Catalytic effects in reactions of chromium(III) — I

j. inorg,nucl.Chem., 1970,Vol.32, pp. 1305to 1311. PergamonPress. Printedin Great Britain

CATALYTIC

EFFECTS IN REACTIONS CHROMIUM(Ill)I

OF

KINETICS OF R E A C T I O N OF C H R O M I U M ( I l l ) WITH E T H Y L E N E D I A M I N E T E T R A A C E T A T E IN THE P R E S E N C E OF C A R B O N A T E , N I T R I T E A N D S U L P H I T E C A T A L Y S T S G. M. P H A T A K , T. R. BHAT* and J. S H A N K A R Bhabha Atomic Research Centre, Chemistry Division, Trombay, Bombay 85, India

(Received 14 February 1969) Al~tract - A detailed study of the kinetics of the reaction between chromium(l I 1) and EDTA has been carried out. Catalytic effect of carbonate, sulphite and nitrite ions on the rate of reaction has been studied at various temperatures. It has been found that it is a first order reaction, with the rate inversely proportional to [H÷]. The catalysts have no effect on either the rate or the order of the reaction. INTRODUCTION

TIlE KINETmS of the reaction of Cr(IlI) with ethylenediaminetetraacetic acid (EDTA) was studied earlier by Harem[l], who found that the slowest step was the formation of the first chelate ring in the product. The catalytic effect of carbonate, sulphite and nitrite ions as well as of some organic solvents such as ethanol, methanol, acetone, dioxane, etc., on the reaction of chromium(Ill) with E D T A was reported by us in an earlier communication[2]. These effects were subsequently found to be of a general nature affecting most of the reactions of chromium, the extent depending on the particular system. In the present study detailed data are presented on the influence of carbonate, nitrite and sulphite catalysts on the reaction of chromium(Ill) with EDTA. This is part of the studies undertaken to elucidate the mechanism of the catalytic action. EXPERIMENTAL Chromium perchlorate was prepared as follows: Freshly precipitated chromium hydroxide was dissolved in dilute nitric acid and the chromium nitrate was crystallized out. The crystals were then dissolved in 2N perchloric acid. The diluted solution was passed through an ion-exchange column. which was washed and then eluted with 2N perchloric acid. Those fractions whose spectrum agreed with that[3] of chromium perchlorate were combined. Chromium perchlorate was preferred to the chloride in order to avoid complex species such as [Cr(H20),~CI] 2÷ etc. EDTA and other chemicals used were of analytical reagent grade purity. As the rate is independent of the concentration of EDTA[I], for maintaining the pH during the course of the reaction, a known excess of the ligand (almost 0-1 M) was added as buffer. Details regarding the experimental procedure have been reported earlier[2]. The concentration of the complex formed was calculated from the absorbance data using the relation * Present address: Semiconductors Ltd., Semiconductor Division, Poona 14, lndia. 1. R. E. Hamm, J. Am. chem. Soc. 75, 5670 (1953). 2. T. R. Bhat and G. M, Phatak, J. inorg, nucl. Chem. 28, 3058 (1966). 3. R. A. Plane andJ. P. Hunt, J . A m , chem. Soc. 79, 3343 (1957). 1305

1306

G . M . P H A T A K , T. R. B H A T and J. S H A N K A R [CrL] = A - ECrtCr(lll)]l ECrL - - ~ C r

where A is the observed absorbance, and the subscript i indicates the amount initially added. The molar absorbance values e for CrL and Cr(l II) were taken as 199 and 8 respectively [2]. The absorbance due to the catalyst or the ligand under the experimental conditions was negligible. The reaction rate was followed to greater than 50 per cent completion and the rate constants were reproducible to 5 per cent or better.

RESULTS

(a) The catalytic effect and the order of reaction The rate of formation of the chromium EDTA complex, both in the presence and absence of catalysts is given by d[CrL]

= ko[Cr3+].

Some typical first order rate law plots in the absence and presence of carbonate, nitrite and sulphite are shown in Fig. 1, from which it is clearly seen that the rates are considerably enhanced by the presence of the catalysts, but the order of the reaction is not altered. It was verified that Beer's law was obeyed in the range of concentrations used. The relative catalytic effect varies as COz2- > SO32- > NO2-. Also the second step (k2) is relatively slower than the first (k0, except in the case of nitrite. These observations are the same as those reported for the chloride system[2]. Some typical values for the two slopes (kl and k2) under different conditions are given in Table 1. It may be pointed out that at higher pH, temperature and catalyst concentration, the second step is not so clear and almost merges with the first. For this reason, the detailed studies were made for the first reaction step only. 2.2

z4 ~ 2.6

~

l.

"l~Lt'~ "~'~x"~'_ -

EDTA

pH-4.6 ' 2.EDTA+2 x 10-4M CARBONATE TEMR. 36"C 3. EDTA+ 1.6x 10-2M NITRITE 4. EOTA÷I-OTxlO-aM SULPHITE

• ,¢o~_o.~o,

2.8

o

3.0

•

4

1

0

I

3.2 3.4

3,6 3.8 4"0

I

0

I

10

I

I

20

I

I

30

f

I

40 Time , rain

I

|

50

I

I

60

J

I

I

70

Fig. 1. T y p i c a l first order rate law plots for the reaction o f C r ( l l I ) perchlorat¢ and E D T A in the presence and absence of the catalyst.

Catalytic effects in reactions of chromium(lI I) - I

1307

Table 1. First order rate law constants for the reaction of chromium(Ill) perchlorate with E D T A in the presence and absence of catalysts at 36 ° and pH = 4-6 concn.

Catalyst

(moles)

(kl x 104)sec-1

(k 2 x 10*)sec-1

7.4 9-6 16.1 22.0 7.4 13-4 17.9 21.0 7-4 23.0 40.2 51-5

3-8 6-3 10.4 16-0 3.8 16.4 25.6 34-5 3-8 18-0 33.3 49.5

Carbonate 0.0002 0-0004 0.001 Nitrite 0-0167 0.0335 0.05 Sulphite 0.001 0.0027 0.0054

(b) Influence of the concentration of catalyst and the relative catalytic effect The effect of concentration of catalyst on the reaction rate was studied in the pH range 4.5 to 6. Some of these values are plotted in Fig. 2. It is seen from this figure that the rate, ko, increases linearly with concentration of the catalyst in such a way as to suggest that it can be resolved into two components, one corresponding to the blank k, and the other to the catalyst kc, thus ko = kb + kc[c] where [c] is the catalyst concentration. The values of kb = (intercepts on the Y-axis) agree well with those experi30

NITRITE

24

1

/

5.5

44

2-

55

36.5

3-

5.0

31

-

_¢ 12(

d 0(

t

o.ol

I

I

I

oo3._o.o4 tCJ mam6/~t

I

3O 24

o

/

O CARBONATE is SULPH ITE

pH t.~'c 1,2- 5.5 44 3,4- 5-5 365 5,6- 5-0 31

1

12( 6( i 0.001

I 0o02

I 0003

I

0004 [C] motes/tit

I

I

0005

0006

Fig. 2. Effect of concentration of catalysts on the reaction rate.

1308

G.M.

P H A T A K , T. R. B H A T and J. S H A N K A R

mentally obtained in the absence of the catalyst. Some typical values of kn and kc are given in Table 2 from which it is clear that the relative catalytic effect varies in the order HCO3- > HSO3- > NOz-. Under the experimental pH conditions the major catalytic species in solution are NO2- and HSO3- in nitrite and sulphite catalysts respectively. But in the case of carbonate at pH = 5, the major species in solution is H2CO3, whereas at pH = 6 the percent HCO3- increases to almost 50. This is also reflected in the values of kc obtained at different pH for carbonate solution (Table 2) which show that kc increases with increasing pH. It may therefore be concluded that HCO3- is a better catalyst than H2COz. (c) Effect of pH The effect of pH on the rate of reaction was studied in the pH range 4 to 6. The plot of log ko vs. pH gave a straight line with slope of n e a r l y - 1, indicating that the rate is inversely proportional to [H+], an unusual observation for a reaction with EDTA. This inverse dependence on [H +] holds good even when the catalyst is present. This is evident from the corresponding linear plots shown in Fig. 3. The values for kc are nearly constant for carbonate, nitrite and sulphite ions in the pH range 5.5-6.2, 5-5.5 and 5-5.5 respectively. At lower pH, the values for k~ are appreciably lowered in all cases. Under the experimental conditions, the solubilities [5] of the probable decomposition products of the catalysts are sufficiently high compared to the concentration of the catalyst used and therefore the observed lowering of the catalytic effect cannot be attributed to the loss of catalyst through decomposition. Rather, this may be due to the formation of undissociated acids H2CO3, HNO2 and H2SO3 (pK values 6.45, 3.35 and 7.20 respectively [6]) which

3.5 3-3 34 I

1- EDTA 2 - EDTA + NITRITE 3- EDTA + CARBONATE 4-EDTA+ SULPHITE

2.9 2,7 == 2-5 2-3 2-1 1,9 1.7

I

4

Fig. 3. Inverse

I "~

5

i

6

i

7 pH

J

8

dependenceof reaction

i

9

10

rate on [H+].

4. H. M. N. H. Irving and W. R. Tomlinson, Chemist Analyst 55, 14 (1966). 5. N. A. Lange, Handbook of Chemistry, pp. 1092-1093. Handbook Publishers, Ohio (1956). 6. Stability constants of metal ion complexes, Spec. Pubis chem. Soc. No. 17, pp. 133, 164, 229 (1964).

4.6 5 5.5 4.6 5 5.5

Sulphite

Nitrite

4.6 5 5.5 5.9

PH

Carbonate

Catalyst

5.9 23.6 61 .O 65.0 0.7 0.55 0.55 0.025 0.035 0.02

1.4 3.6 6.0 10.0

kb kc (lo%b set-‘) (xc-’ mole-‘)

31°C

set-‘)

kb

7.4 38.3 75.0 88.0

( 104kb 2.0 6.0 12.0 12.6 1.1 1.25 1.53 0,035 0.05 0.065

mole-‘)

k-

on k,, and k,

(set-’

36°C

Temperature

Table 2. Effect of pH and temperature

set-‘)

kb

14.5 53.3 123.0 150.0

( 104kb

(see-’

44°C

3.2 8.6 24.0 32.0 2.2 2.1 2.15 0.065 0.23 0.26

mole-‘)

kc

1310

G.M.

P H A T A K , T. R. B H A T and J. S H A N K A R

are less effective in catalysing the reaction. The overall reaction can be represented by ko = k/[H +3 +kc[c]. (d) Effect of temperature The effect of temperature on ko and kc was studied in the temperature range 31 o to 44°C and the values are given in Table 2. From the data the kinetic parameters A H*, A F* and AS were calculated. These are given in Table 3. Table 3. Kinetic parameters for the reaction of c h r o m i u m ( I I I ) perchlorate with E D T A in the p r e s e n c e and absence o f catalysts E Kcal

AH* Kcal

AS cal

AF* Kcal

A

EDTA EDTA + Carbonate EDTA + Sulphite EDTA +

18"2

17"6

- 10"8

20"8

7"8 × 10~

14.1

13.5

-23.9

20.8

1.1 × 1025

14.8

14.2

- 21.8

20.8

1.8 × 10~2

Nitrite

15.1

14.5

-20.7

20.8

7.1 × 10~z

DISCUSSION

Irving and Tomlinson [4] have reported that the addition of a small amount of zinc dust improves the rate of the reaction between chromium(Ill) and EDTA. This has been explained on the basis that part of the Cr(III) is reduced by zinc to Cr(II) which reacts readily with EDTA. The Cr(II)-EDTA complex then undergoes electron transfer reaction with Cr 3+, forming Cr(III)-EDTA complex, thus Cr ~++ E D T A ~ [Cr(II)-EDTA] [Cr(II)-EDTA] + Cr a+ ~ [Cr(III)EDTA] + Cr 2+.

(1) (2)

Both the steps are reported to be quite fast as is also observed in Co 2+ and [Co(III)-EDTA] complexes. These authors assumed that a small amount of Cr 2+ is sufficient to catalyse the reaction. Apparently this possibility should exist also with nitrite and sulphite both of which are reducing agents. However, consideration of oxidation potentials of the systems

HNO2 + H~O ~ NO3- + 3H++ 2eH~SOz+H~O ~ SO4=+4H++2e Cr a+ + e- ~ Cr 2+

-0.94 V -0.17V -0.41 V

would suggest that the amount of Cr 2+ likely to be formed will be negligibly small. Also, in the case of carbonate this type of mechanism is actually excluded. A mechanism has been suggested by Hamm [1] for the possible reaction path between Cr 3+ and EDTA. This assumes that the formation of the first bond between Cr a+ and E D T A is quite fast and the rate determining step is the formation of the chelate ring. The suggested steps are

Catalytic effects in reactions of c h r o m i u m ( I l l ) - I

1311

Cr(H20)63+ + H2Y = ~ B

(1)

B ~K B' + H +

(2)

B' ~

(3)

C.

According to Hamm there is definite evidence for the presence of an intermediate product, whose formation obeys first order kinetics (as indicated by the absorbance reading which corresponds to the fully chelated compound, viz., CrYH20-). Precipitation of chromium does not occur from this intermediate product even at pH = 6, whereas in chromium(Ill) perchlorate, precipitate forms at pH = 4 (pK for hydrolysis of Cr(H20)63+ is 4). This high value compared to chromium perchlorate may be due to partial co-ordination by EDTA. On the basis of these steps Hamm[1] has derived an expression for the first order rate constant kKt [H +] + K

rate =

where k is the rate constant for path (3) and K the equilibrium constant of reaction (2). However, no indication is given whether in the first step any proton is liberated from the ligand. Also, it is not clear whether in the second step the proton is liberated from the ligand or from the coordinated water. The low pK value ( - 3 ) of step (3) is, however, suggestive of its liberation from the coordinated water. If we assume a similar mechanism to operate in the presence of the catalyst, one way of explaining the observed catalytic effect is by assuming an equilibrium of the type B+Cat.. H-Cat.

K1

" B'+H-Cat.

K2 [ C a t . ] + [ H +]

This will lead to an expression for the catalytic path as K1K2 =

[B'] [H +]

[S]

and the overall rate can be represented by rate=

-- K1K2 kt

(K, K2 + [H+])"

This mechanism which is in conformity with most of the experimental data, predicts that the kinetic parameters should remain the same since the rate determining step is not affected by the catalyst but only the actual K value.