The effect of ion-pairing on the kinetics of the acid hydrolysis of chloropentaamine cobalt(III) ion—I. Effect of sulphate on the aquation of chloropentaamine cobalt(III) ion

J. inorg, nucl. Chem., 1974,Vol.36, pp. 3793-3800.PergamonPress.Printedin Great Britain.

THE EFFECT OF ION-PAIRING ON THE KINETICS OF THE ACID HYDROLYSIS OF CHLOROPENTAAMINE COBALT(III) ION--I. EFFECT OF SULPHATE ON THE AQUATION OF CHLOROPENTAAMINE COBALT(Ill) ION N. C. NAIK* and R. K. NANDA'~ Department of Chemistry, Utkal University, Bhubaneswar-4, India (First received 18 May 1971 ; in revised form 18 July 1973) A~tract--The rate of solvolytic aquation of chloropentaammine cobalt(llI) ion has been measured in aqueous medium in the presence of sulphate anion at 40°, 45°, 50° and 55°C. The rate accelerating effect. of sulphate has been ascribed to more labile ion-pairs and the association constants and rate constants of the ion-pair have been calculated. The product analysis has been carried out to investigate the participation of the ion-pairing anion and the activation parameters of the aquation reaction of the ion-pair and different paths of reaction and enthalpy, entropy and free energy changes of the association equilibrium have been computed and discussed.

INTRODUCTION

Tr~ ACIDhydrolysis reaction of chloropentaammine cobalt(III) ion has been extensively studied[l, 2]. The rate of the reaction has been shown to be independent of hydrogen ion concentration below pH7 [3-5] and of ionic strength[5]. However, the presence of certain anions has been known to accelerate the acid hydrolysis of the complex. Garrick[6, 7] first reported that the rate of aquation of chloropentaammine cobalt(III) ion was slightly increased by the addition of chloride, nitrate, chlorate, formate and acetate, whereas sulphate caused a marked acceleration. The accelerating effect of certain anions on the rate of solvolytic aquation of other halopentaammine cobalt(III)[1] and chromium (III)[8] complexes has also been observed. This accelerating influence has been ascribed to ion-pair formation between the complex cation and the added anions. The role of ion-pairs was discussed by Wyatt and Davies[9], who assumed that the free ion and the ion-pair have characteristically different rate constants. Laurie and Monk[4] applied the Wyatt-Davies treatment to study the effect of ion-pairs on the rate of acid hydrolysis of chloropentaammine cobalt(III) complex in an aqueous medium. The results reveal that chloride, azide, acetate, nitrate and glycollate ions cause slight increase in rate while the bivalent anions, sulphate, malonate and phthalate proved to be strong accelerators. Similar results were also obtained from the aquation studies of the chloropentaammine chromium (lII) system[10] in the presence of a variety of anions. * Present address: Department of Chemistry, Rourkela Science College, Rourkela-2, India. t To whom correspondence should be directed. LI.N.C., Vo|. 36, No. 12--K

More recent work of Monk et al.[ll] substantiated~ the accelerating influence of sulphate and some dicarboxylate anions on the aquation kinetics of halopentaammine cobalt(III) complexes. The data were analysed as before[4]. The investigations of Monk et al. clearly show that the ion-pairs, formed between some bivalent anions and halopentaammine cobalt(Ill) or chromium(Ill) cations undergo aquation at a faster rate compared to the free cations. It is now fairly well established that halopentaammine cobalt(Ill) complexes aquate by an essentially dissociative process. The mechanistic pathway for the aquation reaction of the ion-pairs is, however, unknown. The present investigation was undertaken in an attempt, to elucidate the possible path-ways for the aquation reaction of such ion-pairs. Further, it was of interest to find out if the added anions which form ion-pairs of the outer-sphere type participate in any way in the actual act of substitution at the cobalt(III) centre. With this in view, in addition to the determination of the rate constants for the release of chloride ion from the complex, spectrophotometric examination of the products of the reaction has been carried out. The activation parameters of each reaction path have been evaluated.

EXPERIMENTAL Preparation of the complex

Chloropentaammine cobalt(III) perchlorate, [Co(NHa)5C1](C104)2 was prepared and purified as suggested by Chan et al.[ 12].The purity of the sample was checked by estimating its chloride and cobalt contents.

3793

3794

N . C . NAIK and R. K. NANDA

F o u n d : Co, 15.35 per cent; C1, 9.3 per cent: [Co(NH3)sCI ](C104)2 requires Co, 15-57 per cent; CI, 9-4 per cent. The measured molar absorptivity ( c m - l m - 1 ) of Esa 2 = 50.5, is in excellent agreement with the value of 50.5 reported in literature [13]. A q u o p e n t a a m m i n e cobalt(Ill) perchlorate, [Co(NH3) sH 2 0 ] (CIO4)a was prepared and purified as described in the literature[14]. The measured molar absorptivity of E492 = 47.8 is in fair agreement with the values (E49247.7[13] and E49147.3[14]) reported previously. Sulphatopentaammine cobalt(Ill) perchlorate, [Co(NHa)sSO4](C104) was prepared from sulphatopentaammine bisulphate complex, [Co(NHa)sSO,~]. H S O 4 . 2 H 2 0 [15]. To an ice-cold aqueous solution of the bisulphate complex, was added concentrated HCIO 4 slowly with stirring when crystals of the perchlorate salt precipitated out. The crystals were filtered, washed several times with ice-cold absolute alcohol and finally with ether. The sample so obtained was recrystallized three times from water, before being used for kinetic studies. Found: Co, 29-05 per cent, SO~-2, 2%95 per cent: [Co(NH3)5SO4](CIO4) requires Co, 29-28 per cent, SO~ 2, 28.29 per cent. The molar absorptivity of 61.0 at 515 n m is in excellent agreement with the value of 61.0 reported by Buckingham et al.[13]. Sodium sulphate and sodium bisulphate solutions were prepared by mixing together requisite a m o u n t s of sodium hydroxide and sulphuric acid of known concentration. Reagent grade sodium perchlorate was used for adjusting ionic strength. Double distilled water was used for preparing all solutions.

drawn out at known intervals into an ice-cold a c e t o n e water mixture[16]. The chloride ion released by aquation was titrated potentiometricaUy with standard silver nitrate. (ii) Entry of sulphate anion. In order to examine the possibility of the formation of the sulphatopentaammine complex by the entry of the added sulphate anion, reaction mixtures identical with the above were simultaneously prepared and thermostatted. Aliquots were withdrawn from this reaction mixture at known intervals into well-stoppered test tubes (Pyrex) and quenched in an ice-bath. The optical densities of these solutions were measured with a Beckmann D U - 2 Spectrophotometer at appropriate wavelengths. Matched 1 cm quartz cells were used for such measurements and distilled water was used as the 100 per cent transmission standard.

RESULTS

T h e p s e u d o - u n i m o l e c u l a r rate c o n s t a n t s for t h e release o f c h l o r i d e i o n f r o m t h e c o m p l e x were calculated f r o m t h e s l o p e s o f t h e p l o t s o f log(vo~ - vt) vs t i m e t, w h e r e vt is t h e v o l u m e o f silver n i t r a t e c o n s u m e d after t i m e interval o f t, voo b e i n g t h e v o l u m e r e q u i r e d for c o m p l e t e reaction, v~ w a s i n v a r i a b l y c o m p u t e d from the weight of the complex taken. The results o b t a i n e d are collected in T a b l e 1. T h e acid h y d r o l y s i s r a t e c o n s t a n t s o f t h e f r e e c o m p l e x i o n h a v e b e e n d e t e r m i n e d at 40 °, 45 °, 50 ° a n d 55°C. T h e i n t e r a c t i o n o f t h e free c o m p l e x i o n w i t h t h e a d d e d s u l p h a t e a n i o n w o u l d r e s u l t in t h e f o r m a t i o n of o u t e r s p h e r e c o m p l e x e s o r i o n - p a i r s o f t h e type,

Kinetic experiments

[ ( C o ( N H 3)5C1) 2 +,HSO~" J ÷

(i) Acid hydrolysis of the complex. For aquation rate studies, requisite volumes of all reagent solutions excepting the complex were mixed in the reaction vessel and thermostatted. After thermal equilibrium was attained, a weighed a m o u n t of the complex was transferred into the reaction vessel. The solid dissolved rapidly on shaking. The mixture was made up to a certain volume with distilled water, equilibriated at the same temperature. Fiveml aliquots were

and [ C o ( N H 3 ) 5CI) 2 +,SO 2 - ]. T h e o b s e r v e d rate c o n s t a n t in t h e p r e s e n c e o f a d d e d s u l p h a t e i o n w o u l d , therefore, be a c o m p o s i t e t e r m c o n s i s t i n g o f t h e rate c o n s t a n t o f t h e free i o n a n d t h e rate a n d e q u i l i b r i u m c o n s t a n t s o f t h e ion-pairs.

Table 1. Acid hydrolysis of [ C o ( N H 3 ) 5 CI] ( C 1 0 4 ) 2 in aqueous solution in the presence of sulphate, (SO 2-) ion. Concentration of the complex = 5 x 10-3M. Ionic strength = 0.3 M [HCIO4] ( x 10a M) 10

10 10 10 10 ---'--

pH - -

2-12 2-22 2-32 2'50 -----

[HSO2] ( x 103 M)

[SO 2-] ( x 10 a M)

40 °,

---

-5(Na2SO4) 10(Na2SO4)

---

20(Na2SO4) 50(Na2SO4)

0.1(NaHSO4) 0.2(NaHSO4) 0.4(NaHSO4) 0-6(NaHSO4)

10(Na2SO4) 20(Na2SO4)

---

30(Na2SO4)

--

m

5(Na2SO4)

1.02 1.09 1.17 1.33 1.63 --

kobs x 105/see * 45* 50 ° 55°C 1.87 ----2.10 2.30 2.59 2.80

3.36 ----3.57 3.75 4-03 --

5.98 ----6-33 6-63 7.12 --

40 °" . . 2.24 2.26 2.25 2.25 (32-5) ----

ki~ x 10S/se9 45 ° 50 ° . ----4-04 4.05 4.04 4.04 (28.8)

55oc

. ----5-52 5-53 5.52 -(26-1)

----9.89 9.89 9.90 -(23.8)

* The observed rate constants are accurate up to + 1 per cent. The K~p values are given in the parentheses at the end of the system under the k~p readings.

Ion-pairing effects on chloropentaammine Co(Ill) ion

3795

Applying the treatment of Wyatt and Davies[9] to the system under investigation and assuming that the free complex ion, (Cp 2 +); and the bisulphate ion-pair of the complex, CpHSO~ react at the same rate[4, 6, 7], the rate expression at constant ionic strength will have the form,

The concentration of free sulphate ion[SO2-]s was obtained in the following manner.

-dc Rate = - = kobsXC= kl[Cp 2*] + kip[CpSO4] dt

where

(1)

where [Cp 2 +] = [Cp 2+]f + [CpnSO~] c = total initial concentration of the complex k~ = acid hydrolysis rate constant of the free complex ion [Cp2+]; = concentration of the free complex ion [CpSO4] = concentration of the ion-pair of the free complex ion with the sulphate ion kip = acid hydrolysis rate constant of the ionpair CpSO4.

K/~,

+

(Cp 2 +)¢+ (HSO~ b ~ (CpHSO,)

= [so,~-3; + EHSOZ]s + [CpSO,3 + [CpHSO~] + [NaSO,~]

(9)

[SO]-]T = total initial sulphate ligand concentration [SO]-], = total initial concentration of sulphate ion [HSO~]~ = total initial concentration of bisulphate ion [HSO~]y = concentration of the free bisulphate ion. [NaSO~,] = concentration of the ion-pair between (Na+); and (SO]-);. Substituting the values of [CpSO4] and [CpHSO~'] from Eqn (2) and (3) respectively, and that of [NaSO~] from the equation, (Na+); + (SO]-);

Representing the ion-pair equilibrium as (Cp2+)f + (SO2-)f ~ (CpSO4)

[S02-]T = [SO4Z-], + [HSO~],

KAs ~

(NaSO~).

(10)

(2)

(Where Ka, is the association constant of the reaction between (Na+); and (SOl-);), in Eqn (9), one gets

(3)

[SO]-]T = [SO]-]; + [HSO#]; + KI.[Cp2+];[SO]-]; + K,.[CpZ +];[HSO~];

the total initial concentration of the complex c is given by

c = [Cp~+]; + [CpSO4] + [CpHSO;] = [Cp2+];(1 • + Kip[soE-]y + Kip'[HSOg]y).

(4) (5)

+ Ka~[Na+];[SO]-];.

(11)

Assuming [Na+]T = [Na+]i and utilizing the equilibrium (HSO~); ~ (H+)f + (SO2-);

(12)

(where Kd is the second dissociation constant of H2SO,), Eqn (1 l) takes the form [so]-]¢ =

[H+]y

1+ ~

[so~-3T , 2+ [H+]I + Kip[Cp2+]f + Kip[C p ] f - - ~ - -~ Ka~[Na+]T

Substituting for c, [CpSO4], and [CpHSO~] in Eqn (1) and rearranging, one gets kob~ k,(1 + K,p'EHSO;]f) + k,pK,p[SO42-]i = 1 + K,v[SO2-3;+ K,v'[HSO;]I " (6) Assuming Kip' = 21171 and in the range of bisulphate ion concentration 1.0 × 10 -4 M to 4 × 10 -z M (as in the present work), 1 + K~p'[HSOF~]I has values ranging from 1-0002 to 1.08. Therefore, equating (1 + Kip' [HSO;, ]I) to 1, Eqn (6) approximates to

kobs

kl + kipKip[SO2- ]f = - 1 + Kip[SOl-I;

kob, -- k,

1

[H+]s

Kd

[HSO2], [SO]-],

In Eqn (13), 1 + ~[H +3~ + Kip[Cp2+]; + KAs[Na+]T >>Kiff[Cp2+]f [H+]f

Kd'

1

k,p - k~ + (kip - kOK,p[S02-];"

For the system under investigation, at 40°C, [H+]y was measured. Kd being known[18], the value of [l-I+]flKd could be computed. At 45°, 50°, and 55°C, the ratio of the [HSO~-]t/[SO]-]t was maintained at ~ , when it is reasonable to assume that,

(7)

Upon rearrangement, Eqn (7) takes the form 1

(13)

(8)

since the value of Kip[19 ] and K,r'[17 ] are of the order of 30 and 2 respectively. For computation of [SO 2-];,

3796

N. C. NAIKand R. K. NANDA

therefore, the Kip'[Cp2 +]y[I-I+ ] f l K d term is neglected. Eqn (13) then simplifies to

[so~-]f =

[ s o ,~-Jr ['I"I + If 1+ ~ + Kip[Cp2+]i + KA~[Na+]T (14)

Substituting for [Cp2+]f from Eqn (5) in Eqn (14), [SO~-]: is given by

[so,~-]~ =

According to the scheme I, release of chloride ion is visualized to take place through three possible pathways (i) by the aquation of the free complex ion designated by kx path, (ii) by the aquation of the ionpair (A ~ B), designated by kit,' path and (iii) by substitution of chloride by sulphate in the chlorosulphate ion-pair (A) which is designated as k°,t,) path. The ion-pair rate constant, kit,, is, therefore, a combination of kit,' and k°.t,) and kobs is the rate constant as defined by Eqn (1).

[so2-]~ 1 + [H+Jf ~-- +

Kicc + Ka,[Na+]r. 1 + Kip[SO2-]y + K;t,[HSO;]:

(15)

Reasoning as before, Eqn (15) approximates to

[so~-]r =

Es°~-]r

(16)

[H+]f Kipc 1+ ~ + 1 + K,p[SO2-]y t- Ka,[Na+]r

The values of KA~were calculated using published data [17, 20]. Kit,, [SO4Z-]yand kit, were then computed from Eqns (8) and (16) by an iterative procedure using the IBM 1130 computer as follows. Since k t and kob, were measured, plot of 1/(kobs -- kt) vs I/[SO24-]r (Eqn (8)), in the first approximation, gave l / ( k i t, - kt) as the intercept and 1/[(kit, - kOKit,] as the slope. Kip was calculated from these two values. Putting this value of Kit, in Eqn (16), ESO2-]£ could be computed. A plot of I/(kob~ -- k0 vs 1/[SO~-]:, thus gave better values for the slope and the intercept, hence better value of Kit,. Such an iterative operation was carried out till constant values of Kit,, [SO2-]: and kit, were obtained. The Kit, and kit, values have been collected in Table 1. P r o d u c t distribution

The sulphate ion-pair of Co(NH3)sCI 2+ may be visualized to react in two ways leading to the formation of [Co(NHa)sH2 O3+, SO 2-] and Co(NH3)sSO 2. The latter product again may be visualized to result by anation of the aquo-sulphate ion-pair. The various path-ways may be delineated as below:

Evaluation of k°.p) was attempted spectrophotometrically as described below: Since the anation rate constant, k2(k2 --- k2.p) x K~) is known[21], and kobs was experimentally determined, evaluation of k°.p) was possible. Dr, the optical density of a solution containing the Co(NH3)sC12+ and SO~- ions at 510 nm and at any time t may be expressed as. D t = ExA t + E2B t + E3C t

(19)

where A, = concentration of Co(NH3)sC12+ at time t, Bt concentration of Co(NHa)sH2 O3+ at time t,. C, = concentration of C0(NH3)5SO2 at time t, and Ex, E2 and E 3 are the molar absorptivities of the chIoro-, aquo-, and sulphatopentaammine cobalt(III) complexes respectively at a pathlength of 1 cm. According to the reaction scheme I, A, B, C at any time t are given by =

At = Ao e -k°b't

B, =

A°k°b~(e-k

.... -- e - k : )

(20) (21)

k2 - kobs

Scheme I

(17) (18)

Co(NH3)5CI 2+ + H 2 0 ~ Co(NHa)sH2 O3+ + C1Co(NH3)sC12+ + SO~- ~-~ ECo(NH3)sCI,SO4] /

(A)

~o

kaq

Co(NHs)sH20a+ + SO 2- ~ [Co(NH3)sH20,SO,] + + C I - ~2(ip) Co(:NHa)sSO 2 + C1-

(B)

(c)

3797

Ion-pairing effects on chloropentaammine Co(lll) ion

Substituting the values o f A . B t a n d C, f r o m E q n s (20-22) respectively in E q n (19) one gets,

and

k2 Ao kobs [e -k2t

e -k°b~t) ,

ka(it')A° e-k°b't kobs

A°k3(~J')+ Ao (22) ~-

• - e -k2t) D t = E1Ao e -k°b't -F -E2Aokob - - - - - ( e s -x°b't k 2 - kob~

kobs

[-Ao kobs e -k2t

w h e r e Ao refers to the c o n c e n t r a t i o n o f C o ( N H 3 ) s C I 2+ at zero time, a n d 0

2-

katip)Kip[SO4 ]f 1 -q- Kil,[SO2-]f"

ka(iP)

+ E3[

k 2 Ao e -k°b'~ - kob,

k3(ip)AOkobse-kob~+ AOkobsk3(ip)+ Ao1 .

(23)

Table 2. Optical densities of solution mixtures containing Co(NH3)sC12 + and SO 2- ions at different times, and at different sulphate ion concentrations: [Cqmplex] = 5 x 10- a M at 40°C and 1 x 10 -2 M at 45 °, 50 ° and 55°C Temp. (°C.)

[HCIO4] ( x 10aM)

[HSO4] ( × 10 a M)

[SOl-] ( x 10a M)

Time (sec) x 10 -3

1

2

3

4

5

40

10

0.00 7.20 14.40 18.00 21-60 25.20 0.00 7.20 12.60 18.00 21.60 25.20 0.00 3-60 7.20 10-80 14.40 0.00 3-60 9.00 12.60 18-00 21.60

50 (Na2SO4)

45

1.20 (NaHSO4)

60 (Na2SO4)

50

0-80 (NaHSO4)

40 (Na2SO4)

55

0-80 (NaHSO4)

40 (NaESO4)

Optical density (o.d.) Exp.* Calcd 6 7 0.225 0.225 0-227 0.230 0.230 0.232 0.450 0.450 0.460 0.462 0-465 0.470 0.450 0.450 0,455 0.460 0.465 0-450 0.457 0.470 0-482 0.500 0.510

0.225 0.227 0.228 0.229 0.230 0.231 0.450 0.453 0-457 0.463 0-466 0.470 0-450 0.452 0.456 0.460 0-464 0-450 0.457 0-472 0.482 0.494 0-510

* The experimental optical density is accurate up to ± 0.001 for 40°C and + 0.002 for 45 °, 50 ° and 55°C. Table 3. Rate constants for the kip, k2, k3,p) and k°,p) paths of reaction Temp. (°C) [SO ] - ] x 102 M --

--

Pathway

40

45

50

55

kip x 105/sec k 2 x 104/sec k3llp) × 106/sec

2.25 +__0.01 1.12

4.04 +__0-01 2.10

5.52 + 0.01 3-85

9.89 + 0-01 6-95

0.70 _ 0.09 0-94 _+ 0.09 -1.25 ___0-09 -1.36 +_ 0.04

1.03 _+ 0.11 1.14 _+ 0.12 1.45 _ 0-12 -1.48 + 0.11 1.81 _+ 0.07

1.26 + 0.16 1.88 + 0-21 2-22 + 0.21 --3.23 +_ 0.32

3.16 +_ 0-17 4.45 _ 0-17 5.30 +_ 0-18 --5.88 _ 0.13

1.00(Na2SOa) 2.00(Na2SO4) 4~00(Na2SO4) 5.00(Na2SO4) 6.00(Na2SO4) k°~ip) x 106/sec

(24)

Sulphate

Free ion

System

23.4 _ 0.2

AH t {kcal/ mole)

- 6.7 + 0.6

AS* e.u.

kl path

25.5

AG; (kcal/ mole at 40°C)

19.5 + 0.04

AH; (kcal/ mole)

- 17.5 + 0.13

AS* e.u.

kip path

24-98 + 0-02

AG t (kcal/ mole at 40°C)

22.6 + 1.2

AH* (kcal/ mole)

- 13.3 + 3.9

AS* e.u.

o ) path katie

26.8

AG* (kcal/ mole, at 40°C)

r

-4-1 + 0.00

AH (kcal/ mole)

-6.2 + 0.00

AS e.u.

-2.2

AG (kcal/ mole at 40°C)

Ion-association process

Table 4. Activation parameters for the kl, k~p and k°(i~) p a t h s of reactions and t h e r m o d y n a m i c parameters for the ion-association process

7`

7` _>

©

7'

Ion-pairing effects on chloropentaammine Co(Ill) ion E 1, E 2 and E 3 were determined experimentally at 510 nm k2, at the reaction temperature was computed from the data of Taube et a/.[21] using an assumed value of 25 kcal/mole as the activation energy. Ao and kob,, being known, k3tip~was calculated by using Eqn (24) with the help of an one parameter search programme (parameter being ka,p~ which minimized

3799

and attains almost a constant value which is reflected in the levelling off of the rate profile. The activation parameters for the kip path of reaction are almost of the same magnitude as for oxalate, malonate, succinate, maleate and phthalate[22]. If a comparison is attempted in the light of the interpretation advanced by Monk et al.[4], the ratio of the i=n ion-pair rate constant to the free ion rate constant at [ ( Z - D,i)/a(Da)] 2, where Z stands for the right 40°, 45 °, 50° and 55°C is 2.2, 2-16, 1.64 and 1.48 respeci=l hand side of Eqn (24) and a(D,i) is the error in D,. A set tively. The decrease of this ratio with increase in of representative experimental and calculated optical temperature can be attributed to the lower free energy densities for solutions containing the complex and of activation of the ion-pair compared to that of the free different concentrations of sulphate ion at 400, 45 o, 50° ion. The greater reactivity of the ion-pair may be ascribed and 55°C are recorded in Table 2. Utilizing the values of Kip, k3(it,) and [SO,Z-]I, to a process which is predominantly electrostatic in nature. The formation of ion-pairs results in reduction k°op~ was calculated from Eqn (23). The k2, k3tip~and k°0p~values have been recorded in of the overall charge of the complex and thereby favours expulsion of the chloride ion by a dissociative Table 3. mechanism. Such a process would result in a transition Computation of activation and thermodynamic para- state with considerable degree of charge separation and hence loss of activation entropy. The entropy of meters activation of the system under investigation is found The enthalpy and entropy of activation, AH: and to be negative and in accordance with Tobe's observaAS: for the reaction paths, k~pand k°t~p~were calculated tions[23] may be taken to be diagnostic of a disby fitting the rate constants to Eyring's equation, sociative mechanism for the aquation of the ion-pair with a square-pyramidal transition state. T e- AH~;/RTeaS*m, The observed rate constant is a measure of the k = R ~-~ chloride ion released upon aquation of the free complex with the help of a least square programme, which ion and from the ion-pair, as shown in the reaction scheme I. According to this scheme, chloride release minimized the function F as defined by seems to be effected in three possible ways: (i) Aquation of the chloropentaammine cobalt(III) F = ~ Nh e e - ki o'(ki)2 (25) ion leading to the formation of the aquopentaammine where a(ki) is the error in k~. The approximate values of cobalt(III) ion (kl) by a pseudo-unimolecular process. (ii) Aquation of the chloropentaammine c0balt(III) AHt and ASt were chosen from log k vs 1/T plot. The sulphate ion-pair, leading to the formation of the activation parameters have been recorded in Table 4. The enthalpy, free energy and entropy changes of the aquopentaammine cobalt(III)--sulphate ion-pair by a ion-association process, involving Co(NHa)sC12÷ ion pseudo-unimolecular process (k'ip). (iii) The third process is assumed to involve an interand S O l - ion were evaluated from the association change between the innersphere chloride ion and the constants of the ion-pair at different temperatures. associating sulphate ion of the ion-pair. The product The data have been collected in Table 4. complex is sulphato pentaammine cobalt(Ill) ion, [Co0NH3)sSO4] + as represented by the k°,p~ path, in the reaction scheme. The unambiguous interpretation of the mechanism DISCUSSION of such an interchange reaction has not yet been Aquation rate constants of the chloropentaammine advanced from kinetic data. However, this may be cobalt(Ill) ion, measured at 40°, 45°, 50° and 55°C explained in the light of the suggestion put forth by (Table 1) and the activation parameters (Table 3) are Eigen[24] for the anation reaction. in fair agreement with previously reported values[2(a)]. free ions ~- ion-pair The results clearly reveal that in the presence of sulphate, the rate of aquation of the complex is sigion-pair ~ innersphere complex. nificantly accelerated. The observed rate constants increase with increase in the sulphate ion concentra- Such a mechanism is believed to operate in the present tion at a particular temperature, ultimately becoming investigation. Besides the katip~path, the formation of the sulphato asymptotic to the [anion] axis at high anion concentration. The rate acceleration is ascribed to the presence pentaammine cobalt(Ill) ion may be visualized to of sulphate ion-pair, which is more labile than the free occur through an alternative pathway, namely the complex ion. With the increase in the concentration of anation of the aquopentaammine cobalt(III)-sulphate the anion, the concentration of the ion-pair increases ion pair (k2).

3800

N. C. NAIKand R. K. NANDA

The product analysis experiment was carried out to examine the extent to which the sulphatopentaammine complex is formed by way of interchange reaction of o ) the ion-pairs as represented by the k2tip~ and k3tiv path-ways. The plots of the optical density (Table 2) vs time invariably passed through a m i n i m u m which became shallower with increase in temperature and sulphate ion concentration. At 510nm, the molar absorptivities of sulphatopentaammine cobalt(Ill) is 60 and those of chloropentaammine cobalt(Ill) and aquopentaammine cobalt(Ill) are 45 and 43 respectively. The minima in the O.D. vs time plots, therefore, lead us to believe that the formation of sulphato pentaammine cobalt(Ill) predominantly takes place through the anation of the intermediate aquo sulphate ion-pair (k2 path). This is further substantiated by the excellent agreement between the experimental O.D. and the O.D. calculated on the basis of such a reaction scheme. The extent of chloride release from chloropenta0 ammine cobalt(iII) complex through ka,i~ path is limited to about 10 per cent in the maximum compared i 0 to that of the k~, kip and k3tiv ~pathways taken together. Hence it is reasonable to assume that the chloride release takes place predominantly through the formation of the aquo complex. The rate constants for the formation of the sulphatopentaammine cobalt(Ill) complex by interchange reaction of the chloro-sulphate ion-pair and by anation of the aquo-sulphate ion-pair at 25°C were calculated to be 2.08 x 10- 7 sec- ~and 1.3 x 10 -6 sec-1 respectively. From these values it is found that only 16 per cent of the sulphato-complex is formed by the ka,p~ path compared to the k2,p~ path. The energy of activation of the k3,p~ 0 path compares well with the same for the k~ path. The activation entropies for both the paths are negative implying a dissociative mechanism. The ion-association equilibrium for the association of chloropentaammine cobalt(Ill) cation with sulphate ion may now be considered. The free energy, enthalpy and entropy changes of the association process have been calculated to be - 2-2 kcal/mole, - 4.1 + 0.00 kcal/ mole and - 6 . 2 + 0.00 e.u. respectively at 40°C. The negative enthalpy change indicates an exothermic association. Acknowledgements--The authors are thankful to Dr. B. Deo, Professor of Physics, Utkal University and Sri M. K.

Parida for computational work. Thanks are also due to Sri A. C. Dash for many helpful discussions.

REFERENCES

1. F. Basolo and R. G. Pearson, Mechanisms oflnorganic Reactions. 2nd Edn. p. 158. Wiley, New York (1967). 2. D. R. Stranks, Reaction rates of transition metal complexes in Modem Coordination Chemistry (a) p. 129, (Edited by J. Lewis and R. G. Wilkins). Interscience, New York (1960). 3. A. B. Lamb and J. W. Marsden, J. Am. chem. Soc. 33, 1882 (1911). 4. S. H. Laurie and C. B. Monk, J. chem. Soc. 720 (1965). 5. V. D. Panasyuk and V. A. Golub, Russ. J. inorg. Chem. 2, 233 (1969). 6. F. J. Garrick, Trans. Faraday Soc. 33, 486 (1937). 7. F. J. Garrick, Trans. Faraday Soc. 34, 1088 (1938). 8. T. P. Jones, W. E. Harris and W. J. Wallace, Can. J. Chem. 39, 237 (1961). 9. P. A. H. Wyatt and C. W. Davies, Trans. Faraday Soc. 45, 778 (1949). 10. J. B. Walker and C. B. Monk, J. chem. Soc. (A), 1372 (1966). 11. M. B. M. Campbell, M. R. Wendt and C. B. Monk, J. chem. Soc. Dalton 16, 1714 (1972). 12. S. C. Chan, K. Y. Hui, J. Miller and W. S. Tsong, J. chem. Soc. 3207 (1965). 13. D. A. Buckingham, I. I. Olsen, A. M. Sargeson and H. Satrapa, Inorg. Chem. 6, 1028 (1967).. 14. R. C. Splinter, S. J. Harris and R. S. Tobias, Inorg. Chem. 7, 898 (1968). 15. S. M. Jorgensen, J. prakt. Chem. 2, 268 (1885). 16. D. D. Brown and C. K. Ingold, J. chem. Soc., 2680 (1953). 17. D. W. Archer, D. A. East and C. B. Monk, J. Chem. Soc. 720 (1965). The K~p' value at zero ionic strength was corrected for the ionic strength of the present investigation. 18. W. L. Marshal/and E. V. Jones, J. phys. Chem. 70, 4028 (1966). 19. Present work. 20. Reed M. Izatt, Delbert Eatough, James J. Christensen and Calvin H. Bartholomew, J. chem. Soc. 1, 45 (1969). 21. A. Haim and H. Taube, lnorg. Chem. 2, 1199 (1963). 22. Unpublished data. 23. M. L. Tobe, lnorg. Chem. 6, 1261 (1968). 24. M. Eigen, Plenary Lecture Proc. Seventh lnt. Conf. on Coordination Chemistry, p. 97. Stockholm, 1962, Butterworths, London (1963).