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The kinetics of the reaction between sodium hydroxide and diaquo oxalato (ethylenediamine) chromium III ion in aqueous solution
j. inorg, nucl. Chem., 1974,Vol. 36, pp. 3789-3792. PergamonPress. Printed in Great Britain.
THE KINETICS OF THE REACTION BETWEEN SODIUM HYDROXIDE AND DIAQUO OXALATO (ETHYLENEDIAMINE) C H R O M I U M III ION IN AQUEOUS SOLUTION MICHAEL B. DAVIES Science Department, Stockport College of Technology, Wellington Road South, Stockport SKI 3UQ (Received 10 January 1974.)
Abstract--The system Cr en ox (H20)~ + O H - has been studied over a variety of hydroxide ion concenKkl[OH-] . trations at three temperatures. A rate law of the form kobs = ~s suggested. Using the three 1 + K[OH- 1 values of k 1 obtained at three different temperatures by extrapolation an activation energy of 87.57 + 0.54 kJ mol-~ was obtained and log A = 27.35 _+ 0.00. INTRODUCTION 1NARECENTreview[ 1] o f t h e c h e m i s t r y o f a m i n e c h r o m i u m Ill complexes, G a r n e r a n d House have pointed out that there is a great scarcity of data on the base hydrolysis of c h r o m i u m III complexes in aqueous solution. They suggest t h a t this arises, in part, from the ready formation in basic solution ofpolynuclear c h r o m i u m III complexes involving hydroxo or oxo bridges, which complicate any study. As part of a p r o g r a m m e of work on the hydrolysis of oxalato (ethylenediamine) c h r o m i u m III complexes in aqueous solution it was felt that it was of interest to study the kinetics of the reaction of diaquooxalato (ethylenediamine) c h r o m i u m III in aqueous solutions of fairly high pH. Base hydrolysis of a similar cobalt III complex, oxalato bis-(ethylenediamine) cobalt III ion has been studied by a n u m b e r of workers[2-4]. In this reaction the oxalate is fairly readily lost at a b o u t 323°K to form the complex ion Co en2(OH)~ in two stages. This study suggests t h a t the reaction of diaquooxalato (ethylenediamine) c h r o m i u m III with aqueous sodium hydroxide involves replacement of the oxalato group by O H - as in the cobalt III complex above. However, it is suggested that there are fundamental differences in the kinetics which would seem to indicate that a somewhat different m e c h a n i s m is operating. EXPERIMENTAL Reagents
Unless otherwise stated the reagents used were of BDH AnalaR quality. Carbonate-free sodium hydroxide solutions were prepared by the method of Vogel[5] and rejected if they became greater than 2 per cent in carbonate. Diaquooxalato {ethylenediamine) chromium III bromide monohydrate was prepared by the published method[6]. Anal. C, 14,5 ,% : H, 4.3% , N, 8-6~o theory for Crenox (H20)2BrH20, C, 14.3~o; H, 4 . 2 ~ ; N 8-4~.
Solutions containing the tetraquo (ethylenediamine) chromium III ion were prepared by the reaction of 0.5 M perchloric acid with diperoxoaquoethylenediamine chromium IV solution as described by House and Garner[7]. The resulting solution was passed down a 4 x 1 cm Zeocarb 225, 8 per cent DVB, 100-200 mesh cation exchange column in the H + form. The complex cation was then removed from the column with 3 M perchloric acid as described by House and Garner. Product distribution
This was determined by reaction 10-3 M complex with 1-0 x 10-1M solution of sodium hydroxide at 319.7°K. After the required time, 1 cm a of the solution was put on a 4 × 1 cm chromatography column made frpm Isopor Deacidite FF-IP anion exchange resin 100-200 mesh in the hydroxide form. The solution above the column was immediately diluted to 25 cm 3. This mixture was then forced through the column under air pressure to give a flow rate of about 5 cma/min. Aliquots of 10 cm 3 of the eluted solution were taken and the column was washed with water and 0,005 M sodium chloride. The 10 cm 3 aliquots were each analysed for chromium using a Unicam SP 90 atomic absorption spectrophotometer. By repeating this procedure several times using the same concentration of starting material each time, but leaving in the constant temperature bath for varying times, it was possible to build up a picture of the decay and appearance of starting material and products. After each aliquot had been processed, the column was washed with 2 M hydrochloric acid. The above process was repeated using a column made from Permutit Zeo-Karb 225, SRC 15, 100-200 mesh cation exchange resin in the sodium form. This time, however, the 1 cm 3 aliquot was withdrawn and the reaction quenched in 5 cm a 0.i M ice-cold hydrochloric acid before the solution was put on the column. Further, the fractions were eluted with 0.25 M sodium chloride and 3 M hydrochloric acid. The "batchwise" process described for the anion exchange column was not necessary and the same sample was used throughout a given run. After each aliquot had been processed the column was washed with 2 M sodium hydroxide solution.
3789
3790
MICHAEL B, DAVIES
Kinetics The order of the reaction with respect to complex was determined by the method of initial rates[8]. The initial rates with various concentrations of Crenox(H20)~ at a fixed hydroxide concentration of 1.0 × 10 -1 M were determined spectrophotometrically using a Unicam SP 1800 spectrophotometer at a fixed wavelength of 268 nm where the spectrum of the starting material has a shoulder (e = 345) and where a large change in optical density occurs during the reaction. The complex ion solution was thermostatted at the desired temperature and then added to the sodium hydroxide which was in a thermostatted cell holder, using a syringe. The mixture was shaken and the recorder was started immediately the cell was in the instrument, this took 10-15 sees. A plot of log (initial rate) against log (complex concentration) gave a straight line of slope 0-96. The order of the reaction with respect to complex was taken as 1. Similar initial rate experiments were carried out while varying the hydroxide ion concentration and maintaining a constant ionic strength with sodium perchlorate. There was only a fairly small change in initial rate in the range 0.06 to 1.00 M sodium hydroxide and reproducible results could not be obtained. The kinetics of the reaction were followed by taking a thermostatted solution of the complex ( ~ 3 x 10 -3 M) in various concentrations of sodium hydroxide maintained at ionic strength 1.00 with sodium perchlorate. Aliquots of 5 cm 3 of this solution were taken at various times and the reaction quenched by adding to 5-7 cm 3 ice-cold hydrochloric acid of sufficient strength to just acidify the added complex ion solution. The resulting mixture was put on to a cation exchange column similar to that described above. The fraction washed off the column with 0-25 M sodium chloride was collected in a 100 cm 3 graduated flask and the concentration of chromium in each sample was determined using the Unicam SP 90 atomic absorption spectrophotometer. Plots of log (chromium concentration) against time gave excellent straight lines for more than a half life.
Analysis of the products The first product which was washed from a cation exchange column by 3 M acid or from an anion exchange columnby 0.005 M sodium chloride was analysed to determine the ratio of chromium-ethylenediamine-oxalate. Chromium was determined using atomic absorption spectrophotometry on the same sample used to determine ethylenediamine. Oxalate was determined by boiling 100 cm 3 of the product solution containing about 50 ppm Cr with 1 0 ~ potassium hydroxide solution for about 10 min, followed by filtering the chromium III hydroxide produced. The resulting solution was acidified with dilute sulphuric acid and titrated in the usual way with 0.01 M[9] potassium permanganate solution. Ethylenediamine was determined by steam distilling it out of a 100 cm 3 sample of the product, which had been made 10 per cent in sodium hydroxide, into 0.05 M hydrochloric acid. The resulting solution was back-titrated against standard sodium hydroxide. Preliminary experiments with a known solution containing 20ppm ethylenediamine showed that better than 95 per cent of the ethylenediamine was steam distilled into the acid. Many attempts at a Kjeldahl technique for determining nitrogen in the solution were made, but they failed to give reproducible results. The Cr-oxalate-ethylenediamine ratio in the product was found to be 1:0.01:1-1.
I °T'-x. oT "-,\ o
"~
0"~ 0"~ O-t,
0
0"1 --
I
4
r
5
I
I
6
I
7
8
I
9
I
I0
I
II
l
12
pH
Fig. 1. A plot ofpH against optical density for a 1.2 x 1012 M solution of Cr en ox (H20)~- ion. RESULTS AND DISCUSSION (1) The nature of the starting material The complex Crenox ( H 2 0 ) 2 B r H 2 0 was first prepared a n d characterized by Werner[6]. It has been shown to be formed in the acid hydrolysis of oxalatobis (ethylenediamine) c h r o m i u m III ion in aqueous solution either t h e r m a l l y or photochemically[10, 11]. Because it is formed by the removal of ethylenediamine from Cren2ox+(H20)~ " it has been assumed that the substance obtained in the solid form is cis C r e n o x ( H 2 0 ) ; . It is not impossible, however, that a rapid cis-trans isomerism gives rise to a trans product and we are studying this possibility at the moment. The complex cation is adsorbed o n t o an anion exchange column under the conditions of the reaction, i.e. in basic solution. This suggests that the reactant is Crenox(OH)2. This is confirmed by the plot of p H against o.d. for a 1.2 × 1 0 - 2 M solution of the ion (Fig. 1). The estimated approximate p K a values of the first a n d second waters are 5.7 and 8.3 respectively at
f.~
• •
Eluted with 0.25M N o e l from Na + column Eluted with O.O05M N(ICL from OH- column
•
Outed with 3 M HCI from No t column
•
Eluted with 0 . 0 0 5 Noel from O H - column
i
~
,\
o Difference values
E c ~ o4
,a Time,
n~n
Fig. 2. The growth and decay of intermediate speciesformed during base hydrolysis of Cr en ox (H20)~ ion.
Kinetics of the reaction between sodium hydroxide and diaquo oxalato (ethylenediamine) chromium II1 ion 293°K--values which compare favourably with PKa values obtained for other chromium III aquo amine complexes[12]. (211Nature of the products In acid solution oxalato amine complexes of chromium III react primarily by fission of the C r - N bond. It is of interest, therefore to see if a similar reaction occurs in basic solution. The experiments described above show that the first product which is produced in the reaction of diaquooxalato (ethylenediamine) chromium III ion with sodium hydroxide in aqueous solution has virtually no oxalate present and an ethylenediamine-chromium ratio of I.I:I). The obvious possible product in the light of this is therefore tetraaquo (ethylenediamine) chromium III ion [tetrahydroxo (ethylenediamine) chromium III ion under basic conditions]. Accordingly, the spectrum in 1.5 M perchloric acid of the product was compared with that of an authentic sample of Cren (H20)4a+ in the same strength of perchloric acid. They were identical within experimental error. The disappearance of starting material and the appearance and disappearance of the first subsequent products is shown in Fig, 2. The figure contains data accumulated from both anion and cation exchange columns. This confirms that quenching in dilute HC1 does not change the pattern of the reaction when adjusted to the same pH. The spectrum of the first product from the anion exchange column is the same as that from the cation exchange column. The spectrum of the species washed off the anion exchange column with 0.005 M sodium chloride and off the cation exchange column with 0.25 M NaCI both correspond to the spectrum of Crenox (H20)~ under similar conditions. The subsequent products of the reaction are assumed to be dimeric and/or polymeric hydroxo species arising from the reaction of Cren ('I-I20)~ ÷ ion with O H - [7]. When attempts were made to follow the reaction of Cren (H20)~ + with sodium hydroxide solution spectrophotometrically, the complete absence of isosbestic points showed that this reaction does not result in a single product. The behaviour of the later products of the reaction of Crenox (OH)] with O H on cation and anion exchange columns indicates that they must be very highly charged-,negatively in basic solution and positively in acid solution. This is confirmed by the fact that 2.0 M sodium hydroxide solution immediately elutes all these products from a cation exchange column when 5.0 M hydrochloric acid fails to even move them. In the same way hydrochloric acid removes these species from an anion exchange column when saturated sodium chloride solution fails to move them. (3) The kinetics The complexity of this reaction precludes the use of direct spectrophotometric methods to study the kinetics. Recording of spectra at various time intervals
3791
Table 1. Variation of pseudo first order rate constants with hydroxide concentration at various temperatures Temperature (°K) 288.5
[OH]/M 0.0597 0.0797 0.1089 0-1774 1.000 0.597 0.1089 0.183 2.400 0.500 1.0(30
298.7
308.1
0.0555 0-0597 0-1089 0.183 1.000
kobJS-1 x 10s kcajjS-1 x l0 s 8.50 + 0.16 8.25 + 0.39 9.11 +__0.34 9.28 + 0.21 10.47 4- 0.45 30-9 4- 0-71 32-2 4- 1-4 33.9 4- 0.1 33.7 4- 0.14 34-6 4- 0.78 36-4 4- 0.74 95.7 96.9 100.7 103.4 105.6
4- 2.9 4- 2.8 4- 4.0 + 3-3 4- 2.8
8.20 8.65 9.04 9.52 10-2 30.6 32-7 33-8 34.2 35.0 35.3 96-0 96.6 100.8 103.1 108.5
Errors are standard deviations. shows that the mixture does not exhibit isosbestic points. When the disappearance of the starting material is studied at 268 nm, a plot of log (D-D~o) against time is riot linear. The starting material is efficiently washed off a cation exchange colunm with 0-25 M sodium chloride. The first product is eluted by 3.0 M hydrochloric acid. The difference between these concentrations means that the disappearance of the starting material is readily followed by eluting it with 0-25 M NaC1.
?,
o
5
~
~° 95F-"
%
32 r
27 IO~
IOC _~
,~,~ I ~ ~
e~'~°
~°
I / [OH-],
mole-I drn3
Fig. 3. A plot of reciprocal of hydroxide ion concentration against reciprocal of observed first order rate constant for base hydrolysis of Cr en ox (H20)2~.
3792
MICHAEL B. DAVIES
The observed first order rate constants obtained using the cation exchange method are shown in Table 1. The variation of rate constants with hydroxide ion concentration when the latter is in large excess at a given temperature is small, but increases with temperature. A plot of reciprocal of hydroxide ion concentration against reciprocal of observed first order rate constant is shown in Fig. 3 for three temperatures. Fairly good straight lines are obtained at all three temperatures. This suggests a rate law of the type: ~obs
Kkl[OH-] = 1 + K[OH-]
1 1 From the plot of kZoo, against ~ K k t and the intercept will be
the slope will be
.
A mechanism which gives rise to a rate law of this type must involve a preequilibrium step, with equilibrium constant K, followed by a first-order ratedetermining step whose rate constant is k t . For the base hydrolysis of Crenox(H20)~- a mechanism of the following type would give the above rate law: K
Crenox(OH)~- + OH = ~ Crenox(OH)~"one ended" oxalate
Table 2. Values of kt and K at various temperatures Temperature (°K)
kl/S- 1
K/~ 1 dm 3
288-5 298.7 308.1
1.04 × 10 -4 + 0.04 3-57 × 10-4 _ 0.06 10.64 x 10 -4 + 0.03
167 + 9 I01 _ 25 64 +__17
Errors are standard deviations. Fig. 4. From this a value of 87.57 + 0.54 kJ m o l - 1 was obtained for the activation energy. The values of kl calculated from the intercepts at various temperatures are shown in Table 2. Also included in this table are values of K calculated from the slopes of the three graphs. A similar rate law to the above could be produced by a conjugate base mechanism in which the hydroxide ion abstracts a proton from the ethylenediamine ligands[13]. While this remains a possibility, it seems unlikely that such a mechanism would explain the relatively small effect of hydroxide ion o n the rate. The small effect of hydroxide ion implies a fairly large equilibrium constant for the pre-equilibrium reaction and amines in transition metal complexes are not often very acidic.
kl
Crenox(OH)J-
s~ow
ratedetermining
) Cren(OH)a + ox 2-
Cren(OH)a + O H - f ~ Cren(OH)2 ,LoHpolynuclear species. A computer was used to determine the best straight line through the experimental plots. The computed values are shown in Table 1 and the values found are well within experimental error ( ~ + 3per cent). The intercept was calculated using the computed line and k~ was calculated from this at the three tem1 peratures. A plot of T-K against log kt is shown in 4"0 D
B*
3'8 3'6 i
3.4 3"2 3"0
3-3
3"4 I/T,
*K - i
Fig 4. Arrhenius plot for the base hydrolysis of Cr en ox (r~Ot:+.
Acknowledgements--The author wishes tO thank Mrs. B. Bradley for technical assistance and Miss O. Harper for computing. A personal grant from the Chemical Society and an S.R.C. grant towards an SP 1800 spectrophotometer are gratefully acknowledged.
REFERENCES 1. C. S. Garner and D. A. House, Transition Metal Chemistry 6, 211 (1970). 2. S. Sheel, D. T. Meloon and G. M. Harris, Inorg. Chem. 1 170 (1962). 3. C. Andrade and H. Taube, J. Am. chem. Soc. 86, 1328 (1964). 4. M." E. Forago and C. F. V. Mason, J. chem. Soc. (A), 3100 (1970). 5. A. I. Vogel, A Textbook of Quantitative Analysis. p.241. Longmans, London (1961). 6. A. Werner, Ann. Chem. 4115,212 (1914). 7. D. A. House and C. S. Garner, lnorg. Chem. 5, 840 (1966). 8. A.A. Frost and R. G. Pearson, Kinetics and Mechanism, pp. 45 and 46 2nd Edn. Wiley, New York (1961). 9. A.I. Vogel, A Textbook of Quantitative Analysis. p. 284, 3rd Edn. Longmans, London (1961). 10. M. B. Davies, J. W. Lethbridge and Othman Nor. J. Chromatogr. 68, 231 (1972). 11. A. D. Kirk, K. L. Moss and J. G. Valentin, J. Chromatogr. 36, 332 (1968). 12. C. S. Garner and D. A. House, Transition Metal Chemistry. 6, 177 (1970). 13. M. L. Tobe, Inorganic Reaction Mechanisms. p.95. Nelson, London (1972).
THE KINETICS OF THE REACTION BETWEEN SODIUM HYDROXIDE AND DIAQUO OXALATO (ETHYLENEDIAMINE) C H R O M I U M III ION IN AQUEOUS SOLUTION MICHAEL B. DAVIES Science Department, Stockport College of Technology, Wellington Road South, Stockport SKI 3UQ (Received 10 January 1974.)
Abstract--The system Cr en ox (H20)~ + O H - has been studied over a variety of hydroxide ion concenKkl[OH-] . trations at three temperatures. A rate law of the form kobs = ~s suggested. Using the three 1 + K[OH- 1 values of k 1 obtained at three different temperatures by extrapolation an activation energy of 87.57 + 0.54 kJ mol-~ was obtained and log A = 27.35 _+ 0.00. INTRODUCTION 1NARECENTreview[ 1] o f t h e c h e m i s t r y o f a m i n e c h r o m i u m Ill complexes, G a r n e r a n d House have pointed out that there is a great scarcity of data on the base hydrolysis of c h r o m i u m III complexes in aqueous solution. They suggest t h a t this arises, in part, from the ready formation in basic solution ofpolynuclear c h r o m i u m III complexes involving hydroxo or oxo bridges, which complicate any study. As part of a p r o g r a m m e of work on the hydrolysis of oxalato (ethylenediamine) c h r o m i u m III complexes in aqueous solution it was felt that it was of interest to study the kinetics of the reaction of diaquooxalato (ethylenediamine) c h r o m i u m III in aqueous solutions of fairly high pH. Base hydrolysis of a similar cobalt III complex, oxalato bis-(ethylenediamine) cobalt III ion has been studied by a n u m b e r of workers[2-4]. In this reaction the oxalate is fairly readily lost at a b o u t 323°K to form the complex ion Co en2(OH)~ in two stages. This study suggests t h a t the reaction of diaquooxalato (ethylenediamine) c h r o m i u m III with aqueous sodium hydroxide involves replacement of the oxalato group by O H - as in the cobalt III complex above. However, it is suggested that there are fundamental differences in the kinetics which would seem to indicate that a somewhat different m e c h a n i s m is operating. EXPERIMENTAL Reagents
Unless otherwise stated the reagents used were of BDH AnalaR quality. Carbonate-free sodium hydroxide solutions were prepared by the method of Vogel[5] and rejected if they became greater than 2 per cent in carbonate. Diaquooxalato {ethylenediamine) chromium III bromide monohydrate was prepared by the published method[6]. Anal. C, 14,5 ,% : H, 4.3% , N, 8-6~o theory for Crenox (H20)2BrH20, C, 14.3~o; H, 4 . 2 ~ ; N 8-4~.
Solutions containing the tetraquo (ethylenediamine) chromium III ion were prepared by the reaction of 0.5 M perchloric acid with diperoxoaquoethylenediamine chromium IV solution as described by House and Garner[7]. The resulting solution was passed down a 4 x 1 cm Zeocarb 225, 8 per cent DVB, 100-200 mesh cation exchange column in the H + form. The complex cation was then removed from the column with 3 M perchloric acid as described by House and Garner. Product distribution
This was determined by reaction 10-3 M complex with 1-0 x 10-1M solution of sodium hydroxide at 319.7°K. After the required time, 1 cm a of the solution was put on a 4 × 1 cm chromatography column made frpm Isopor Deacidite FF-IP anion exchange resin 100-200 mesh in the hydroxide form. The solution above the column was immediately diluted to 25 cm 3. This mixture was then forced through the column under air pressure to give a flow rate of about 5 cma/min. Aliquots of 10 cm 3 of the eluted solution were taken and the column was washed with water and 0,005 M sodium chloride. The 10 cm 3 aliquots were each analysed for chromium using a Unicam SP 90 atomic absorption spectrophotometer. By repeating this procedure several times using the same concentration of starting material each time, but leaving in the constant temperature bath for varying times, it was possible to build up a picture of the decay and appearance of starting material and products. After each aliquot had been processed, the column was washed with 2 M hydrochloric acid. The above process was repeated using a column made from Permutit Zeo-Karb 225, SRC 15, 100-200 mesh cation exchange resin in the sodium form. This time, however, the 1 cm 3 aliquot was withdrawn and the reaction quenched in 5 cm a 0.i M ice-cold hydrochloric acid before the solution was put on the column. Further, the fractions were eluted with 0.25 M sodium chloride and 3 M hydrochloric acid. The "batchwise" process described for the anion exchange column was not necessary and the same sample was used throughout a given run. After each aliquot had been processed the column was washed with 2 M sodium hydroxide solution.
3789
3790
MICHAEL B, DAVIES
Kinetics The order of the reaction with respect to complex was determined by the method of initial rates[8]. The initial rates with various concentrations of Crenox(H20)~ at a fixed hydroxide concentration of 1.0 × 10 -1 M were determined spectrophotometrically using a Unicam SP 1800 spectrophotometer at a fixed wavelength of 268 nm where the spectrum of the starting material has a shoulder (e = 345) and where a large change in optical density occurs during the reaction. The complex ion solution was thermostatted at the desired temperature and then added to the sodium hydroxide which was in a thermostatted cell holder, using a syringe. The mixture was shaken and the recorder was started immediately the cell was in the instrument, this took 10-15 sees. A plot of log (initial rate) against log (complex concentration) gave a straight line of slope 0-96. The order of the reaction with respect to complex was taken as 1. Similar initial rate experiments were carried out while varying the hydroxide ion concentration and maintaining a constant ionic strength with sodium perchlorate. There was only a fairly small change in initial rate in the range 0.06 to 1.00 M sodium hydroxide and reproducible results could not be obtained. The kinetics of the reaction were followed by taking a thermostatted solution of the complex ( ~ 3 x 10 -3 M) in various concentrations of sodium hydroxide maintained at ionic strength 1.00 with sodium perchlorate. Aliquots of 5 cm 3 of this solution were taken at various times and the reaction quenched by adding to 5-7 cm 3 ice-cold hydrochloric acid of sufficient strength to just acidify the added complex ion solution. The resulting mixture was put on to a cation exchange column similar to that described above. The fraction washed off the column with 0-25 M sodium chloride was collected in a 100 cm 3 graduated flask and the concentration of chromium in each sample was determined using the Unicam SP 90 atomic absorption spectrophotometer. Plots of log (chromium concentration) against time gave excellent straight lines for more than a half life.
Analysis of the products The first product which was washed from a cation exchange column by 3 M acid or from an anion exchange columnby 0.005 M sodium chloride was analysed to determine the ratio of chromium-ethylenediamine-oxalate. Chromium was determined using atomic absorption spectrophotometry on the same sample used to determine ethylenediamine. Oxalate was determined by boiling 100 cm 3 of the product solution containing about 50 ppm Cr with 1 0 ~ potassium hydroxide solution for about 10 min, followed by filtering the chromium III hydroxide produced. The resulting solution was acidified with dilute sulphuric acid and titrated in the usual way with 0.01 M[9] potassium permanganate solution. Ethylenediamine was determined by steam distilling it out of a 100 cm 3 sample of the product, which had been made 10 per cent in sodium hydroxide, into 0.05 M hydrochloric acid. The resulting solution was back-titrated against standard sodium hydroxide. Preliminary experiments with a known solution containing 20ppm ethylenediamine showed that better than 95 per cent of the ethylenediamine was steam distilled into the acid. Many attempts at a Kjeldahl technique for determining nitrogen in the solution were made, but they failed to give reproducible results. The Cr-oxalate-ethylenediamine ratio in the product was found to be 1:0.01:1-1.
I °T'-x. oT "-,\ o
"~
0"~ 0"~ O-t,
0
0"1 --
I
4
r
5
I
I
6
I
7
8
I
9
I
I0
I
II
l
12
pH
Fig. 1. A plot ofpH against optical density for a 1.2 x 1012 M solution of Cr en ox (H20)~- ion. RESULTS AND DISCUSSION (1) The nature of the starting material The complex Crenox ( H 2 0 ) 2 B r H 2 0 was first prepared a n d characterized by Werner[6]. It has been shown to be formed in the acid hydrolysis of oxalatobis (ethylenediamine) c h r o m i u m III ion in aqueous solution either t h e r m a l l y or photochemically[10, 11]. Because it is formed by the removal of ethylenediamine from Cren2ox+(H20)~ " it has been assumed that the substance obtained in the solid form is cis C r e n o x ( H 2 0 ) ; . It is not impossible, however, that a rapid cis-trans isomerism gives rise to a trans product and we are studying this possibility at the moment. The complex cation is adsorbed o n t o an anion exchange column under the conditions of the reaction, i.e. in basic solution. This suggests that the reactant is Crenox(OH)2. This is confirmed by the plot of p H against o.d. for a 1.2 × 1 0 - 2 M solution of the ion (Fig. 1). The estimated approximate p K a values of the first a n d second waters are 5.7 and 8.3 respectively at
f.~
• •
Eluted with 0.25M N o e l from Na + column Eluted with O.O05M N(ICL from OH- column
•
Outed with 3 M HCI from No t column
•
Eluted with 0 . 0 0 5 Noel from O H - column
i
~
,\
o Difference values
E c ~ o4
,a Time,
n~n
Fig. 2. The growth and decay of intermediate speciesformed during base hydrolysis of Cr en ox (H20)~ ion.
Kinetics of the reaction between sodium hydroxide and diaquo oxalato (ethylenediamine) chromium II1 ion 293°K--values which compare favourably with PKa values obtained for other chromium III aquo amine complexes[12]. (211Nature of the products In acid solution oxalato amine complexes of chromium III react primarily by fission of the C r - N bond. It is of interest, therefore to see if a similar reaction occurs in basic solution. The experiments described above show that the first product which is produced in the reaction of diaquooxalato (ethylenediamine) chromium III ion with sodium hydroxide in aqueous solution has virtually no oxalate present and an ethylenediamine-chromium ratio of I.I:I). The obvious possible product in the light of this is therefore tetraaquo (ethylenediamine) chromium III ion [tetrahydroxo (ethylenediamine) chromium III ion under basic conditions]. Accordingly, the spectrum in 1.5 M perchloric acid of the product was compared with that of an authentic sample of Cren (H20)4a+ in the same strength of perchloric acid. They were identical within experimental error. The disappearance of starting material and the appearance and disappearance of the first subsequent products is shown in Fig, 2. The figure contains data accumulated from both anion and cation exchange columns. This confirms that quenching in dilute HC1 does not change the pattern of the reaction when adjusted to the same pH. The spectrum of the first product from the anion exchange column is the same as that from the cation exchange column. The spectrum of the species washed off the anion exchange column with 0.005 M sodium chloride and off the cation exchange column with 0.25 M NaCI both correspond to the spectrum of Crenox (H20)~ under similar conditions. The subsequent products of the reaction are assumed to be dimeric and/or polymeric hydroxo species arising from the reaction of Cren ('I-I20)~ ÷ ion with O H - [7]. When attempts were made to follow the reaction of Cren (H20)~ + with sodium hydroxide solution spectrophotometrically, the complete absence of isosbestic points showed that this reaction does not result in a single product. The behaviour of the later products of the reaction of Crenox (OH)] with O H on cation and anion exchange columns indicates that they must be very highly charged-,negatively in basic solution and positively in acid solution. This is confirmed by the fact that 2.0 M sodium hydroxide solution immediately elutes all these products from a cation exchange column when 5.0 M hydrochloric acid fails to even move them. In the same way hydrochloric acid removes these species from an anion exchange column when saturated sodium chloride solution fails to move them. (3) The kinetics The complexity of this reaction precludes the use of direct spectrophotometric methods to study the kinetics. Recording of spectra at various time intervals
3791
Table 1. Variation of pseudo first order rate constants with hydroxide concentration at various temperatures Temperature (°K) 288.5
[OH]/M 0.0597 0.0797 0.1089 0-1774 1.000 0.597 0.1089 0.183 2.400 0.500 1.0(30
298.7
308.1
0.0555 0-0597 0-1089 0.183 1.000
kobJS-1 x 10s kcajjS-1 x l0 s 8.50 + 0.16 8.25 + 0.39 9.11 +__0.34 9.28 + 0.21 10.47 4- 0.45 30-9 4- 0-71 32-2 4- 1-4 33.9 4- 0.1 33.7 4- 0.14 34-6 4- 0.78 36-4 4- 0.74 95.7 96.9 100.7 103.4 105.6
4- 2.9 4- 2.8 4- 4.0 + 3-3 4- 2.8
8.20 8.65 9.04 9.52 10-2 30.6 32-7 33-8 34.2 35.0 35.3 96-0 96.6 100.8 103.1 108.5
Errors are standard deviations. shows that the mixture does not exhibit isosbestic points. When the disappearance of the starting material is studied at 268 nm, a plot of log (D-D~o) against time is riot linear. The starting material is efficiently washed off a cation exchange colunm with 0-25 M sodium chloride. The first product is eluted by 3.0 M hydrochloric acid. The difference between these concentrations means that the disappearance of the starting material is readily followed by eluting it with 0-25 M NaC1.
?,
o
5
~
~° 95F-"
%
32 r
27 IO~
IOC _~
,~,~ I ~ ~
e~'~°
~°
I / [OH-],
mole-I drn3
Fig. 3. A plot of reciprocal of hydroxide ion concentration against reciprocal of observed first order rate constant for base hydrolysis of Cr en ox (H20)2~.
3792
MICHAEL B. DAVIES
The observed first order rate constants obtained using the cation exchange method are shown in Table 1. The variation of rate constants with hydroxide ion concentration when the latter is in large excess at a given temperature is small, but increases with temperature. A plot of reciprocal of hydroxide ion concentration against reciprocal of observed first order rate constant is shown in Fig. 3 for three temperatures. Fairly good straight lines are obtained at all three temperatures. This suggests a rate law of the type: ~obs
Kkl[OH-] = 1 + K[OH-]
1 1 From the plot of kZoo, against ~ K k t and the intercept will be
the slope will be
.
A mechanism which gives rise to a rate law of this type must involve a preequilibrium step, with equilibrium constant K, followed by a first-order ratedetermining step whose rate constant is k t . For the base hydrolysis of Crenox(H20)~- a mechanism of the following type would give the above rate law: K
Crenox(OH)~- + OH = ~ Crenox(OH)~"one ended" oxalate
Table 2. Values of kt and K at various temperatures Temperature (°K)
kl/S- 1
K/~ 1 dm 3
288-5 298.7 308.1
1.04 × 10 -4 + 0.04 3-57 × 10-4 _ 0.06 10.64 x 10 -4 + 0.03
167 + 9 I01 _ 25 64 +__17
Errors are standard deviations. Fig. 4. From this a value of 87.57 + 0.54 kJ m o l - 1 was obtained for the activation energy. The values of kl calculated from the intercepts at various temperatures are shown in Table 2. Also included in this table are values of K calculated from the slopes of the three graphs. A similar rate law to the above could be produced by a conjugate base mechanism in which the hydroxide ion abstracts a proton from the ethylenediamine ligands[13]. While this remains a possibility, it seems unlikely that such a mechanism would explain the relatively small effect of hydroxide ion o n the rate. The small effect of hydroxide ion implies a fairly large equilibrium constant for the pre-equilibrium reaction and amines in transition metal complexes are not often very acidic.
kl
Crenox(OH)J-
s~ow
ratedetermining
) Cren(OH)a + ox 2-
Cren(OH)a + O H - f ~ Cren(OH)2 ,LoHpolynuclear species. A computer was used to determine the best straight line through the experimental plots. The computed values are shown in Table 1 and the values found are well within experimental error ( ~ + 3per cent). The intercept was calculated using the computed line and k~ was calculated from this at the three tem1 peratures. A plot of T-K against log kt is shown in 4"0 D
B*
3'8 3'6 i
3.4 3"2 3"0
3-3
3"4 I/T,
*K - i
Fig 4. Arrhenius plot for the base hydrolysis of Cr en ox (r~Ot:+.
Acknowledgements--The author wishes tO thank Mrs. B. Bradley for technical assistance and Miss O. Harper for computing. A personal grant from the Chemical Society and an S.R.C. grant towards an SP 1800 spectrophotometer are gratefully acknowledged.
REFERENCES 1. C. S. Garner and D. A. House, Transition Metal Chemistry 6, 211 (1970). 2. S. Sheel, D. T. Meloon and G. M. Harris, Inorg. Chem. 1 170 (1962). 3. C. Andrade and H. Taube, J. Am. chem. Soc. 86, 1328 (1964). 4. M." E. Forago and C. F. V. Mason, J. chem. Soc. (A), 3100 (1970). 5. A. I. Vogel, A Textbook of Quantitative Analysis. p.241. Longmans, London (1961). 6. A. Werner, Ann. Chem. 4115,212 (1914). 7. D. A. House and C. S. Garner, lnorg. Chem. 5, 840 (1966). 8. A.A. Frost and R. G. Pearson, Kinetics and Mechanism, pp. 45 and 46 2nd Edn. Wiley, New York (1961). 9. A.I. Vogel, A Textbook of Quantitative Analysis. p. 284, 3rd Edn. Longmans, London (1961). 10. M. B. Davies, J. W. Lethbridge and Othman Nor. J. Chromatogr. 68, 231 (1972). 11. A. D. Kirk, K. L. Moss and J. G. Valentin, J. Chromatogr. 36, 332 (1968). 12. C. S. Garner and D. A. House, Transition Metal Chemistry. 6, 177 (1970). 13. M. L. Tobe, Inorganic Reaction Mechanisms. p.95. Nelson, London (1972).