ACIDHYDROLYSIS OF C~~-DICHLORODIAMMINEPLATINUM( 11)
June 5 , 1961
in explaining the invariance of the energy of activation from one halo complex to another. The reorientation of the solvent around the reacting complex must then be a real effect. The variable portion of this effect must be associated with the leaving group since it would appear to be the only variable within the system. If this is a correct assumption, then the differences in the entropy of activation between the complexes ought to roughly parallel the differences in the thermodynamic entropies of the reactions corresponding to I. Latimer and Jollyz5 have suggested a procedure for the estimation of the thermodynamic entropy of any reaction of this type. In general form this reaction is rewritten halogen (bound)
+ HzO(l)
---f
HzO(bound) f halide (aq)
(111)
Values that have been estimatedz6for the various entropy terms in 111 are 9.4 e.u. for bound water, 21 e.u. for free bromide, 16.7 for free water and 9 for bound bromide. These values lead to a calculated value for A S for reaction l of f4.7 e.u. A similar treatment applied to the hydrolysis of the chloro and iodo complexes leads to values of +O.S and 7.2 e.u., respectively. These values have (25) W. hI. Latimer and W. L. Jolly, J . A m . Chem. Soc., 76, 1548 (1953). %'. hl. Latimer, "Oxidation Potentials," 2nd Ed., Prentice(26) 1 Hall, Inc., S e w York, N. Y . , 1952.
[CONTRIBUTION FROM
THE
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approximately the same relationship to one another as do the entropies of activation for the hydrolysis reaction. The principal difference between the thermodynamic entropies of the reaction leading to the replacement of halogen by water in the complex is the entropy of formation of the various halide ions in solution ((21-, 15 e.u., Br-, 21 e . u . ; I- 27 e . ~ . ) .A~ correlation ~ between the entropy of activation for the hydrolysis reaction and the thermodynamic entropy of hydration of the ion seems to be reasonable on the basis of this data. The relative order and magnitudes of these values is not greatly different from that observed between the entropies of activation for the three complexes. Such considerations as these indicate very strongly that the solvent sheath is a direct participant in both the PH independent and the QH dependent reactions. Furthermore the similarity of the variation of the entropy terms in the two reactions which behave so differently to changes in the pH indicates that both have transition states which interact with the solvent sheath in the same way. There is however, no indication as to whether a single solvent molecule is specifically involved in the transition state. It is not possible then to assign these reactions on the basis of the observations reported in this paper to either of the S x l or S Nclassifications. ~ (27) W. M. Latimer, K S. Pitzer and C. M. Slansky, J . Chem P h y s , 7, 108 (1939).
INSTITUTE FOR ATOMIC RESEARCH ASD DEPARTMENT OF CHEMISTRY, IOIVASTATE USIVERSITY, AMES,IOWA]
cis-Dichlorodiammineplatinum(I1). Acid Hydrolysis and Isotopic Exchange of the Chloride Ligands1 BY JOHN W. REISHUSAND DONS.MARTIN,JR. RECEIVEDOCTOBER 21, 1960 The acid hydrolysis of cis-[Pt(NH3)~Cl~] has been studied a t 25 and 35" For the first acid hydrolysis G ~ X - [ P ~ ( N H ~ ) ZHz0 C I ~+ ] c k [ P t ( XHa)?Cl(HzO)] C1the equilibrium constant, K1, is 3.3 X mole/l. and the rate constant, k l , is 2.5 X 10-5 set.-' a t 25'. There is no significant direct exchange between the chloride ligands of cis-[Pt(KH,)2Cla]and C1-. For the second acid hydrolysis ~ i ~ - [ P t ( X H 3 ) z C l ( H z+0 ) ] HzO +cis-[Pt(SH3)z(HzO)z] f C1the equilibrium constant KS is 4 X mole/l., a t 25'. The exchange of chloride with cis-[Pt(r'jH3)~C1(H~0)] occurs a t a rate which is chloride-independent and is characterized by a first order rate constant, kz = 3.3 X set.-' a t 25'. I t appears likely t h a t this exchange also occurs by only an acid hydrolysis mechanism.
+
+
+
+
++
Introduction When the compound which was originally designated by Wernerza as cis- [Pt(NH3)2C12] is dissolved in HzO.the solution possesses the very low conductivity of a non-electrolyte. However, an increase in the conductivity of its solutions has been noted by a number of investigator^.^-^ Jensenj proposed that the increase in conductivity
of the solutions upon aging can be explained in terms of acid hydrolysis or aquation of the chloride ligands according to
(1) Contribution No. 924. Work was performed in t h e Ames Laboratory of t h e U. S. Atomic Energy Commission. (2) (a) A. Werner, Z . anopg. Chem.. 3, 267 (1893). (b) A. Werner and A. Miolati. Z . phypik. Chem., 12, 49 (1893). (3) H. D. K . Drew, P.1". Pinkard. W.Wardlaw and E . G. Cox, J . Chem. SOC.,988 (1932). (4) D. Banerjea, F. Basolo and R. G. Pearson. J . A m . Chem. Soc., 79, 4055 (1957). ( 5 ) I 1 are obtainable under certain conditions. It was found t h a t the rate of exchange of the cis-[Pt- high. Equilibrium constants, K1 and K 2 , were (NH&C12] fraction was relatively insensitive t o the rates calculated by the use of Eq. 8 from each combinaRZ or R". It was not possible to isolate the [Pt(NH&- tion of two titrations a t a given temperature. The
+ +
-
2460
if'. REISHVSAiYD
JOHN
s. AlrZRTIN, JR.
ir01. 83
I .o
0.9 0.8
07 0.6
:
.4
-
05
-.
8
0.4
2
0.3
,..-
3
-
.3
a
I
w $
EQUIL. T l T H t = 1 . 9 7 ~IO
1 0
0
20
40
60
80 100
200
I i
I
l
'
i
1
300
400
1 500
(
(
-4
-
I
~~~~
0 iT
.
600
700
- MIN. Fig. 1.-Titer of cis- [Pt(NHB)zCI?]solutions as. time, t c m p . = ?So,1 = 0.318. 0 , initial concentration = 0,0025 n i i i l c ~ l . 0, , initial concn. = 0.0050 mole/]. TIME
averages of these calculated constants are in Table I. The concentration of each of the complex species and the titer, calculated from the equilibrium constants, have been included for each experiment in Table I. The agreement of these calculated titers with the observed values attests t o the consistency of the equilibrium constants. It is noted t h a t the calculated concentration of the [Pt(KH~)Q(H?O ) ~was ] ++ never large and, therefore, the constant K 2 was not well defined. If for t h e 25' experiments, K 2was arbitrarily set equal to zero and K 1 was calculated from the experiment at a = 5.00 X rnole'l. and b = 0, the calculated mole '1. and b = 0 was titre for n = 2.30 X
