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The reaction between oxalatotetraammine-cobalt(III) ion and hydroxide ion in aqueous solution
J. inorg, nucl. Chem. Vol. 40, pp. 1661-1664 © Pergamon Press Ltd., 1978. Printed in Great Britain
0022-1902/78/0901-1661/$02.G(I/0
THE REACTION BETWEEN OXALATOTETRAAMMINE-COBALT(III) ION AND HYDROXIDE ION IN AQUEOUS SOLUTION L. S. BARK Chemistry Department, University of Salford, Salford, England and MICHAEL B. DAVIES* and MALCOLM C. POWELL Science Department, Stockport College of Technology, Stockport SKI 3UQ, England
(Received 2 September 1977; received[or publication 8 February 1978) Abstract--The rate of the reaction between sodium hydroxide and oxalatotetraamminecobalt(III) ion was measured for a variety of hydroxide ion concentrations and at four temperatures. The rate law below 333 K is given by kobs= ko+ k2[OH-]2 and above 333 K is shown to be kob,= ko+ kl[OH-]. The reaction proceeds with a single rate controlling step, which is interpreted as oxalate ring opening. This is followed by a rapid oxalate loss step. INTRODUC~TION There have been several studies of the base hydrolysis of oxalatobis (ethylenediamine) cobalt(Ill) ion[I--4]. The two most recent[2,4] both suggest that the reaction occurs in two stages, the first is the formation of a monodentate oxalato species and the second is the loss of oxalate from this to form the dihydroxo product. However, the two studies differ in that whilst Chan and Harris[4] propose the rate determining step is oxalate ring opening followed by rapid loss of oxalate, Farago and Mason[2] interpreted their data in terms of a fast ring opening reaction followed by the much slower loss of oxalate. These latter workers could find no evidence for a monodentate malonato complex. They have however, isolated monodentate substituted malonato complexes. Interpretation of the mechanisms of base hydrolysis reactions has been made more difficult because there are few data on activation parameters[7]. The purpose of this paper is to provide activation energy data for the oxalato-tetraammine complex and to see if additional kinetic data for a closely related complex can be used to elucidate the mechanisms of the oxalato and malonato ethylenediamine complex base hydrolyses. EXPERIMENTAL
Reagents. All reagents used were B.D.H. AnalaR quality unless otherwise stated. The sodium hydroxide solutions were prepared carbonate free and analysed by the method of Vogel[8]. Oxalatotetraammine-cobalt(IIl) chloride was prepared by a published method[9]. (The perchlorate salt was too insoluble for use in these studies.) Analysis of samples prepared by this method always gave a value for ammonia which was about 3% lower than theoretical despite several recrystallisations in the recommended way. When a solution was passed through a cation exchange column, (zeroiit 225, SCR, 100-200 mesh) it was found that there were two bands. One was eluted by 0.5 M sodium chloride in a way characteristic of a unipositiv¢ cation. The other did not move even with 5.0 M hydrochloric acid. The fraction which was eluted using 0.5 M sodium chloride had a spectrum which was identical, within the experimental error, with the published spectrum of oxalatotetraammine-cobalt(II[) ion. Solutions prepared in this way were used. *Author for correspondence.
Analyses. Cobalt was determined using a Unicam SP90 Atomic Absorption Spectrophotometer or iodometrically. Ammonia was determined using the Kjeldahl technique. Oxalate was determined by titration against potassium permanganate after the complex had been destroyed by boiling it for at least 15 rain with concentrated sodium hydroxide solution. Kinetics. Some kinetic runs were carried out by mixing previously thermostatted reagents in a cell which was then returned to the thermostatted cell holder of a Unicam SP 1800 UV-visible spectrophotometer. The reaction was initially followed by scanning from 350 to 610rim every 4rain and later the reaction was followed at a wavelength of 360 nm. The temperature in the cell was accurate to +-0.1K. When the reaction had proceeded for about a half an hour at room temperature, it was found that a faint brown precipitate of cobalt(Ill) hydroxide began to be visible. This precluded the use of spectrophotometry to follow the reaction. • At known time intervals, 5.0 cm; aliquots of the thermostatied reaction mixture containing the complex ion and sodium hydroxide solution at the appropriate ionic strength were taken. They were passed down 6 cm x 0.5 cm Zerolit 225, S.C.R. 100200 mesh cation exchange columns. The mixture was immediately chilled. The columns acted as a filter for the insoluble cobalt(III) hydroxide and also removed coloured cations while oxalate ion passed straight through. The solutions were forced through the columns under pressure at about 10 cm3 min-~. The eluants were titrated with 0.002 M potassium permanganate solution. A value of V® was obtained by carrying out the same procedure after 24 hr. Kinetic data were obtained using plots of log (V®- Vt) against time. These were linear for at least 75% of the total reaction time. From the plots, pseudo-first order rate constants could be calculated. The thermostat bath maintained the temperature to better than +-0.05K. All runs were repeated at least twice and usually three times and they were usually reproducible to better than +-3% in the pseudo first order rate constant. Throughout the runs the ionic strength was maintained at 1.00M by adding the appropriate weights of solid sodium chloride to each solution prior to the run.
RESULTS AND DISCUSSION
Preliminary studies. Consideration of earlier work[2] suggests convenient temperatures to follow the loss of oxalate from this complex to be 323 K and above, but it was thought possible that there might be a faster reaction which would be too fast at such a high temperature, but which would have a convenient rate at 298 K. After the 1661
1662
L.S. BARK et al.
addition of the base, at room temperature there was no visible appearance of cobalt(III) hydroxide for 30 rain. However at 298 K there was rapid increase in the optical density at 360 nm. After about 5 min, the rate of this reaction decreased and a much slower increase in absorbance was noted. This was considered to be caused by the main reaction. The initial fast reaction gave excellent first order plots using the Geggenheim method [10] and using estimated D® values. The plots were linear for 3 half lives. Furthermore, the "rate constants" obtained from these plots gave the expected first order dependence on hydroxide concentration. Acidification of the reaction mixture caused the spectrum of the "intermediate" to return to that of the starting material. However, it was found that if the reaction mixture was carefully filtered at any point during this apparent initial reaction, the spectrum of the filtrate was almost exactly that of the original mixture and that a slow change in optical density, similar to that expected for the reaction involving the loss of oxalate, ensued (Fig. 1). Although there was no visible appearance of cobalt(III) hydroxide during the "initial reaction", the above results are interpreted as arising from the formation of colloidal suspensions of cobalt(III) hydroxide in the solutions, which gives rise to the observed increase in optical density. The apparent slowing of the reaction may be due to the formation of larger particles of cobalt(III) hydroxide which, in the unstirred solutions, begin to sink below the light path of the spectrophotometer. Many attempts were made chromatographically to separate an intermediate containing monodentate oxalate, such as [Co(NH3)4C204OH]°, from the reaction mixture under a wide variety of hydroxide ion concentrations. However, at no point in the reaction did coloured species pass directly through the column. A number of experiments were performed to try to isolate the intermediate reported to have been detected spectrophotometrically in a kinetic study[2] of the base hydrolysis of oxalatobis(ethylenediamine)cobalt(III) ion. Using conditions identical to those reported and passing solutions down cation exchange columns in the sodium form, it was not possible to isolate any zero charged or
negatively charged species. All cobalt containing ions were always retained on the column. It is possible that the decrease in the hydroxide ion concentration, caused by the necessary dilution of the solution prior to introduction to the column allows the recombination of the oxalate linkage to give the chelate complex. However for solutions at pH 7-8, using a similar technique to that described above, Chan and Harris[4] reported a relatively stable monodentate complex. The increase in optical density observed by Farago and Mason could have been caused by a small amount of precipitation of cobalt(III) hydroxide at the high hydroxide concentrations used. The kinetics
The rate constants for the reaction between sodium hydroxide and oxalatotetraammine-cobalt(III) ion at four temperatures and at a variety of hydroxide ion concentrations and at different ionic strengths are shown in Table 1. When these data are plotted it is clear that the dependence of the rate constant on the hydroxide ion concentration differs according to the temperature of the reaction mixture. Workers on systems similar to that studied here have obtained linear plots of kobs/[OH-] vs [OH-]. For the oxalatotetraammine cobalt(III)-hydroxide ion system, such plots are not linear. It seems that the reason is that, unlike other similar systems, our data gave a small but persistent intercept when kob~was plotted vs [OH-]2. Below 333 K such plots gave good straight lines. Above 333 K, straight lines could only be obtained using plots of kob~ vs [OH-]. Such plots are shown in Fig. 2. We interpret this situation as arising from a rate law of the type: kob~= ko+ kl[OH-] above 333 K and: kobs = ko + k2[OH-] 2 below 333 K.
2.0
I
1.9
J.8
2
I
I
l
0 Time, rain
Fig. 1. (1) Variation of optical density with time before filtering. (2) Variation of optical density with time after filtering.
J 4O
Oxalatotetraammine-cobalt(III) ion Table 1. Rate constants for release of oxalate in the base hydrolysis of oxalatotetraammine-cobalt(IlI) ion
TempfK
Concentration of OH-/M
Pseudo first order rate constant
0.202 0.306 0.370 0.507 0.628 0.050 0.096 0.212 0.367 0.502 0.622 0.105 0.198 0.244 0.363 0.502 0.622 0.095 0.175 0.244 0.297 0.352 0.3%
2.5 --.0.13 4.28 -+0.26 6.56 -+0.43 11.8 2 0.5 17.5 -+ 1.7 0.63 --+0.07 1.820.1 4.92-.0.3 13.0 2 0.6 22.5 -+ 1.5 35.9 -+ 1.5 7.5 -+0.4 12.1 -+0.6 17.2-+ 1.5 30.6 ± 2.6 54.4 -+2.7 83.2 2 6.5 26.4 -+2.6 43.4 2 1.6 68.6 -+4.5 83.1 -+3.5 97.04 2 5.69 116.5 2 3.5
t663
120 IO3 ,~aO
kobsl s -1
b 313
323
333
343
4O 20
Ol
343 333 323 313
kl/s
iM i
k21s-i
2.97 + 0.18 x 10-3 0.67 2 0.12 x 10-3t 0.25 × 10-3t --
AH~-~= 110.1_+12.6kJmol ~ ASi-+= 24.7-+ 1.1JK -I tool-'
02
r
I
!
I
OI
0.2
03
(34
[oH-]
Fig. 2. Plot of k,,b, vs [OH] and tON ]:.
Table 2. Rate constants and activation parameters derived from plots of [OH-] and [OH-] 2 against kobs Temp/K
[oH_]2
M-2
-2.06 2 0.03 x 10-3 0.89-+ 0.02 × 10-3 0.43 20.01 x 10-3
AH2±=65.0_+3.1 kJmol-~ AS2± = - 102.5-+0.3 JK -1 mol-t
fCalculated from the first few points of plots of kobsvs [OH-]. Values of ko, k~ and k2 at various temperatures derived from such plots are given in Table 2. Also included with these data are activation parameters calculated using an Eyring plot[11]. Measurement of the rate of loss of oxalate from oxalatotetraammine-cobalt(III) ion cannot in itself give an unambiguous answer to the question of the nature of the reaction pathway. It cannot distinguish between the possibilities of slow oxalate loss after rapid ring opening and slow ring opening followed by rapid oxalate loss. The early production of insoluble cobalt(Ill) hydroxide precludes a spectrophotometric study to show the existence of a rapid equilibrium and there is no possibility of studying the stereochemical arrangement of the products for the same reason.
The anation of diaquotetraammine-cobalt(IIl) by oxalate has been shown to occur in two stages when the pH is in the range 3--4112]. This was interpreted by the authors as arising from an initial rapid formation of a monodentate oxalato complex followed by a slower reaction. We have used the cation exchange column method of following the release of oxalate with the product of the first reaction, the monodentate oxalato species. A hydroxide ion concentration similar to that used in the runs carried out with the bidentate species was chosen. The oxalate release was extremely rapid, too fast to be measured using the column method. This result and the fact that we have failed to detect any zero charged complex in the reaction mixture suggests that slow ring opening is the first step in the reaction occurring here. This conclusion is in agreement with that of Chan and Harris[4] for the bis(ethylenediamine)oxalato-cobalt(IIl) complex in dilute hydroxide ion solution. The occurrence of terms in [OH-] and [OH ]2 in the rate law is a fairly common feature of base hydrolysis of carboxylato amine cobalt(III) complexes[13, 14]. The kt term has been ascribed to cobalt-oxygen bond fission and the k2 term to carbon-oxygen bond breaking. The [OH-] term is a square term because it is thought that there is a "concerted attack" by the OH- on the carboxyl carbon[7]. This paper confirms a similar situation in the oxalatotetraammine-cobalt(llI) system. Furthermore, the activation enthalpy arising from the k, term is considerably higher than that for the k2 term and this is also in line with data obtained by other workers[13]. Table 3 compares the activation parameters for some other ring opening reactions. The oxalato ring opening has a substantially higher activation enthalpy than either the benzylmalonato or ethylmalonato ethylenediamine complexes. The activation enthalpy of the unsubstituted malonato ethylenediamine complex, however, is very similar to that of the oxalato tetraammine complex. This situation possibly reflects the high stability, of the five membered oxalato ring system [3] and the increased stability of the unsubstituted malonato complex due to the
Table 3. Activation parameters for a variety of ring opening reactions Complex Co en2etmal+ Co en2benzmal+ Co enzmal+ Co(NH3)4ox+
AHt ±
AH2 ±
ASI -~
El mo1-1
El mol-l
JK -I tool-~
65.0
-92 -50.2 95.3 24.7
60.6 73.1 124.1 110.1
AS2-* JK -~ mol t
Ref.
- 102.5
[6] {6] [5] This work
1664
L.S. B A R K et at.
formation of a non reactive resonance stabilised ring when a proton is removed in basic solution from the malonate ring[6]. The very small contribution from a term independent of [OH-] is unusual. Such a term has not been observed in other chelate complex base hydrolysis reaction. The rate constant is attributed to the aquation of the complex.
Acknowledgements--The authors wish to thank Mr. D. Dawson and Mrs. B. Bradley for technical assistance and Dr. J. W. Lethbridge for valuable discussions. One of us (M.B.D.) wishes to thank the Deutsche Akademische Austauschdienst and Prof. K. Wieghardt for provision of facilities at the University of Hanover where part of this work was carried out. REFERENCES
1. S. Sheel, D. R. Meloon and G. M. Harris, lnorg. Chem. 1, 170 (1962). 2. M. E. Farago and C. F. V. Mason, J. Chem. Soc. (A), 3100 (1970).
3. C. Andrade and H. Taube, J. Am. Chem. Soc. 86, 1328 (1964). 4. S. C. Chart and G. M. Harris, lnorg. Chem. 10, 1317 (1971). 5. V. Carunchio, F. Giannetta and G. G. Strazza, J. lnorg. Nucl. Chem. 33, 3025 (1971). 6. M. E. Farago and ]. M. Keefe, Inorg. Chim. Acta 15, 5 (1975). 7. R. G. Wilkins, The Study o[ Mechanism o/Reactions of Transition Metal Complexes, p. 212. Allyn & Bacon, Boston (1974). 8. A. I. Vogel, A Textbook o/Quantitative Analysis, 3rd Edn, p. 284. Longmans, London (1961). 9. W. G. Palmer, Experimental Inorganic Chemistry, p. 547. Cambridge University Press, Cambridge (1959). 10. E. A. Guggenheim, Phil. Mag. 2, 538 (1926). II. Ref. [7], p. 81. 12. M. B. Davies and M. C. Powell, to be published. 13. N. S. Angerman and R. B. Jordan, lnorg. Chem. 6, 1376 (1967). 14. R. Davies, G. B. Evans and R. B. Jordan, Inorg. Chem. 8, 2025 (1969).
0022-1902/78/0901-1661/$02.G(I/0
THE REACTION BETWEEN OXALATOTETRAAMMINE-COBALT(III) ION AND HYDROXIDE ION IN AQUEOUS SOLUTION L. S. BARK Chemistry Department, University of Salford, Salford, England and MICHAEL B. DAVIES* and MALCOLM C. POWELL Science Department, Stockport College of Technology, Stockport SKI 3UQ, England
(Received 2 September 1977; received[or publication 8 February 1978) Abstract--The rate of the reaction between sodium hydroxide and oxalatotetraamminecobalt(III) ion was measured for a variety of hydroxide ion concentrations and at four temperatures. The rate law below 333 K is given by kobs= ko+ k2[OH-]2 and above 333 K is shown to be kob,= ko+ kl[OH-]. The reaction proceeds with a single rate controlling step, which is interpreted as oxalate ring opening. This is followed by a rapid oxalate loss step. INTRODUC~TION There have been several studies of the base hydrolysis of oxalatobis (ethylenediamine) cobalt(Ill) ion[I--4]. The two most recent[2,4] both suggest that the reaction occurs in two stages, the first is the formation of a monodentate oxalato species and the second is the loss of oxalate from this to form the dihydroxo product. However, the two studies differ in that whilst Chan and Harris[4] propose the rate determining step is oxalate ring opening followed by rapid loss of oxalate, Farago and Mason[2] interpreted their data in terms of a fast ring opening reaction followed by the much slower loss of oxalate. These latter workers could find no evidence for a monodentate malonato complex. They have however, isolated monodentate substituted malonato complexes. Interpretation of the mechanisms of base hydrolysis reactions has been made more difficult because there are few data on activation parameters[7]. The purpose of this paper is to provide activation energy data for the oxalato-tetraammine complex and to see if additional kinetic data for a closely related complex can be used to elucidate the mechanisms of the oxalato and malonato ethylenediamine complex base hydrolyses. EXPERIMENTAL
Reagents. All reagents used were B.D.H. AnalaR quality unless otherwise stated. The sodium hydroxide solutions were prepared carbonate free and analysed by the method of Vogel[8]. Oxalatotetraammine-cobalt(IIl) chloride was prepared by a published method[9]. (The perchlorate salt was too insoluble for use in these studies.) Analysis of samples prepared by this method always gave a value for ammonia which was about 3% lower than theoretical despite several recrystallisations in the recommended way. When a solution was passed through a cation exchange column, (zeroiit 225, SCR, 100-200 mesh) it was found that there were two bands. One was eluted by 0.5 M sodium chloride in a way characteristic of a unipositiv¢ cation. The other did not move even with 5.0 M hydrochloric acid. The fraction which was eluted using 0.5 M sodium chloride had a spectrum which was identical, within the experimental error, with the published spectrum of oxalatotetraammine-cobalt(II[) ion. Solutions prepared in this way were used. *Author for correspondence.
Analyses. Cobalt was determined using a Unicam SP90 Atomic Absorption Spectrophotometer or iodometrically. Ammonia was determined using the Kjeldahl technique. Oxalate was determined by titration against potassium permanganate after the complex had been destroyed by boiling it for at least 15 rain with concentrated sodium hydroxide solution. Kinetics. Some kinetic runs were carried out by mixing previously thermostatted reagents in a cell which was then returned to the thermostatted cell holder of a Unicam SP 1800 UV-visible spectrophotometer. The reaction was initially followed by scanning from 350 to 610rim every 4rain and later the reaction was followed at a wavelength of 360 nm. The temperature in the cell was accurate to +-0.1K. When the reaction had proceeded for about a half an hour at room temperature, it was found that a faint brown precipitate of cobalt(Ill) hydroxide began to be visible. This precluded the use of spectrophotometry to follow the reaction. • At known time intervals, 5.0 cm; aliquots of the thermostatied reaction mixture containing the complex ion and sodium hydroxide solution at the appropriate ionic strength were taken. They were passed down 6 cm x 0.5 cm Zerolit 225, S.C.R. 100200 mesh cation exchange columns. The mixture was immediately chilled. The columns acted as a filter for the insoluble cobalt(III) hydroxide and also removed coloured cations while oxalate ion passed straight through. The solutions were forced through the columns under pressure at about 10 cm3 min-~. The eluants were titrated with 0.002 M potassium permanganate solution. A value of V® was obtained by carrying out the same procedure after 24 hr. Kinetic data were obtained using plots of log (V®- Vt) against time. These were linear for at least 75% of the total reaction time. From the plots, pseudo-first order rate constants could be calculated. The thermostat bath maintained the temperature to better than +-0.05K. All runs were repeated at least twice and usually three times and they were usually reproducible to better than +-3% in the pseudo first order rate constant. Throughout the runs the ionic strength was maintained at 1.00M by adding the appropriate weights of solid sodium chloride to each solution prior to the run.
RESULTS AND DISCUSSION
Preliminary studies. Consideration of earlier work[2] suggests convenient temperatures to follow the loss of oxalate from this complex to be 323 K and above, but it was thought possible that there might be a faster reaction which would be too fast at such a high temperature, but which would have a convenient rate at 298 K. After the 1661
1662
L.S. BARK et al.
addition of the base, at room temperature there was no visible appearance of cobalt(III) hydroxide for 30 rain. However at 298 K there was rapid increase in the optical density at 360 nm. After about 5 min, the rate of this reaction decreased and a much slower increase in absorbance was noted. This was considered to be caused by the main reaction. The initial fast reaction gave excellent first order plots using the Geggenheim method [10] and using estimated D® values. The plots were linear for 3 half lives. Furthermore, the "rate constants" obtained from these plots gave the expected first order dependence on hydroxide concentration. Acidification of the reaction mixture caused the spectrum of the "intermediate" to return to that of the starting material. However, it was found that if the reaction mixture was carefully filtered at any point during this apparent initial reaction, the spectrum of the filtrate was almost exactly that of the original mixture and that a slow change in optical density, similar to that expected for the reaction involving the loss of oxalate, ensued (Fig. 1). Although there was no visible appearance of cobalt(III) hydroxide during the "initial reaction", the above results are interpreted as arising from the formation of colloidal suspensions of cobalt(III) hydroxide in the solutions, which gives rise to the observed increase in optical density. The apparent slowing of the reaction may be due to the formation of larger particles of cobalt(III) hydroxide which, in the unstirred solutions, begin to sink below the light path of the spectrophotometer. Many attempts were made chromatographically to separate an intermediate containing monodentate oxalate, such as [Co(NH3)4C204OH]°, from the reaction mixture under a wide variety of hydroxide ion concentrations. However, at no point in the reaction did coloured species pass directly through the column. A number of experiments were performed to try to isolate the intermediate reported to have been detected spectrophotometrically in a kinetic study[2] of the base hydrolysis of oxalatobis(ethylenediamine)cobalt(III) ion. Using conditions identical to those reported and passing solutions down cation exchange columns in the sodium form, it was not possible to isolate any zero charged or
negatively charged species. All cobalt containing ions were always retained on the column. It is possible that the decrease in the hydroxide ion concentration, caused by the necessary dilution of the solution prior to introduction to the column allows the recombination of the oxalate linkage to give the chelate complex. However for solutions at pH 7-8, using a similar technique to that described above, Chan and Harris[4] reported a relatively stable monodentate complex. The increase in optical density observed by Farago and Mason could have been caused by a small amount of precipitation of cobalt(III) hydroxide at the high hydroxide concentrations used. The kinetics
The rate constants for the reaction between sodium hydroxide and oxalatotetraammine-cobalt(III) ion at four temperatures and at a variety of hydroxide ion concentrations and at different ionic strengths are shown in Table 1. When these data are plotted it is clear that the dependence of the rate constant on the hydroxide ion concentration differs according to the temperature of the reaction mixture. Workers on systems similar to that studied here have obtained linear plots of kobs/[OH-] vs [OH-]. For the oxalatotetraammine cobalt(III)-hydroxide ion system, such plots are not linear. It seems that the reason is that, unlike other similar systems, our data gave a small but persistent intercept when kob~was plotted vs [OH-]2. Below 333 K such plots gave good straight lines. Above 333 K, straight lines could only be obtained using plots of kob~ vs [OH-]. Such plots are shown in Fig. 2. We interpret this situation as arising from a rate law of the type: kob~= ko+ kl[OH-] above 333 K and: kobs = ko + k2[OH-] 2 below 333 K.
2.0
I
1.9
J.8
2
I
I
l
0 Time, rain
Fig. 1. (1) Variation of optical density with time before filtering. (2) Variation of optical density with time after filtering.
J 4O
Oxalatotetraammine-cobalt(III) ion Table 1. Rate constants for release of oxalate in the base hydrolysis of oxalatotetraammine-cobalt(IlI) ion
TempfK
Concentration of OH-/M
Pseudo first order rate constant
0.202 0.306 0.370 0.507 0.628 0.050 0.096 0.212 0.367 0.502 0.622 0.105 0.198 0.244 0.363 0.502 0.622 0.095 0.175 0.244 0.297 0.352 0.3%
2.5 --.0.13 4.28 -+0.26 6.56 -+0.43 11.8 2 0.5 17.5 -+ 1.7 0.63 --+0.07 1.820.1 4.92-.0.3 13.0 2 0.6 22.5 -+ 1.5 35.9 -+ 1.5 7.5 -+0.4 12.1 -+0.6 17.2-+ 1.5 30.6 ± 2.6 54.4 -+2.7 83.2 2 6.5 26.4 -+2.6 43.4 2 1.6 68.6 -+4.5 83.1 -+3.5 97.04 2 5.69 116.5 2 3.5
t663
120 IO3 ,~aO
kobsl s -1
b 313
323
333
343
4O 20
Ol
343 333 323 313
kl/s
iM i
k21s-i
2.97 + 0.18 x 10-3 0.67 2 0.12 x 10-3t 0.25 × 10-3t --
AH~-~= 110.1_+12.6kJmol ~ ASi-+= 24.7-+ 1.1JK -I tool-'
02
r
I
!
I
OI
0.2
03
(34
[oH-]
Fig. 2. Plot of k,,b, vs [OH] and tON ]:.
Table 2. Rate constants and activation parameters derived from plots of [OH-] and [OH-] 2 against kobs Temp/K
[oH_]2
M-2
-2.06 2 0.03 x 10-3 0.89-+ 0.02 × 10-3 0.43 20.01 x 10-3
AH2±=65.0_+3.1 kJmol-~ AS2± = - 102.5-+0.3 JK -1 mol-t
fCalculated from the first few points of plots of kobsvs [OH-]. Values of ko, k~ and k2 at various temperatures derived from such plots are given in Table 2. Also included with these data are activation parameters calculated using an Eyring plot[11]. Measurement of the rate of loss of oxalate from oxalatotetraammine-cobalt(III) ion cannot in itself give an unambiguous answer to the question of the nature of the reaction pathway. It cannot distinguish between the possibilities of slow oxalate loss after rapid ring opening and slow ring opening followed by rapid oxalate loss. The early production of insoluble cobalt(Ill) hydroxide precludes a spectrophotometric study to show the existence of a rapid equilibrium and there is no possibility of studying the stereochemical arrangement of the products for the same reason.
The anation of diaquotetraammine-cobalt(IIl) by oxalate has been shown to occur in two stages when the pH is in the range 3--4112]. This was interpreted by the authors as arising from an initial rapid formation of a monodentate oxalato complex followed by a slower reaction. We have used the cation exchange column method of following the release of oxalate with the product of the first reaction, the monodentate oxalato species. A hydroxide ion concentration similar to that used in the runs carried out with the bidentate species was chosen. The oxalate release was extremely rapid, too fast to be measured using the column method. This result and the fact that we have failed to detect any zero charged complex in the reaction mixture suggests that slow ring opening is the first step in the reaction occurring here. This conclusion is in agreement with that of Chan and Harris[4] for the bis(ethylenediamine)oxalato-cobalt(IIl) complex in dilute hydroxide ion solution. The occurrence of terms in [OH-] and [OH ]2 in the rate law is a fairly common feature of base hydrolysis of carboxylato amine cobalt(III) complexes[13, 14]. The kt term has been ascribed to cobalt-oxygen bond fission and the k2 term to carbon-oxygen bond breaking. The [OH-] term is a square term because it is thought that there is a "concerted attack" by the OH- on the carboxyl carbon[7]. This paper confirms a similar situation in the oxalatotetraammine-cobalt(llI) system. Furthermore, the activation enthalpy arising from the k, term is considerably higher than that for the k2 term and this is also in line with data obtained by other workers[13]. Table 3 compares the activation parameters for some other ring opening reactions. The oxalato ring opening has a substantially higher activation enthalpy than either the benzylmalonato or ethylmalonato ethylenediamine complexes. The activation enthalpy of the unsubstituted malonato ethylenediamine complex, however, is very similar to that of the oxalato tetraammine complex. This situation possibly reflects the high stability, of the five membered oxalato ring system [3] and the increased stability of the unsubstituted malonato complex due to the
Table 3. Activation parameters for a variety of ring opening reactions Complex Co en2etmal+ Co en2benzmal+ Co enzmal+ Co(NH3)4ox+
AHt ±
AH2 ±
ASI -~
El mo1-1
El mol-l
JK -I tool-~
65.0
-92 -50.2 95.3 24.7
60.6 73.1 124.1 110.1
AS2-* JK -~ mol t
Ref.
- 102.5
[6] {6] [5] This work
1664
L.S. B A R K et at.
formation of a non reactive resonance stabilised ring when a proton is removed in basic solution from the malonate ring[6]. The very small contribution from a term independent of [OH-] is unusual. Such a term has not been observed in other chelate complex base hydrolysis reaction. The rate constant is attributed to the aquation of the complex.
Acknowledgements--The authors wish to thank Mr. D. Dawson and Mrs. B. Bradley for technical assistance and Dr. J. W. Lethbridge for valuable discussions. One of us (M.B.D.) wishes to thank the Deutsche Akademische Austauschdienst and Prof. K. Wieghardt for provision of facilities at the University of Hanover where part of this work was carried out. REFERENCES
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