June20, 1963
HYDROLYSIS O F cHLOROAMMINERHODIUM(III) COMPLEXES
1741
57.4 (106+)+( 7 8 + )+ 28 from ring-labeled species. Observed retentions of 6.5% from the a-isomer and about 100% from the to arise by loss of CO from the parent ion. The ortho and meta are consistent with either or both of the above paths as the source of most of the C~HE.+ many similarities in the spectra of tropone and benzene -especially the profiles at m/e 50 to 52 and 73 to 78yield.
T h a t the label is retained in even a small part of the CsH5+ yield from benzaldehyde-a-d implies the existence of y& another contributing path. A metastable peak 76.0
( 7 8 + )--f (77+)
+1
in the spectrum of unlabeled benzaldehyde, not detected in the spectrum shown in Table I but clearly visible in the one measured on the high-mass instrument, identifies C6Hs+ as another C&S+ precursor. If the six hydrogen atoms in C&E.D+ lose positional identity, 83~&neglecting possible isotope effectsof the C&S+ ions arising from this precursor should retain the label. Comparison with Tropone Comparison of the spectra of benzaldehyde and tropone supports the view that most fragment ions from the former arise by paths not involving ring expansion. The spectra differ markedly-a rather different state of affairs from that found with toluene, cycloheptatriene and the four other C7Hs isomers, the mass spectra of which have been reported.’ Loss of a hydrogen atom to form C7H60+ contributes but little t o the spectrum of tropone. The most abundant ion is C!~HB+,shown by the metastable peak
[CONTRIBUTION FROM
THE
suggest that most of the smaller fragment ions formed from tropone stem from a C&Ie+ intermediate and, indeed, that common states of C H e + are involved in decomposition of tropone and benzene. The reaction apparently occurring in tropone-two C-CO bonds rupturing and a new bond forming across the gap from which CO is lost-seems to occur also in benztropone and in many quinones and aromatic ketones in which the carbonyl group is incorporated in a ring,26~30,31 but little or no loss of CO occurs from the parent ions of aliphatic aldehydes.l9 Loss of CO from benzaldehyde may arise from a ringexpanded, tropone-like intermediate. The driving force for other decomposition paths in benzaldehyde is evidently centered in the carbonyl group; ring expansion in C6He+formation would parallel the behavior of other benzene derivatives and thus reflect the influence of the benzene ring. (30) J. H . Beynon and A. E. Williams, Agpl. Specfry., 14, 156 (1960). (31) In cycloalkanones. in contrast, loss of CO is slight. High-resolution measurements show that the ions of parent mass less 28 units in the spectra of these compounds arise chiefly by loss of CzH4 [J. H. Beynon, R . A. Saunders and A. E. Williams, A p p l . S p e c f r y . , 14, 95 (1960)], paralleling the loss of CzH, from cycloalkanes [S. Meyerson, T. D. Nevitt and P. N . Rylander, in “Advances in Mass Spectrometry,” Vol. 2, R . M. Elliott, Ed. Pergamon Press, New York, N. Y., 1963, p. 3131.
CHEMICAL LABORATORIES OF NORTHWESTERN UNIVERSITY, EVANSTON, ILL.]
Kinetics and Mechanisms of the Hydrolysis of Some Chlorosinerhodium(II1) Complexes1r2 BY S. A. JOHN SON,^ F. BASOLOAND K. G. PEARSON RECEIVED FEBRUARY 20, 1963 The rates of hydrolysis of certain chloroamminerhcdium(111) complexes were investigated as a function of steric effects, the extent of chelation and other variables. It was observed that various nucleophiles have no effect on the rate of chloride ion release from these systems. These results are similar to the behavior of analogous cobalt(II1) systems. Striking differences were also found: ( 1 ) the rate of reaction of a rhodium(II1) complex is insensitive to the charge on the complex, (2) alkali has little or no effect on the rate of hydrolysis and (3) reactions of rhodium(111) complexes occur with almost complete retention of configuration. These results are discussed in terms of probable mechanisins for reactions of rhodium(111) complexes.
Introduction The literature contains numerous descriptions of the reactions of the ammine complexes of cobalt(II1). The diacidotetraammines, [CoA4X2] +, and the acidopentaammines, [CoASX]+, have been of particular i n t e r e ~ t . ~Extensive studies of the rates of hydrolysis of the complexes [ C O A ~ X Zand ] + [CoA6XI2+have been made as a function of a large number of variable^.^ I t is, therefore, somewhat surprising that so little information has appeared on the analogous complexes of the second and third row transition element members of the cobalt triad, rhodium and iridium. I t would certainly be of interest to ascertain the effect of increasing size and effective nuclear charge of the central metal atom upon the reactions of the various complex (1) This research was supported in part by the Air Force Office of Scientific Research of the Air Research and Development Command, under contract No. AF 49(038)315, and in part by the U.S. Atomic Energy Commission under contract AT(ll-1)-1087. (2) Taken in part from the Ph.1). Thesis o f S. A. Johnson, Northwestern University, lYfil, (3) National institutes of Health Fellow, l Y f i & l Y C i l . (4) In these general formulas, Aa and Aa indicate coordination to cobalt of four and five amine nitrogens, respectively; X represents a mononegative anion such as CI-, B r - , I - , C N S - , Na-, Nor or OH-. (R) F Basolo and K . G. Pearson, “Mechanisms of Inorganic Reactions,” John Wiley & Sons, Inc., New York, N . V., 1958, Chapter 3.
ions. The only quantitative kinetic studies of the reactions of acidoamminerhodium (I1I) and -iridium(111) complexes are those reported by Lamb6 for [Rh(NH3)6C1]2+and [Rh(NH3)sBr]2+ and by Lamb and Fairhall’ for [Ir(NH3)5X]2+, X = C1-, Br-, I- and NOS-. The dearth of information available on the reactions of diacidotetraamminerhodium (I 11) complexes may be logically ascribed to the fact that until very recentlys only a few isolated members of this group had been prepared. Any study of those compounds whose syntheses had been described would have precluded all but the sketchiest of comparisons to the members of the very prolific cobalt(II1) family; indeed no examples of either geometric or optical isomerism were known. The successful synthesis of many of the necessary diacidotetraamminerhodium(111) complexes now makes possible quantitative comparisons between these systems and the analogous cobalt(II1) compounds. This paper reports kinetic data on reactions of rhodium( I I I) systems and discusses these data in terms of the previously reported results on cobalt( I I I ) . A. B . Lamb, J . A m . C h c m . Soc., 61, I i Y Y (1Y8!4) (7) A . B. Lamb and 1,. T . Fairhall, i b i d . , 46, 378 (1928). ( 8 ) S. A. Johnson and F. Basolo, I n o r g . C h e w , 1, 925 (l!Jli2). (I;)
S. A. JOHNSON, F. BASOLO AND R. G. PEARSON
1742
Vol. 85
TABLE I RATESOF HYDROLYSIS IN ACIDICSOLUTION Complexa
Medium
T, OC.
k o b X lW,
RClequil
"ki" X 1W,sec-1
trans-[Rh(NH,)4C1z] 0.1 MHNOI 80 12.1 39 4.7 . I MHNOa 80 10.4 26 2.7 trans- [Rh(en)zClz] + . l M HNOa 80 85 60 51 cis-[Rh(enhCl2] + . l M HNOa 55 15 47 7.1 cis-[Rh(trien)Cl~]+ trans-[Rh(m-bn)zCl~] . l M "03 80 85 32 27 trans- [ Rh(dl-bn)~Ch]+ . I M HNOa 80 12 30 3.6 trans- [Rh(tetrameen)&L] H20 25 160 -50 -80 cis-[Rh(tetrameenhCl:] + Hz0 25 110 -50 -55 truns-[Rh(bipyhCl~]+ 0 . 1 M HNOa 80 ... 0 ... [Rh(NH3)6CIl *+ 0 . 1 M HNOa 80 6.2 69 4.3 trans-[Rh(czo4~clz]~H2O 80 1.6 100 1.6 Hz0 80 2.3 100 2.3 cis-[Rh(CzO,)sClz]3 en = ethylenediamine; trien = triethylenetetramine; m-bn = meso-butylenediamine; dl-bn = dl-butylenediamine; tetrameen = (C,C')-tetramethylethylenediamine; bipy = 2,2'-bipyridine. +
+
+
Experimental and Results A. Preparation of Materials.-trans-[Rh(NH~)&l~]NOa,cisand trans- [RH(enhClz]NOS, cis- [Rh(trien)Cl?]NOa, trans-[Rh( mbn)zClz]NOa, trans-[Rh(dl-bn)zClz]NO~, 625- and trans-[ah(tetrameen)~Clz]C1 and [Rh(NHs)Xl] (N03)2 were prepared and analyzed as described previously.* Dichlorobis-(2,2'-bipyridine)-rhodium(111) chloride was prepared by the method of Jaeger and van Dijkg by the reaction of an aqueous solution of RhCla.3HzO with 2,2'-bipyridine in ethanol. The product, a fluffy yellow crystalline substance, was washed with a little water, then with acetone to remove any unreacted bipyridine. The yield was approximately quantitative. Anal. Calcd. for [Rh(bipy)zC1~1C1-2Hz0: C, 43.07; H, 3.62; N, 10.05; Rh, 18.45. Found: C, 42.45; H, 3.31; N, 10.24; Rh, 18.60. Potassium cisdichlorobis-( oxa1ato)-rhodate(111) was prepared by the method of Delepine.lo The crude product was recrystallized from water-ethanol. The fluffy golden yellow powder was obtained in 50% yield. Anal. Calcd. for K ~ [ R ~ ( C Z O ~ ) Z C ~ Z ]C, . H 9.90; Z O : H, 0.44; C1, 14.6. Found: C, 10.13; H , 0.36; C1, 14.6. B. Determination of the Rates of Chloride Ion Release from Chloro Complexes.-The rates of chloride ion release in these reactions were determined by three different methods: (!) Titration of Chloride Ion.-Whenever possible, a hydrolysis reaction was followed by direct titration of the chloride ion as it was released. In general, sufficient complex to make its concentration -2.5 X 10-3 M was added to 50 ml. of HzO (containing the other reactants if any) which had been thermostated a t the desired temperature (f0.1")in a constant temperature bath. At appropriate intervals a 5.00-ml. aliquot was removed from the reaction mixture and cooled to room temperature or slightly below. The sample was then made acidic by the addition of 1 ml. of 6 M "01 and titrated with 0.01 M Hg(N0a)Z using sodium nitroprusside as the indicator .I1 Rate constants were determined graphically from the relationship
(C,
-
Co) Ct) where C ,, CO,and Ct represent the experimental concentration of chloride ion a t infinite time, zero time and time t , respectively. ( 2 ) Conductometric Method.-A reaction was. followed conductometrically when a significant change in the conductivity was expected during the course of the reaction. In general, the finely powdered complex (1-2.5 X 10-8 M ) and any other solid reactant were placed in one leg of a Y-shaped conductivity cell and 25 ml. of distilled water was pipetted into the other leg. The cell was then thermostated a t the desired reaction temperature (f0.1") in a constant temperature bath. When temperature equilibrium had been attained, the cell was inverted causing the water to run into the leg containing the complex which dissolved to begin the reaction. Measurements of the resistance of the solution were made at appropriate times using an industrial instruments conductivity bridge, model R C 16. Rate constants 2.303
k = - log t (C,
were evaluated graphically from equation 2 where Ro,R, and (9) F. M.Jaeger and J. A. van Dijk, 2. anorg. allgem. Chem., 227, 323 (1936). (10) M. Delepine, .4nales soc. espan. 5s. quim., 27, 487 (1929). (11) I. M. Kolthoff and E. B . Sandell, "Textbook of Quantitative Inorganic Analysis," The Macmillan Co., New York, N. Y.,1953, pp. 460, 547-549.
Rt are the measured resistances of the solution at zero time, infinite time and time t , respectively. was often possible t o (3) Spectrophotometric Method.-It find some wave length in the visible or ultraviolet region where the absorptions of the starting complex and the reaction product were sufficiently different that the reaction might be followed spectrophotometrically. Most of the reactions were followed at 80'. For these runs the solvent water containing any other reactants was first thermostated in the 80" (f0.1"),constant temperature bath and the complex (-2.5 X lo-' M )was then added t o it. At appropriate intervals a 3-ml. aliquot was removed from the reaction mixture and pipetted into an ice-cooled beaker. After the sample had been cooled t o room temperature or slightly below it was placed in a l-cm. quartz cell and its optical density was measured at the chosen wave length using a Beckman DU spectrophotometer. When the reaction could be followed at 55" or less, the cell compartment of the DU was thermostated at the desired temperature by circulating water around it from the constant temperature bath. A 3-ml, sample of the reaction mixture prepared as described above was placed in a l-cm. quartz cell in the thermostated cell compartment and periodic measurements were made of its optical density. Rate constants were calculated graphically from the relationship 2.303 A , - Ao k = - t (3) log A x t where A,, AO and At are the measured optical densities at infinite time, zero time and time t , respectively. Table I gives results of measurements of the rates of acid hydrolysis of the various complexes. Duplicate runs were made in every case and agreement was usually within 2-3%. First-order kinetics were found to be followed out to two or more half-lives. These solutions did not hydrolyze completely; less than half of the total coordinated chloride was generally found by titration at equilibrium. In such a case the observed rate constant calculated from eq. 1 is not a simple quantity but a composite of all four rate constants for the two reversible reactions [Rh&CLI
+
ki
+ H2O
__
[Rh&ClH%O]'+
kr kJ
[Rh&ClHiO]'+ f H2O
[Rh&(HzO)da+
+ C1-
(4)
+
(5)
C1-
k4
Furthermore, by analogy with other anation reactions, the reverse steps of 4 and 5 undoubtedly depend on the chloride ion concentration. A somewhat similar situation was met by Lamb' in his study of the acid hydrolysis of [Rh(NH3)sX]"+. In this example where only one halide ion can be released, Lamb showed that the observed constant is given by kob
=
ki
+ kz[X-I
(6)
where [X-1 is an average value of the halide concentration during a run. Further, k, could be calculated from a knowledge of the equilibrium concentration of halide ion, C,
June20, 1963
HYDROLYSIS OF kl
= (Cm/Co)koba
CHLOROAMMINERHODIUM(III) COMPLEXES
hubs
=
-k k o ~ - [ O H - l
~H&I
(8)
where k H 2 0 is the rate constant for acid hydrolysis and k o H - is a second-order rate constant for reaction with hydroxide ion. The relative values of k H n O and k O H - determine the apparent order with respect to hydroxide ion in different p H ranges. The data of Table I1 allow estimates of the constants in eq. 8 to be made as described in the following section. The reactions of dilute solutions of trans- [Rh(en)2cl~]+ with a variety of possible nucleophilic reagents were also studied. The reactions were found to go to completion. In every case simple first-order kinetics were observed. The rate constants are taken from the first-order slopes and are given in Table 111. The rates are seen to be essentially independent of the nature and concentration of the reagent X. The observed product, however, is that in which X is in the coordination sphere. The products were assessed by spectra and by isolation and analysis.8 In the case of "3, the preparative run gave trans- [Rh(en)zNHoC1I2+. However, this was for a short period of time (10 minutes a t reflux). The kinetics run showed that after longer periods of time (several hours a t 80') both chlorides were released. In the case of thiourea, no product could be identified by preparation, but in the (12) R. G. Pearson and F. Basolo, J
TABLE I1 RATESOF HYDROLYSIS IN BASICSOLUTION k X Id, T,
(7)
While such a procedure cannot be valid in cases where two replaceable chlorides are present, nevertheless "kl" was calculated from eq. 7 and the values obtained are given in Table I. They will be compared later with rate constants for base hydrolysis. It may be noted that for complexes such as cisand trans- [ C ~ ( e n ) ~ C +,lthe ~ ] second chloride ion aquates a t a rate only one-hundredth of that of the first. This is usually considered partly an effect of the increased positive charge in [ C 0 ( e n ) ~ H ~ 0 C land ] ~ + partly as a result of dative s-bonding from chlorine to metal in the dichloro complex. l 2 Although we have no distinct data on complexes such as [RhA4H20C1]2+,the data in Table I show little effect of charge on the rates of acid hydrolysis. For this reason we assume that kl and k 3 are, if not equal, a t least very close, and kl and k4 also comparable. This view is supported by the first-order behavior observed. Whenever hydrolyses of [RhA4C12]+ were carried out in basic solution, 100% of the coordinated chloride was released a t equilibrium and found by titration. Again the reaction is a two-stage process and we have no direct information on the behavior of [RhAdOHC l ] + intermediates. However, i t was found that a t constant pH, good first-order kinetic plots were obtained over two or more half-lives. This can mean either that [RhA40HCl]+ reacts much faster than [RhA4C12]+,or that the two react a t very nearly the same rate. As indicated above, it seems more likely that the latter is true. In any case, the values of the rate constants given in Table I1 are simply the observed first-order constants with no corrections. The rates of base hydrolysis of [RhA4C12]+ were found to be independent of the concentration of hydroxide ion for all trans isomers for concentrations of base as high as 0.10 M. In the hydrolyses of cis[RhA4C12]+ and [Rh(NH3)&1I2+, the order with respect to hydroxide ion was found to vary between zero and one. For comparison, [CoA4C12]+hydrolyses are strictly first order in hydroxide ion above a p H of 8 or so. A t constant p H all of these systems may be described by the equation
A m Chem. S o c . , '78, 4878 (1956).
1743
Complex
OC.
9H"
sec. -1
80 8.5 8.9 80 8.9 9.9 8.6 80 12.7 70 1.5 trans- [Rh(en)~Clz]+ 8.9 80 4.2 8.9 11.6b 90 8.9 4.4 80 9.9 SO 5.1 12.7 55 6.3 cis-[Rh(en)~Cl~] 8.9 14 55 9.9 7.3 11.9 35 14 40 11.9 27c 45 11.9 9. I d 25 12.7 16 cis-[Rh(trien)Cl~] 55 8.9 52 55 9.9 19 11.9 25 1 . 9 x 102' 12.7 25 41 80 8.9 so 38 9.9 46 80 12.7 80 5.6 9.9 6.1 80 12.7 1.0 x 102 trans-[Rh(tetrameen)~CI~] 25 12.7 -7.5 0 12.7 cis- [Rh(tetrameen)~Ch]' Negligible 80 12.7 trans-[Rh(bipy)~Clr] 80 9.1 8.9 [ Rh( IL"3),C1] 26 80 9.9 Fast 80 12.7 Dec. 80 tr~ns-[Rh(Cz04)zCI~] 8.9 Dec. 80 ci~-[Rh(C?O1)2Clz] a8.9 a p H values are those measured with a Beckman p H meter, model G, of the following media a t 25': 8.9, 0.1 M &BoaNaZB407; 9.9, 0.1 A4 NaHC03-NazC03; 11.9, 0.1 M NaZHP04NaOH; 12.7, 0.1 M KOH. Corrections were not made for the E. = change in pH of a buffer with increasing temperature. 25.5 kcal. E, = 25 kcal. Second-order rate constant k = 2 X 10-3 M-1 set.-' a t 25'. e Second-order rate constant k = 3 X M - l set.-' a t 25".
h n s - [ R h ( NH3)4CIz]
+
+
+
+
'+
+
'-
"
kinetic run both chlorides appeared to be released and a product different from the hydroxy or aquo species was formed. In addition, qualitative studies of the rates of reaction of trans- [Rh(en)2C1~]+ with Br-, CNS-, CN-, s 2 0 3 * - , S032- and N3- were made. In no case was the rate of reaction significantly greater than those given in Table 111. TABLE I11 RATESOF THE REACTION trans-[Rh(en)tCl~] 2X -+ truns-[Rh(en)~X~]"+ f 2c1- FOR VARIOUS REAGENTS AT 80' +
k X lo5, X
sec. - 1
+
k X 105, X
sec. - 1
OH- (0.1 M ) 5.1,b4.0a I - (0.05 M ) 5. l C NOz- (0.1 M ) 4.2b 3 ~ 1 -(0.01 hi) 4.0 NOz- (0.05 M ) 4.2b Thiourea (0.1 M ) 4.9" I - (0.1 M ) 5.2O NHI (5.0 M ) 4.0° Extrapolated from data a t 78' obtained by U. Klabunde using chloride titration. By chloride titration. Spectrophotometric. Conductometric.
These results can be explained most readily by assuming that the rate-determining step in every case is the hydrolysis of the dichloro complex. This is followed by the rapid replacement of coordinated water by the group X. This suggests that the value of k in Table I11 is the true rate of acid hydrolysis, k H , O , in eq. 8 . The difference between this value of about 5 X set.-' and the value of k, equal to 2.7 X sec.-l given in Table I can be attributed to the inapplicability of eq. ?. This equation assumes, for example, that
S. A.
1744
JOHNSON,
F. BASOLOAND R. G. PEARSON
TABLE IV A COMPARISON OF THE HYDROLYSIS BEHAVIOR OF cis-
AND
Vol. 85
trans-[Rh(en)zClz]+ AND [Co(en)zClz]+ pH where
kazo, Complex
koa-, M - 1 sec.-l
sec. - 1
