The aquation of the bis(oxalato)ethylenediaminechromate(III) ion

Notes chelate is more active than the parent compound against B.subtills, S.typhi, S.aureus, S.dysentricus, P.mirabilis and equally active against B.anthracis, C.ovis and E.coli. The cobalt complex has greater activity against B.subtilis, S.typhi, S.aureus, S.dysentricus, B.anthracis, P.mirabilis, C.ovis and E.coli. All the four compounds are found to be active against M.tubercdosis. Metal chelates in every ease have been found to be more potent than the parent compound which could be explained due to their greater lipid solubility and subsequent cellular penetration.

Acknowledgement--The authors wish to thank C.S.I.R. New Delhi for the award of a research fellowship to P.J. Department of Chemistry Holkar Science College lndore India

PRABUDDHA JAIN KAMAL K. CHATURVEDI

REFERENCES

1. P. Ray and A. K. Mukherjea, J. Ind. Chem. Soc. 32,604 (1955). 2. P. Jain and K. K. Chaturvedi, J. Ind. Chem. Soc. 52, 805 (1975). 3. P. Jain and K. K. Chaturvedi, J. Ind. Chem. Soc. 52, 1220,1225 (1975).

~3

4. P. Jain and K. K. Chaturvedi, J. lnorg. Nucl. Chem. 38, 799 (1976). 5. Takeo Tsukumato and Kennosuke Yuhi, Yakugakuzasshi 78, 706 (1958); Chem. Abstr. 52, 14561b (1958). 6. Takeo Tsukamato, Takeda Pharmaceutical Industries Ltd. Japan 5766 (1960); Chem. Abstr. 55, 5879f (1961). 7. P. Jain and K. K. Chaturvndi, Defence Science J. in press. 8. P. W. Selwood, Magnetochemistry, p. 92. Interscience, New York (1955). 9. A. Weissberger, Techniquesof Organic Chemistry, Vol. IX, p. 394. Interscience, New York (1956). 10. A. R. Katritsky, Quart. Rev. 13, 359 (1949). 11. P. Teysee and J. J. Charetee, Spectrochim. Acta 19, 1407 0963). 12. L. E. Clongherty, J. A. Sousa and G. M. Wyman, J. Org. Chem. 22, 462 (1957). 13. K. Ueno and A. E. Martell, J. Phys. Chem. 60, 1270 (1956). 14. L. J. Bellamy, The IR Spectra of Complex Molecules. Methuen, New York (1964). 15. P. Jain and K. K. Chaturvedi, J. Ind. Chem. Soc. 53, 360 (1976). 16. K. K. Chaturvndi, S. Siddiquei, B. K. Agrawal and R. Kaushal, Ind. J. Pharm. 37, 85 (1975).

.l. inorg, nucl. Chem., 1977, Vol. 39, pp. 903-905. P¢rgamon Press. Printed in Great Britain

The aquation of the bis(oxalato)ethylenediaminechromate(HI) ion (Received 14 April 1976) We have been working for some time with complexes of chromium(HI) containing both oxalate Cox) and 1,2 diaminoethane (ethylenediamine, en) or ammonia as iigands. These complexes are of particular interest since they can show aquation with either chromium-nitrogen or chromium-oxygen bond breaking. A recent publication [I] has shown that competitive aquatlon can occur for the ion, Cr(en) (ox)2-, I, and suggests that the complete aquation of this ion in acid solutions (pH > 5) needs to be represented by a scheme of the following type: kI

Cr(en)(ox)2(I)

Cr(en)(ox)(H~O)~+ (IV)

were separated using anion and cation exchange columns, as described previously and individual chromatographic fractions were analysed for chromium using an SP90 atomic absorption spectrophotometer [5]. Spectroscopic data were obtained using a Pye-Unicam SP 1800 spectrophotometer coupled to an AR 25 chart recorder.

Kinetics of the initial reaction

Weighed portions (12 mg) of the solid salt K [Cr(en)(oxh]. 2H20 were placed in a 20 mm cell in the thermostatted cell compartment k2 of the spectropbotometer. Aqueous acid (5 cm 3) of the desired > Cr(enH)(ox)e(H20)° ~ Cr(ox)2(H20)2concentration and ionic strength (HCI/NaCI) at the required temperature, was added directly to the cell and, after a short (II) (III) period for mixing, absorbance readings were recorded directly. The aquation of I was followed spectrophotometrically at 476 nm at which wavelength the species II and III have an isosbestic point. a, , Cr(enH)(ox)(H20)s e+ k7 , Cr(ox)(I-I2q), ÷ The neutral complex II was isolated chromatographically and its aquation followed spectrophotometrically at 390 nm. (V) (VI)

The results published on the basis of the above scheme were obtained at pH < 1 and were particularly concerned with reactions involving oxalate loss (k3). The published results are in agreement with those obtained in our laboratories. We publish here results of parallel studies undertaken in the pH range 1-2. This work is a continuation of earlier work on the labilization of chromiumnitrogen bonds by coordinated oxalate[2, 3] and hence is particularly concerned with the reactions involving stepwise loss of ethylenediamine (k,, k2). We have also sought to deal with some previously unresolved problems associated with the interpretation of kinetic data on this system. EXPERIMENTAL The salt K [Cr(en)(ox)2]. 2H20 was prepared by previously published methods [4]. The products of the aquation of the ion, I,

RESULTS AND DISCUSSION In an earlier study at pH > I, using spectroscopic and other techniques, Schiller[6] observed three main stages during the aquation of I, a fast initial reaction, only observed using pH and refractometer measurements (reaction (1), rate constant k,), a slower reaction observed using both pH and spectrophotometric measurements (reaction (2), rate constant k,, with ks -~ 30-80 kl) and a final, very much slower reaction, observed only spectropbotometrically (reaction (3), rate constant k~). Reactions involving oxalate loss were not considered and it has now been shown[l] that such reactions are of minor importance above pH 1.05. The above reaction stages were interpreted in terms of (1) ethylenediamine ring opening to give a neutral species Cr(enH)(ox)2(H20), II; (2), the aquation of II to give the known ion cis-Cr(ox)2(H~O)~-, cis-III; (3), the slow equilibration of

904

Notes to lead to a shift in the positions of the band maxima, whilst reversible water substitution into the inner coordination sphere of chromium(Ill) is known to be a slow process[9], [10]. We conelude that the fast reactions observed here are due to the equilibration of the Cr(en)(ox)2- ion with a small concentration of a 5-coordinated form of the ion having a monodentate oxalate group, possibly accompanied by a steric rearrangement. Such an explanation is consistent with earlier observations of a fast reaction involving pH change, since it might be expected that the one-ended oxalate group would also be available for protonation. The isolation of a monodentate oxalate complex has been reported upon dissolution of the cis-salt K[Cr(ox)2(H20)2]2H20 in acid at low temperature [l l]. Ready oxalate ring opening has already been suggested as an explanation of the remarkably facile acid-catalysed oxalate loss in the present system, the rate of which is 30 times faster than for cis-IIl [1], but no direct evidence of such ring opening was reported. It is now clear that the second stage reaction observed by Schliifer is a two-step process,

cis-llI with its trans-IIl isomer. It follows from this interpretation, with k, >>k,, that a rapid and substantial build up of the neutral intermediate, II, should occur. The species II was not isolated at the time but has now been isolated and characterized in solution, however the maximum yield obtained at p H I was 25% [1]. In our own work at higher pH we have found the species builds up to 40% of the concentration of starting complex but not higher. It is therefore clear that this interpretation of this first fast reaction is not in agreement with more recent work. No alternative explanation has been offered for the initial fast reaction. We have found that when the solid salt K[Cr(en)(ox)2]2H20 dissolves in aqueous solution a rapid decrease in absorbance occurs around the wavelength of maximum absorbance (396 nm). The decrease is approximately exponential with time, half-life about one minute at room temperature, the total decrease being about 3% of total absorbance. Quantitative observations were limited because of the slow dissolution of the solid salt at low temperature and the onset of other aquation reactions at higher temperatures. Conventional log (A - A~) vs time plots gave evidence for a two step reaction. Using estimated A® values for the two steps reasonable straight line plots were obtained for one to two half-lives and values for rate constants obtained. Values are given in Table 1 assuming a reaction scheme,

I

[H ÷]

292.7

0.01 0.05 0.10 0.01 0.05 0.10

282.6

ka/10-2 sec

kdl0 -4 sec-1

1

1.65 1.44 1.60 1,04 1,08 1.28 /~

ko 'lb

h ~iii

for which the rate constants, kl and k2 differ by less than a factor of two. By working at pH 1-2 we have extended published data for the rate of ethylenediamine loss from I and II. Our results for the two steps are compared with published values in Table 2, which also contains data for two related reactions [2],

Table 1. Approximate rate constants for initial fast reactions upon dissolution of K[Cr(en)(ox)2]2H20 Temp/K

k~ ~II

Cr(en)(ox)(H20)2 + IV

10.4 8.2 13,8 6,8 6.7 6.4

k6 , Cr(enH)(ox)(H20)~ V

The earlier failure to distinguish a two step reaction during this stage of aquation arises because of a fortuitous relationship between the relative values of the rate constants k~ and k2 and the relative values of the extinction coefficients for the three species involved, over the wavelength ranges where the total change in absorbance is greatest (350-390, 490-550 nm). In Table 3 the observed absorbance changes for a solution of I are compared with values calculated (i) using appropriate extinction coefficients and rate constants for a two step reaction and (ii) for a simple first order reaction. At the acid concentration used for the reaction shown in Table 3 (0.1 M) there is a small contribution to the reaction by paths involving oxalate loss (12% of the overall product, see Ref. [1]), nevertheless the fit of the data is good for a simple first order reaction and entirely within exlmdmental error for a two step reaction. These results highlight the di~culties inherent in obtaining meaningful kinetic parameters without full chemical analysis of the reacting system. The suggestion that the third stage of the reaction represents the cis-III/trans-III equilibration is now known to be incorrect since the latter is a much faster reaction[7] and although it is likely that cis-III is the initial product of the aquation of II the relative rates of the reactions are such that only an equilibrated mixture of cis-IIl and trans-III could be observed. Reactions involving loss of oxalate and the

~ 'I.

Values of the rate constants have an estimated precision of _+10% and within these limits appear to have no significant acid dependence and only a relatively small temperature dependence. The values of the rate constants reported in Table 1 are consistent with the earlier observation of an initial fast reaction [6]. Since the solid salt K[Cr(en)(ox)2].2H20 can be recovered from its solutions the reactions observed here are probably reversible equilibration reactions. The small decrease in absorbance is not accompanied by any marked shift in the position of the wavelengths of the band maxima. There is now good evidence that reactions leading to isomerization[7] and racemization[8] of cis-III and to 180 exchange with coordinated oxalate[9] occur with fast, reversible oxalate ring opening. The rate of the reactions reported here are of the same order of magnitude as for isomerization. Nitrogenchromium bond breaking might be expected to be irreversible and

Table 2. Rate constants and activation parameters for the stepwise loss of ethylenediamine

Reaction ]

kl

II

~ II

k2 ,III

k/lO-5 sec-I at 298 K

AHS;IkJ mol-~

AS*/JK-t mol '

Reference

1.7

86.1

-47.7

[1]

1.71 -+0.07

83.2 -+0.5

-57.3 -+ 1.3

t

1.9

--

1.42 - 0.05

81.8 - 3.0

-63.5 -+ 1.5

[1] t

IV

k6 ~V

0.63-+0.03

81.5-+1.2

-71-+4

[2]

V

k7 >VI

0.25---0.01

92.4-+0.4

-41.8+-4.2

[2]

tThis work.

k, , Cr(ox)(H20)4 ~ VI

Notes Table3. Comparison of observed and calculated absorbance values for the aquation of Cr(enXox): ion. Temperature 318 K, [H +] 0.1 M, ionic strength 1.0 M

Time/min

Abs. obs.

0 20 40 60 80 100 120 140 160 180 200 220 240 260 280 300 320

0.9OO 0.848 0.803 0.765 0.729 0.698 0.670 0.648 0.625 0.607 0.590 0.578 0.563 0.553 0.544 0.536 0.530 0.480 Standard deviation

(i) Abs. calc. 0.9OO 0.848 0.804 0.764 0.729 0.698 0.671 0.647 0.626 0.607 0.591 0.577 0.564 0.553 0.544 0.535 0.528 0.480

Diff. 0 0.001 0.001 -0.001 0 0 0.001 -0.001 0.001 0 0.001 -0.001 0,001 0 0 -0.001 -0.002 0 0.001

(ii) Abs. calc. 0.9O0 0.848 0.802 0.762 0.727 0.6% 0.669 0.645 0.625 0.607 0.591 0.577 0.565 0.554 0.545 0.537 0.530 0.480

Diff. 0 0 -0.001 -0.003 -0.002 -0.002 -0.001 -0.003 0 0 0.001 -0.001 0.002 0.001 0.001 0.001 0 0 0.001

the only step for which rate data have not now been published. The isomeric composition of the species IV and V are not known exactly but it is likely that only one form of each species is stable in solution[2]. Although reactions involving oxalate loss can be suppressed by working at higher pH, nevertheless the problems originally encountered in the interpretation of kinetic data for this system are understandable. Finally it should be noted that the complexity of the reactions of the Cr(enXoxh ion arises because of the mutual labilization of the coordinated groups. Thus oxalate loss is more rapid than from the aquated ion Cr(oxh(H20)2-[12] or from the ions Cr(NH3)2(ox)(H2Oh +[3] and Cr(en)2(ox)+[2] for which oxalate loss has not been detected below pH O, and ethylenediamine loss (ring-opening) is more rapid than for Cr(en)(H20)43+[13] or Cr(en)(ox)(H20)2 +[2]. The reasons for this lability are not clear but may be linked with the presence of the ethylenediamine in an anionic complex.

Acknowledgement--We thank Miss O. Harper for valuable assistance in the preparation of computer programmes used for the analysis of our kinetic data. Science Department Stockport College o[ Technology Stockport SK 1 3 UQ England

Calculations for (i) based on k~=l.66x104sec ' k2 = 1.22×10 -4 sec-~. Calculations for (ii) based on k= 1.11 × 10-4 sec-'. Absorbance readings at 3% nm, the relevant extinction coefficients are Cr(en)(oxh- 100; Cr(enH)(ox)2(H20), 71; Cr(oxh(H2Oh , 55 cm-' M ~. formation of the ion Cr(ox)(H20), +, IV, are observed in the final slow reaction following the aquation of I. CONCLUSION The scheme given earlier for the aquation of the ion Cr(en)(oxh- (I) needs to be extended to include two preliminary reversible steps In.

•lb.

•I

905

~11. etc.

if consideration is also to be given to the rapid reactions occurring on dissolution of the salt K[Cr(en)(oxh].2HzO. Values are now known for almost all the rate constants in the extended scheme. The rate constants for reactions involving oxalate loss, k3, k4 and k~ are second order rate constants for acid-catalysed reactions, these reactions are very slow above pH 1. The reaction II-,V is

J. W, LETHBRIDGE MICHAEL B. DAVIES

REFERENCES 1. T. W. Kallen, lnorg. Chem. 14, 2687 (1975), 2. M. B. Davies, J. W. Lethbridge, Othman Nor and L-Y. Goh, J. lnorg. Nucl. Chem. 37, 175 (1975). 3. M. B. Davies, J. W. Lethbridge and M. S. Mirrlees, J. Inorg, Nucl, Chem. 35, 3358 (1973). 4. A. Werner, Ann. 406, 286 (1914), 5. M. B. Davies, J. W. Lethbridge and Othman Nor, J. Chromat. 68, 231 (1972). 6. H. L. Schliifer, J. lnorg. Nucl. Chem. 13, 101 (1%0). 7. M. B. Davies and J. W. Lethbridge, J. lnorg. Nucl. Chem. 37, 141 (1975). 8. F. L. Welch and R. E. Hamm, Inorg. Chem. 2, 295 (1%3). 9. J. Aggett, I. Mawson, A. L. Odell and B. E. Smith, J. Chem. Soc. (A), 1413 (1%8). 10. R. A. Plane and H. Taube, J. Phys. Chem. 56, 33 (1952). 11. J. C. Chang, Inorg. Nucl. Chem. Lett. 5, 587 (1%9). 12. D. Banerjea and M. S, Mohan, J. lnorg. Nucl. Chem. 26, 613 (1%4). 13. R. F. Childers Jr., K. G. Vander Zyl Jr., D. A. House, R. G. Hughes and C. S. Garner. Inorg. Chem. 7, 749, 2678 (1%8).

J, inorg,nucl.Chem.,1977.Vol.39. pp, 905-907. PergamonPress. Printedin GreatBritain

The critical

temperatures of

the elements

(Received 3 June 1976) Experimental values of the critical constants of the noble and elementary gases are available and the numerical values are the subject of a recent review by Mathews[1]. In the case of the metallic elements, direct measurements are limited to mercury [25] and caesium[6--8], although by reasonable extrapolation critical data for other alkali metals may be obtained[6]. With the possible exceptions of lithium and cadmium, experimental values for the metallic elements are expected to remain beyond the limits of existing measurement techniques. Consequently, a con-

]INC Vol.39, No. 5--K

siderable effort has been applied to the prediction of values for the critical constants of the metallic elements[9]. In general, the methods involve either extended extrapolation of low temperature physical data or, alternatively interpolation by comparison with the known behaviour of the non metallic elements. Grosse[10], has adopted the former procedure and outlines two essentially independent methods whereby the critical constants of the metallic elements may be estimated. In one of the methods, experimental liquid density and vapour pressure data