes, whose most characteristic feature consists in the formation of well-defined solution complexes in which the solvent molecules and/or the anions play the role of ligands coordinated to the metal cations. A number of different coordination complexes are usually possible in the given system, their relative contents varying with varying concentration of the solution. As a result, the chemical equilibrium approach to the interpretation of properties of the respective solutions becomes of larger importance for this class of electrolytes than for the more “typical” ones, such as the alkali or tetraalkylammonium salts, usually considered from the stand point of the physical theories. Adopting the first approach we assume that increasing concentration of the solution brings about reversible changes in the cootdination state of the solution, separable in the effects they produce from changing interactions between ions, whether complex or simple. A number of our earlier observations have shown that coordination complexes of a given type once formed, usually exhibit properties to a good approximation independent of the nature of the central metal ion[lG3]. This introduces a certain simplification to the formidable task of interpreting solution properties of transition metal salts in terms of coordination equilibria. However, there remains the important question of the main factors affecting the relative stabilities of the different types of coordination complexes appearing in solution, viz. those determining their dependence on the nature of all : the metal cation, the anion, and the solvent. The present series of papers aims at obtaining partial answers to this general question for salt solutions of the divalent transition metals belonging to the Mn to Zn series in the now popular solvents like acetonitrile or dimethyl sulphoxide. The main task consists in determining the nature and relative stabilities of complexes formed in closely related systems. Our recent papers were con-

cerned with solutions of NiCl, and CoCl, in dimethylsulphoxide[4,5]. Presently, we report the results obtained for solutions of CuCl, in this solvent. Copper(U) chloride reacts with DMSO to form the CuCl, * 2DMS0 solid compound, whose structure is known from X-ray diffraction analysis[6]. Consecutive formation of the CuCl+, CuCli, CuCl;, and CuCl:complexes in DMSO solution of variable chloride concentration, inferred from a spectrophotometric and potentiometric study, was reported by Suarez et ar.[7]. On the other hand, little is known as regards the nature of species present in solutions of C&l, in dimethyl sulphoxide[8]. In the present study we use conductometry and spectrophotometry as the source of information. An auxiliary study of mixed solutions of variable chloride anion concentration provides information on the spectra of the single chlorocomplexes of copper in DMSO.

EXPERIMENTAL

Reagent grade dimethyl sulphoxide was dried over CaO and distilled under reduced pressure, Specific conductivity of the final material was 1.2 x lo-’ S cm-‘. The DMSO solvated copper(U) perchlorate, CU(C~O~)~ .6DMSO, was obtained by dissolving the hydrated salt in the anhydrous solvent, followed by evaporating the solution under reduced pressure at room temperature. The blue crystals which deposited from the solution were recrystallized twice from the pure solvent and stored in a desiccator. Crystals of CuCl, *2DMS0 were obtained in an analdgous way using anhydrous CuCl, as the starting material. Reagent grade tetraethylammonium chloride was purified by repeated crystallizations from acetonitrile, and dried in vacua at 60°C. The product was used to prepare the stock solution in DI&O by weighing. Stock solutions of Cu(ClO,), and of CuCl, in DMSO were prepared from the DMSO solvated salts and analyzed for copper by standard EDTA titration

1034

W.

LIBUS,

M.

PILARCZYK AND

using murexide as indicator, gravimetrically as CuNCS, and electrogravimetrically. Several determinations were made by each method. The results obtained by the different methods were consistent to within 0.1 per cent. Conductivities were determined as described in the preceding pape+b] using a Luder type bridge and ac of lo3 Hz. at (25 + O.Ol)“C. Absorntion snectra were measured’ b; m&ms of a Zeiss VSU-2 spectrophotometer equipped with a thermostated cell compartment. Preparations of the solutions and further manipulations were carried out in a dry box. 1

RESULTS AND Auxiliary

DISCUSSION

study of the Cu(ClO&

+ Et,NCl

mixed

solutions Figure 1 shows the visible absorption spectra of a series of solutions containing Cu(ClO& at a constant concentration and tetraethylammonium chloride at a number of different concentrations, up to 0.174 M at 25°C. The concentration independent band at 860 nm of the pure Cu(ClO,), solution in DMSO (curve 1) is known to be due to the 2 T,,(D) 6 ‘E,(D) transition of the Cu(DMSO)i+ tetragonally distorted solvocomptex[9]. Increasing concentration of Et,NCI added brings about a sharp increase in intensity of the

T.

SZUCHNICKA

band and its concomitant small red shift, indicating coordination of the chloride anion in octahedral configuration. Closer examination of the set of curves shows that beginning from the Cl : Cu mole ratio of 1 absorption remains constant within the 700-800 nm spectral range over a certain concentration range of EthNCl, as illustrated in Fig. 2. It follows, that it is formation of the CuCl+ complex, apart from its real nature, that is responsible for the spectral changes observed at lower concentrations of l$,NCl. It also follows that the complex is rather unexpectedly stable, as compared with the analogous CoCl+ and NiCl+ complexes in DMSO solution. Obviously, the abovementioned constancy of absorption within a certain spectral range should be considered as resulting from the existence of an isosbestic point corresponding to the intersection of the respective absorption curves of the single solution complexes at a very small angle. A large scale plot indicated that the true isosbestic point is located somewhere at 780 nm. It is most probable that the isosbestic point, in Fig. 1 indicated as P,,, corresponds to the equilibrium cooexistence of the CuCl f and CuCl: consecutive complexes. Appearance of the next higher complex, presumedly CuCl;, in the set of absorption curves now under consideration is manifested by the fact that the respective curves cease passing through P,,. Appearance of the next isosbestic

1201

I

110 loo-

80.

DUJ

I””

800

900

1000

1100

1200

1300

Koo

l5m

h/nm Fig. 1. Absorptionspectra of mixed solutionsCu(CIO,), (constant concentrationof 0.0044 M) + EtdNCI (variable concentration) in DMSO at 25°C. Concentrations of Et,NCI: 1 - 0.0 M; 2 - 0.00175 M; 3 0.00245M;4-0.0034SM;5-0.00509M;6-0.00691M;7-0.00756M;8-0.01143M;9-0.01508M;100.02166M;11-0.02913M;12-0.04398M;13-0.0566lM;14-0.08520M;15-0.1738M.Indicatedare the derived spectra of the CuCl+ (curve 16) and CuCIi (curve 17) complexes.

1035

Ionization equilibria of copper(H) chloride Range I c

C

‘V

(1)

820nm

i>

at A,,, and

800nm ,.

(2)

780nm

at any A. Range II (3) E cz

0

t

,

Ca-/ht2+

2

-

El

(4)

c E23 -

El

at A,,, and C.? -

CIEl

&2 -

Fig. 2. Dependence of the mean molar absorption coefficient of copper(H) on the Cl-Cu mole ratio at selected constant wavelengths.

(5)

C2 at

any A.

Range III c,

E 23

=c-

Fig. 3. Ranges of appearance of isosbestic points in the set of absorption curves of the Cu(CIO,), + Et4NCI (variable concentration) mixed solutions.

point, located at 1040 nm, closely follows the disappearance. of the former one with increasing concentration of the chloride anion. This we take as an indication that a new two-species equilibrium is now established, viz. that involving the CuCli and CuCl; complexes. The above inferences are summarized in Fig. 3 showing, in a schematic way, ranges of occurrence of the single chloro-complexes of copper(I1) in the Cu(ClO& + Et,NCl mixed solutions in DMSO. Indicated are also the ranges of Et,NCl concentrations in which the abovementioned isosbestic points Pi2 and P,, are observed. It may be readily shown from Beer’s law and respective material balances that the following relations should be valid in the three ranges indicated in Fig. 3:

-

E (6)

E23

-

El

co, cl, and c2 denote at A,,. In these relations equilibrium concentrations of the CL?+, CuCl+, and @Cl; solution species, respectively (irrespective of their solvation), c denotes total concentration of copper( co, sl, and e2 are the molar absorption coefficients of the respective species, &i2 and sZ3 are the molar absorption coefficients of copper(I1) at the respective isosbestic points, and E is the measured mean molar absorption coefficient of copper(H) in the given solution. Equilibrium concentrations of the single chloro-complexes of copper(I1) calculated from the spectrophotometric data for the Cu(ClO,), + Et4NCl mixed solutions in Fig. 4 are plotted us total chloride anion concentration. On the other hand, plots of e,(L) and e2(A) calculated from the same data are indicated in Figs 1 and 5 together with the spectra of the respective solutions.

Cl02

a01 c Et, NCl/M

Fig. 4. Mole per cent contents of single chloro-complexes of copper(l1) in the Cu(ClO.Ja (0.0044 M) + Et,NCI mixed solutions in DMSO as function of the concentration of Et,NCI;

at 25°C.

W. LIBUS,M. PILARCZYKANDT. SZUCHNICKA

1036

14 I . 6GU

700

800

900

loo0

llcu

1200

1300

l4Oc

?+-n Fig. 5. Absorfition spectra of the CuCl, solutions in DMSO at 25°C. Concentrations: 2 - 0.00043 M; 3 0.0079M;4-0.00116M;5~0.00227M;6-0.00304M;7-0.00711M;IndicatedarespectraoftheC~~~ (curve l), CuCl+ (curve 16), and CuClp (curve 17) single species.

Solutions of CuCl, in DMSO Figure 5 shows the visible absorption spectra of CuCl, in DMSO solution at a number of different concentrations, while plots of the mean molar absorp tion coefficient us concentration at seveial selected wavelengths are shown in Fig. 6. As is seen, all the spectra are intermediate between those of the CuCl+ and CuClz complexes derived from the study on the mixed solutions, as described in the preceding Section.

. 95ora-n -v 0

0

05

ti

lo l@C/p.q

1.5

Fig. 6. Concentration dependences of the mean molar absorption coemcient of copper(N) at selected constant wavelengths for CuC12 in DMSO solution at 25°C.

It follows, that the second-step cuc1+

association

+ Cl - = cuc1;

equilibrium (7)

predominates over a broad concentration range in the CuCl, in DMSO solutions now under consideration. This is cjuite an exceptional situation, not encountered in the other divalent transition metal chloride solutions in DMSO. It provides favourable conditions for the derivation from the conductometric data of the respective equilibrium constant and, what is particularly interesting, of the equivalent conductivity of the intermediate cation. The molar conductivity curve of CuCI, in DMSO solution at 25°C is shown in Fig. 7, while Table 1 lists numerical values of the conductivity. We may note that the limiting molar conductivity of the Cu(DMSO), ‘+ .2Cl- complex electrolyte, expected as the only solution form of CuCI, in DMSO at infinite dilution, would have the value of 88.8 S cm’ mol- I, as calculated from the known ionic conductivities[lO, 111. It becomes clear that extensive association must occur already at the lowest concentrations accessible experimentally, although it is not a priori evident from the conductometric data alone that only the second step association occurs within a certain concentration range of CuCl,. Assuming this being thecase, as inferred from the spectrophotometric results, the data were analyzed using the Shedlovsky’s method[12,13]. We found it impossible to use the more advanced methods of conductivity evaluation, as

Ionization equilibria of copper

I8

2 "E

CL?+ + cl-

321.. I b

while the

=cuc1+

(8)

overlaps with the presently considered one at concentrations below some 10e4 M. The final Shedlowsky’s plot for CuCl, in DMSO solution corresponding to the assumed equilibrium (7) is shown in Fig. 8. Values of the limiting conductivity r\FI of the intermediate electrolyte CuCl+ Cl- and its association constant K:, are found from the linear least squares approximation in accordance with the relation

inu \

r

1037

they apply to the very lowest concentrations, first step association equilibrium

40

T

chloride

4

MO5

CIA

where S(Z) is a tabulated

function[14]

defined as

while

01

02

03

‘z (C/M?

S denoting the limiting Onsager slope (92.572 for DMSO at 25”(Z), and the other symbols having their usual meanings. The limiting Debye-Hiickel equation was assumed as a sufficiently good approximation for the activity coefficient y, and the normal iterative procedure was used in deriving the best values of the two parameters :

K; = (8.00 + 0.2) x lo’, Fig. 7. Concentration dependence of the molar conductivity of C&I, in DMSO solutions at 25°C.

and hyj = 40.1 S cm2 mol- l, at 25°C. A check on the above derived association constant is provided by the diagram shown in Fig. 9, in which the mole per cent contents of the single solution complexes are plotted us totai concentration of CuCl,. At concentrations exceeding 0.0005 M the equilibrium concentrations were calculated from the above value of K:

Table 1. Molar conductivitics of copper(H) chloride in dimethyl sulphoxide solution at 25°C 103C

mol dn-’ 0.00707 0.00957 0.01132 0.01515 0.01852 0.05046 0.1195 0.1569 0.1966 0.2925 0.3944 0.5799 0.7761 0.9675 1.1645 1.7318 2.4733 3.2170

Am S cm2 mol-1 62.95 56.95 53.70 51.19 49.25 41.72 39.34 31.19 35.78 34.00 31.80 29.49 27.80 26.49 25.24 22.70 20.35 18.60

WC mol dm-3 4.687 6.463 9.757 12.934 16.873 25.593 32.202 35.460 38.591 45.176 51.518 57.360 63.711 70.620 76.974 83.697 96.862

4n S cm2mol-’ 16.46 14.55 12.58 11.42 10.50 9.32 8.75 8.45 8.40 8.11 7.86 7.61 7.47 7.32 7.16 7.03 6.80

W. LIB&, M. PILARCZYKANDT. SZUCHNICKA

1038

solution properties of transition metal chlorides important is determination of the first step association constant as well. Within the limits of the present study an estimation of the latter constant was possible from the molar conductivities measured at concentrations below 2 . low5 M. Corresponding part of the conductivity curve is shown in the upper part of Fig. 7. It is clear, that molar conductivities measured at as high dilutions cannot be of high accuracy, because of the solvent conductivity becoming comparable with that of the solution. As a result, but the simple Davies method[15] was used in their evaluation. Applied to the present case it was underlined by the assumption that

(12)

Fig. 8. Determination of the association constant of reaction 7 by the Shedlovsky’s method; final plot.

100 ,r ..

j Lcu*+

Fig. 9. Mole Per cent contents of single chloro-complexes of copper as function of CuCl, concentration in DMSO solution at 25°C ; broken lines calculated from the formation constants derived from conductometric data.

where A,, and A,, are molar conductivities of the CuCl+ .Cland Cu2+ . 2Cl- electrolytes, respectively, c1 and c0 denoting their respective concentrations. On the other hand, taking into account the material balances as well as the equilibrium expressions for reactions (7) and (8) the relation is obtained 2c = (1 - cK,)[Cl]

+ K,[CIIZ

+ KjK,[Cl]3

(13)

whereK, = Ki/Y, and K2 = Kf$Y2,K,,Kz,Y,,and Y, are equilibrium concentration quotients and activity coefficient quotients of reactions (7) and (8), respectively. Assuming the limiting Onsager dependences for the molar conductivities, viz. A, 1 = 40.1 55.8J1 and Al2 = 44.1 - 92.6,/I, and the limiting Debye-Hiickel relations for the activity coefficients, viz. log Y, = 4.4441 and log Y, = 2.22JI, values of KY and Ki were found for which the standard deviation of the calculated conductivities from the experimental ones was at minimum. An Odra 1204 computer was used in performing the calculations with the result that KY = (1.65 f 0.15) * lOsand Ki = (8f 0.5)102, at 25°C. The latter value agrees well with that derived from the conductometric data at higher concentrations, as described above. CONCLUSIONS

using the Debye-Hiickel equation involving the ionsize parameter % for the activity coefficient. The value of 8.2 A was estimated for B from the crystallographic radius of the Cl- anion and the molecular model of the CuCl+ cation, on the assumption that the real nature of the latter corresponds to the formula CuCl(DMSO)z, as discussed in the next Section. Resulting curve in Fig. 9 is indicated by the broken line. On the other hand, the full points were calculated directly from the spectrophotometric data using eg (4) of the preceding Section. AS is seen, the agreement between the spectrophotometric and conductometric results is reasonable at moderate concentrations and practically complete at the lowest concentrations. In view of the above results, there seems little doubt that the second step association equilibrium (7) in fact predominates in the diluted and moderately concentrated solutions of CuCIZ in DMSO. However, for the general question of the regularities governing

Combined with those of our preceding works[5] the present results provide a partial pattern of the variation in stability within the first transitional series for the monochloroand dichloro-complexes in DMSO solution. Respective plots are shown in Fig. 10. Numerical values of the equilibrium constants are collected in Table 2, along with those reported by other authors[7,16-181. The latter values roughly agree with ours with the exception of the data for CuCl+. It should be noted, however, that some of the reported formation constants relate to concentrated ionic media and not to the pure solutions in DMSO. We should also note that complete absence of the monochloroand dichloro-complexes of cobalt(U) in DMSO solution was claimed by Magnell and Reynolds[l9]. Apart from the marked uncertainties in the derived values of the formation constants, it becomes rather clear that the complexes under consideration do not conform to the Irving-Williams series, contrary to what is commonly believed to be the case[l7]. In this respect the chloro-complex formation in DMSO so-

Ionization

66 5

“2

OI 39 2l-

Cl’

co

Ni

n-2

I

cu

REFERENCES

Table 2. Formation constants of consecutive chlorocomplexes of cobalt(II), nickel(U), and copper(H) in DMSO solution at 25°C. Included are literature values of the constants log K;

log K;

log KY

Co(H)

2.78 f 0.09

2.48 &- 0.08

3.67 * 0.05

Ni(II)

2.47 + 0.06 2.7[16,17]

0.84 * 0.2

Cu(I1)

5.22 * 0.05 4.4[18,17] 4.5[19,17]

2.90 + 0.03 3.1[18,17] 3.0[19,17]

to that occurring in water[20], except MCI+ type complexes in water are conless stable than in DMSO. The latter has usually been ascribed to stronger solthe halide anions in protic solvents[17]. An “inverted Irving-Williams series” was also observed for the second step association constants of the divalent transition metal cations with the NO; anion in acetonitrile solution[21]. We believe that this is a general property of octahedral complex formation of the divalent transition metal cations with the relatively weakly coordinating anions when occurring in the strongly coordinating solvents like dimethyl sulphoxide or acetonitrile. It seems that the inverted Irving+Williams series may be rationalized in exactly the same way as was the normal one, while taking into account that it is the displaced solvent molecules that are the stronger ligands than the newcoming anionic ligands. Characteristically, the complexes under conis similar

that the siderably difference vation of

1039

chloride

Acknowledgement ~-The authors are indebted to Dr Roman Pastewski for writing the computer program and performing the calculations.

Fig. 10. Variations of the formation constants of the MCI+ and MCl$ type complexes in DMSO solution within the Co to Cu series of metals ; 25°C.

lution

of copper(I1)

sideration are usually formed in markedly endothermic reactions and owe their stability to the positive entropies of reaction[22]. The position of copper(H) in the inverted series, like in the normal one, may be due to increasing ligand field stabilization gained by the cupric complex with increasing tetragonal distortion when one or both of the axial solvent molecules of the parent hexasolvo-complex are subsituted by the weaker field anionic ligands.

l-l-1

J v

4.

equilibria

1. W. Lib&, Section Lectures of the X111-th. International Conference on Coordination Chemistry, CracowZakopane p. 241 (1970). 2. W. Libub and H. Strzelecki, Electrochim. Acta 15, 703 (1970); ibid. 16, 1749 (1971); ibid. 17, 577 (1972). 3. W. Lib&, M. Kluczkowski and W. Nierzwicki, J. Chem. Sot. Faraday Tmns. I 70, 1057 (1974); ibid. 72, 2552 (1976). Roczniki 4. W. LibuS, M. Pilarczyk and T. Szuchnicka, Chem. 49, 1291 (1975). 5. W. LibuS, M. Pilarczyk and T. Szuchnicka, Elecrrochim. Acta 20, 831 (1975); ibid. submitted. 6. R. D. Willett and Kun Chang, Inorg. Chim. Acta 4, 447 (1970). and J. Kleinberg, Inorg. 7. T. E. Suarez, R. T. Iwamoto Chem. Actn 7, 292 (1973). and L. Hiibner, Monatsh. 92, 1261 (1961). 8. V. Gutmann 9. H. Schlafer and H. P. Optiz, 2. Elektrochem. 65, 372 (1961). 10. W. LibuS and M. Pilarczyk, Bul. Acad. Polon. Sci.,Ser. sci. chim. 20, 539 (1972). 11. D. E. Arrington and E. Griswold, J. Phys. Chem. 74, 123 (1970). 12. R. M. Fuoss and T. Shedlovsky, J. Am. Chem. Sot. 71, 1497 (1949). and R. L. Kay, J. Phys. Chem. 60, 151 13. T. Shedlovsky (1956). 14. H. M. Dagget, 1. Am. Chem. Sot. 73,4977 (1951). p. 16. Butterworths, 15. C. W. Davies, Ion Association, London (1962). 16. F. Dickert and H. Hoffman, Ber. Bunsenges. Phys. Chem. 75, 1320 (1971). 17. S. Ahrland and N. Bjtirk, Coord. Chem. Rev. 16, 115 (1975). Bull. Sot. 18. A. Foil, M. Le DCmCzet and J. Courtot-Coupez, Chim. France 408 (1972). 19. K. R. Magnell and W. L. Reynolds, Inorg. Chim. Acta 6, 571 (1972). J. Solution Chem. 4, 1011 20. Z. LibuS and H. Tiatowska, (1975). L. Fraczyk and H. Strzelecki, 21. W. Libul, B. Chachulski, Roczniki Chem. 49, 19 (1975). 22. W. Lib&, B. Chachulski and L. Fr4czyk. PO!. J. Chem. 52, 493 (1978).