330
F. D. GRAZIANO AND G. M. HARRIS
R. A. P L A N E . - - F O I the racemization, as for the aquation of the hexammines of Cr(II1) and Co(III), the significant differencemay lie in the nature of the lowest spin-forbidden state. In the Cr(II1) doublet a de orbital can be available to enhance non-oxidation-reduction reactions which is not the case for the lowest Co(II1) triplet. Incidentally, for Cr(C204)8-4,Schlafer reports the lowest doublet, the spacing for which (relative to the lowest uartet) is intermediate between that of Cr(H20)6+3and 8r(NHa)s+3; so that I would have expected a measurable quantum yield.
Vol. 63
R. G. PEARsoN.-It is proposed that in CrSf a doublet state in which two electrons are paired in a single d orbital may be reached easily. This leaves a vacant lower d orbital which may allow for easier substitution reaction, according to the suggestion of H. Taube, by a bimolecular displacement path. However, in the crystal field theory, a vacant lower d orbital allows for easier substitution reaction by either a bimolecular or unimolecular path. This is because ligands in the plane of the vacant orbital can be rearranged without excessive loss of crystal field stabilization energy.
SUBSTITUTION REACTIONS OF OXALATO COMPLEX IONS. I. OXALATE EXCHANGE REACTIONS OF THE TRIS-OXALATO-COBALT(II1) AND CHROMIUM(II1) COMPLEX ANIONS BY F. D. GRAZIANO' AND G. M. HARRIS Department of Chemistry, Universit~05 Bufalo, Buffalo, N . Y . Received October 4 , 1068
A study has been made in aqueous solution of the reactions between the ions M(CQO&~ and carbon-14-labeled C20a", where M is Co+++or Cr+++. The exchange waa followed by precipitation and radioactive assay of the free oxalate as calcium oxalate monohydrate. Ligand exchange of the cobaltiate proceeds a t a rate negligible compared to that of thermal decomposition. The est,imated exchange half-time is 2130 hours at 50" and pH 8, compared to 6.5 hr. for the decomposition. It is thereby confirmed that the decomposition proceeds by a mechanism not involving appreciably reversible equilibration of C20r radicals or ions. The chromiate exhibits no sign of dark thermal decomposition during several weeks a t 75" and 4 < pH < 6, while ligand exchange shows half-times of less than a day under similar conditions. Outside the quoted acidity limits, decomposition becomes measurable and exchange is accelerated. The rate of exchange is practically independent of ( H + )in the 4 < pH < 6 range, but becomes first-order in (H+) a t lower pH. The reaction is fist order in (com lex) and partiall f i s t order in (oxalate). A mechanism consistent with cognate studies of the chromiate ion can be fittea t o the data. d o s e analogies with the corresponding rhodiate behavior are observed.
Introduction Reactions of the tris-oxalato complexes of the transition elements have been the subject of several kinetics investigations. Johnson2 was the first to attempt to correlate resistance toward ligand substitution with bond-type, utilizing as his criterion the retention of optical activity. Thus, the Al(III), Mn(II1) and Fe(II1) complexes were classed as ionic, since resolution was impossible, while the readily resolvable Co(II1) and Cr(II1) analogs were stated to be covalent. Long3 substantiated this classification in his pioneer isotopic exchange study, using carbon-1 1-labelled oxalate. Rapid and complete oxalate exchange occurred at room temperature with the non-resolvable complexes. The Co(111) and Cr(II1) complexes, however, underwent no observable exchange (at 50 and 35", respectively) within the short time intervals it was possible t o follow t.he reaction with carbon-11, in spite of the fact that racemization of these compounds is extensive under similar conditions. Clearly, a simple dissociative racemization mechanism is not operative. An analogous lack of relationship between exchange and racemization exists for the complex ion Coen2C03+, though in this case racemization is by far the slower p r o ~ e s s . ~ A study of cobalt electron-transfer exchange between the species C o ( C ~ 0 4 and ) ~ ~ CO(CZO~)~-~ has (1) Work done by F. D . Graziano as part of Ph.D. requirement of University of Buffalo, 1958. Complete report available from University Microfilms, Ann Arbor, Michigan. (2) C. H. Johnson, Trans. Faradav Soc., 28, 845 (1932); 31, 1612
(1935). (8) F. A. Long, J . A m . Chem. Soc., 61,570 (1939): 68, 1353 (1941). (41 J. 9. Holden and G. M. Harris, ibid., 77, 1934 (1955).
been made by Adamson and co-workers.6 The thermal decomposition of the former ion complicates the interpretation, but a slow exchange was observed. The mechanism proposed parallels closely that recently suggested for both thermal and photochemical decomposition reactions of the Co(II1) speciese and also provides a path for exchange of its oxalate. However, later work by Stranks and co-workers' supports an alternative interpretation which necessitates no appreciable back-exchange of oxalate ion into the Co(C2Or)a" group, notwithstanding the measureable electrontransfer process between the Co(I1) and Co(II1) oxalato complexes under the reaction conditions . _ employed. The only other studs made so far of the Cr(C204)8e/C204-exchange system is that of Carter, Odell and Llewellyn,* who found that H2Ols exchanged oxygens with the Cr(II1) complex much faster than exchange occurred with carbon-13labeled oxalate ion. Further, all twelve oxygen atoms of the complex underwent exchange, suggesting a "one-ended" dissociation of oxalate ion as the initial step in this process
+
Cr(C204)3g H20
C~(CZO~)~OCZO~.HZO"
This provides a common mechanism for racemiza( 6 ) A. W. Adamson, H. Ogata, J. Grossman and R. Newbury, J . Inorg. Nuclear Chem., 6 , 319 (1958). (6) T. B. Copestake and N. Uri, Proc. Roy. Soc. (London), A%'28, 252 (1955). (7) W. Schneider, D. R. Stranks and F. S. Dainton, private corn. munication from D. R. Stranks. (8) J . H. Carter, A. L. Odell and D. R. Llewellyn, private communication from J. H. Carter.
L
March, 1959
SUBSTITUTION REACTIONS OF OXALATOCOMPLEX IONS
33 1
tion, oxygen exchange and mono-aquation, each of which can be very rapid relative to oxalate exchange. The present investigation had two purposes: (a) t o confirm and supplement the findings of LongS relative to the oxalate exchange reaction of Co(CZO4)F and Cr(Cz04)3’;, making use of the much longer-lived isotope carbon-14, and (b) to carry out detailed kinetic studies of the exchanges, if possible, with a view toward elucidating their exact mechanisms. As will be seen, no measureable oxalate exchange takes place with the Co(II1) complex notwithstanding its rapid irreversible thermal decomposition, but the kinetics of the exchange reaction of the Cr(II1) species has been amenable to thorough study. Experimental
contrived to be in the neighborhood of 0.84 mg./cm.2. This corres onds to the maximum in the empirical selfabsorptionjelf-scattering correction curve under the experimental conditions, thus tending to minimize errors in the application of the correction.*6 I n absence of decomposition processes, self-buffering of the reaction solutions was provided by the oxalate-acid oxalate mixture. pH measurements were made after rapid cooling of the sample of reactant to room temperature, employing a Beckman Model G pH-meter Spectrophotometry was done on Beckman DU or DK-2 instruments. Half-times of exchange were determined from the conventional McKay plots, when applicable, and the results expressed in terms of the rate of exchange R in the usual manner,13 neglecting possible isotope effects.
ate decomposition mixture prepared for counting in the usual manner. Thin-end-window counting of solid samples was employed throughout, with corrections applied in the case of the pure calcium oxalate samples for background, self-absorption and self-scattering, as determined in separate experiments. Whenever possible, these samples were
(15) The curve has a rather flat maximum (see G. B. Cook and J. F. Duncan, “Modern Radiochemical Practice,” The Clarendon Press, Oxford, England, 1952, p. 236) so that the correction is nearly
Results and Discussion A. The Tris-oxalato-cobalt(JI1)Ion.-The thermal instability of this complex rendered it impossible to carry out conventional exchange studies with free oxalate ion. Preliminary experiments A. Preparation, Purification and Analysis of Compounds.-Potassium tris-oxalato cobaltiate was prepared based on spectrophotometric determinations according to standard procedures9J’J and purified by suc- showed that at 50” the decomposition proceeded at cessive recrystallizations from ice-water. Cobalt was de- the rate predicted by Copstake and Uri’s results,6 termined by the l-nitroso-2-naphthol method11 and oxalate by Sorensen’s technique.9 Calcd. for K ~ C O ( C ~ O ~ ) ~ . ~ ~and / ~ Hthat ~ O :it was independent of the concentration Co, 11.71; C Z O ~ 52.46. , Found: Co, 11.75 & 0.06; of added free oxalate in the acidity range of the czo4,52.51 f 0.03. present study (6 < pH < 9). The change of pH Potassium tris-oxalato chromiate was prepared by stand- with time of a typical reaction mixture ((complex) ard rnethods’o and purified in the same way as the cobaltiate. Chromium was determined by the conventional technique12 = 0.03505 M ; (oxalate) = 0.01460 M ) a t 50” was and oxalate as before. Calcd. for K3Cr(CzOc)3.3Hz0: Cr, followed. There was a drop in pH within the first 10.67; (2204, 54.19. Found: Cr, 10.92 f 0.10; CZO4, hour from 7.0 to 6.1, then a steady rise to 8.7 a t 54.81 f 0.10. complete decomposition after many hours, correCommercial C-14-labeled oxalic acid solution of high specific activity was diluted with A. R. grade inactive po- sponding t o the expected pH of pure KzCz04at the tassium oxalate solution to give a reagent of activity such concentration achieved. The initial drop in pH that when precipitated as calcium oxalate monohydrate, can no doubt be ascribed to the setting up of the the thin-sample count was approximately 600 per minute per C02/H20 equilibrium as decomposition commilligram. COz-free water obtained from an Illcoway Remences. The exchange investigation was made by search Model Ion Exchanger was used for all solutions. B. Apparatus and Technique.-The procedure varied determining the change in activity of both comlittle from that used in similar previous studies in these plexed and uncomplexed oxalate a t various stages 1ab0ratories.l~ Calculated volumes of standard solutions during the thermal decomposition process. The of the complex, inactive neutral or acid oxalate, and radioactive oxalate were separately brought to reaction tempera- results of such an experiment are presented in ture and the run commenced by mixing. Possible effects Table I. of light on both the cobaltiate and the chromiate exchange It is clear from these data that no appreciable reactions were avoided by use of foil-wrapped vessels.14 A back exchange of oxalate ion into the complex oc“zero time” 2-ml. sample could be withdrawn and the free oxalate precipitated within a few minutes, utilizing a Ca- curs during a time interval long enough to allow exCL/NHICl/NHdOH mixture N / 2 in each constituent as tensive decomposition of the complex (about 33% precipitant. After centrifugation and a double washing in this case). The constant activity of the prewith water, an alcohol slurry of the calcium oxalate mono- cipitated residual material (last column of Table I) hydrate was transferred to a weighed nickel-dated steel planchet and dried. When complexed oxalate activity of can reasonably be ascribed to separation contamithe cobaltiate was to be determined, the complex was de- nation.17 It may therefore be concluded that the composed by heating with alkali in the presence of precipi- decomposition does not provide a ready path for tating agent after separation of free oxalate, and the result- incorporation of labeled oxalate ion into the undeing calcium oxalate, calcium carbonate and cobaltous oxal-
constant over a range of sample thicknesses on either side of this maximum. (16) No attempt has been made t o apply temperature corrections to the pH data in the interpretation of the exchange kinetics. I t has been simply assumed that the correction factors for various hydro(9) s. P. L. sorensen, 2.a n o r g . Chem., 11, 1 (1899). (10) “Inorganic Syntheses,” Vol. I, 1st Edition, McGraw-Hi11 gen ion concentrations at a given elevated temperature are roughly Book Co., Inc., New York, N. Y., 1939, p. 37. equal. Thus they affect only the abiolute values of (H+) at that tem(11) W. F. Hillebrand and G. E. F. Lundell. ”Applied Inorganic perature, not the relative values. Analysis,” 2nd Edition, John Wiley and Sons, New York, N . Y., (17) It is seen that in terms of activity per milligram of the two 1953, p. 417-22. precipitates, the contamination does not exceed 476, nor does the total amount of this activity change appreciably from sample to sample. (12) H. H. Willard, N . H. Furman and C. E. Bricker, “Elements of Quantitative Analysis,” 4th Ed., D. Van Nostrand Co.,New York, The figures of column 4 of the Table indicate complete retention of the oxalate activity in the free oxalate portion throughout the 7.5 N. Y.,1956, p. 250-251. (13) J. E. Boyle and G. M. Harris, J . Am. Chem. SOC.,80, 782 hour period, within the experimental error. With the latter taken (1958). to be as great as 4%, corresponding to a maximum unobservable exchange during the period of the experiment, the oxalate exchange half(14) Detailed studies have shown the rapid photo-decomposition of the tris-oxalato complexes of CofIII) (ref. 6) and Fe(II1) (C.A. time cannot be less than about 130 hours. The half-time of decomParker, Proc. R o y . SOC.(London), 8230, 104 (1952); A2.36, 518 (1956)). position under these mme conditions IS cloee to 6.5 hoqrp.6
F. D. GRAZIANO AND G. M. HARRIS
332
Vol. 63
TABLE I mization of the cobaltiate independent of other exDETERMINATION OF A c m v I r r I E s OF FREEAND COMPLEXEDchange processes. Reactions 1 , 2 and 3 provide the unidirectional path for the decomposition by interOXALATED U R I N Q DECOMPOSITION OF K&!O(C~O~)~ AT 50" Initial conditions: (complex) = 0.03505 0.01460 M ; . DH . 7.0
min.
=
Activity, ct./min. Cor. Per mg. total oomfree plexed CIOl CaOd
Sample wt. (mg.) ComFree plexed CSOP CnOib
Time,
M; (oxalate)
0 3.78 34.88 822 3.9 . . .c 846 ...0 30 4.92 60 5.40 34.30 850 3.3 90 6.06 36.03 844 2.9 120 6.66 32.91 845 2.8 180 7.81 32.77 848 2.9 240 8.80 30.39 842 2.9 300 9.74 29.41 850 3.2 360 10.59 26.89 846 2.9 420 11.36 24.65 846 3.1 450 11.76 28.68 843 3.0 Pure calcium oxalate monohydrate samples. b Residual complex decomposition mixture samples. Sample spoiled. Uncorrected for self-absorption or scattering.
composed complex, in confirmation of similar findings by other workers.' The mechanism proposed for the thermal decompositiona has as the initial rate-determining step the reaction
+
CO(CZH~)~- Ha0 +HCo(C204)~-
+
c204-
+ OH-
This reaction would have to be essentially unidirectional to conform with our data. Stranks, et aE.,' who based their argument on the results of a study of the radioactivity of the C02 liberated from the decomposing exchange system, prefer the iollowing unidirectional initial reaction cO(C204)s" --f cO(cz04)z-
+ Con + co2-
This reaction would be consistent with the observed non-exchange of oxalate. The reversible initial process CO(cz0i)s'
I ' co(cz04)2- + Czo4-
proposed by Adamson, et aL16for the mechanism of cobalt electron exchange in the system Corn(11)oxalate/Co(C204)sE would be tenable only if the reverse step is slow, since rapid electron exchange between C204- and C Z O ris very likely. In a re-investigation of the cobalt exchange, Stranks, et al.,' find favor for a path involving a Co(II)/Co(III) binuclear oxalato complex intermediate. Their scheme to account for the various known reactions of the Co(C20&= ion is
+
CO(c204)~' HzO Co(C204)z.0Cz03.HzOe ( 1 ) Co(C204)z.OCzOs.HzO' + Co(Cz0i)zCOz COzHIO (2) cO(C~od)3~ COz- +C O ( C ~ O ~ ) ~ COz - ~ (3) C O ( C ~ O ~ ) ~ . O C ~ O ~ *CO*( H ~ OCz04)z(HzO)z" CO(CZ0&(H~O)-OC-C-O-CO*( CzOa)z(Hz 3 ) --O
+
[1 '
+
+
+
1111
+
+ +
1
Co*(Cz04)z.0Cz03.HzOM Co(CzO,)z(H,O)z(4)
Reaction 1 is proposed by analogy with the results of Llewellyn, et aL,B on the chromiate; it allows for rapid monoaquation, oxygen exchange and race-
nal electron transfer, and reactions 1 and 4 allow for the cobalt electron exchange process. Our results are consistent with this scheme, provided exchange of oxalate with the CO(C~O~)~.(H,O)~ion is quite slow, as seems likely in view of its relatively high stability.' B. The Tris-oxalato-chromium(1II) Ion.-The stability of aqueous solutions of was tested in two ways. First, the change in pH on long standing at the usual reaction temperature (75') was determined. It was found that in the range 4 < pH < 7, the acidity was constant for many hours. A slight decrease was detectable after one day, a period corresponding to several half-times of exchange for most runs. Similarly, the ultraviolet and visible spectra of the solution remained practically constant under the same conditions. In more acidic solutions, however, appreciable changes in both pH and spectrum were observed, as illustrated in Table 11.
c
TABLE I1 CHANGEOF pH AND ABSORBANCY OF ACIDIFIED AQUEOUS SOLUTION OF K&r( C104)8ON HEATING AT 75" Time (tw.)
%H +.(422mp) A.(573mp)
0 2.5 0.51 0.40
4.25 18 2.7 2.7 0.45 0.45 0.35 0.35 ( A log roll)
44 2.7 0.45 0.35
90 2.7 0.44 0.34
-
It is apparent that an equilibrium is rapidly attained, followed by no further re-adjustments. Other experiments showed that this equilibrium state was established within an hour, and that it involved liberation of a small amount of oxalate from the complex. The significance of this will be referred t o later. I n any case, since the attainment of the chemical equilibrium occurred within a time which was quite short relative to the exchange halftime, measurements of the rate of exchange were not subject to corrections due to this pre-equilibration. This fact was attested to by the complete linearity of the McKay plots of runs made a t low pH after a short initial period of slight curvature. When the pH exceeded 7, another type of decomposition process commenced, which eventually involved precipitation of Cr(OH)3. A thorough study of the reactions in this region has not been made, except to note the accelerating influence of high pH on the exchange rate. Experiments were carried out a t three temperatures (65, 75 and 85") under various reactant concentration conditions, The most comprehensive set of data was obtained a t 75", and the results of the various exchange runs appear as points in Figs. 1, 2 and 3. Entirely analogous curves can be plotted from the data a t 65 and 85", which are summarized in Table 111. A partial study of the effect of ionic streng;thl8 on the exchange has been made, as illustrated in Figs. 1 and 4. (18) A11 ionio strength adjustments were made by addition of the requisite amount of A. R. sodium perchlorate.
4
SUBSTITUTION REACTIONS OF OXALATO COMPLEX IONS
March, 1959
TABLE I11 RATEOF EXCHANQE IN THE Cr(Cp04),'/Cn1404" SYSTEMS (Free oxalate), M X 101
(Complex), M X 108
PH
(at R.T.)
A. R u n s a t 65' 4.57 3.95 1.88 3.95 4.57 7.90 1.86 7.90 15.81 4.56 15.81 1.86 15.81 4.48 15.81 1.92 15.81 4.49 15.81 1.94 15.81 4.49 15.81 1.98 B. R u n s a t 85" 5.27 3.95 4.64 5.27 3.95 1.98 5.27 7.90 4.64 5.27 7.90 1.95 5.27 15.81 4.55 5.27 15.81 1.94 15.81 15.81 4.52 15.81 15.81 2.04 31.62 15.81 4.52 31.62 15.81 2.04 47.43 15.81 4.50 47.43 15.81 2.07 Ionic strength = 0.300 in all runs.
5.27 5.27 5.27 5.27 5.27 5.27 15.81 15.81 31.62 31.62 47.43 47.43
6.6 X 7.2 X 6.9 X 8.2 X 7.1 x 10.0 x 1.9 x 3.0 x 4.0 X 5.9 x 5.6 X 9.0 x 3.6 X 2.6 X 4.0 X 4.1 x 4.3 x 7.4 x 1.2 x 1.5 X 2.4 x 3.1 X 3.7 x 4.7 x
=
(complex)[ka
+ kb(oxa1ate) +
k,,(H+)
lo-@ 10-8 10-8 10-n 10-8 10-8
10-7
10-7 10-7
. 200 -
P 100 -
0
A
3
-2 8
a:
-
80 605040-
-
ao2015-
-
0
10 -
A
8I
-
0 I
I
1
-
I
10-7 10-8 10-7 10-7
I
10-7 10-7 10-8
c; T%2 0
'
t
d
/ i
I
+ kd(H+)(oxalate)]
+
Cr(C204)2.OCzOsH.Ha0' Kt Cr(C204)2.0C203.H201 Hi0 Cr(CzO4)2(H@)z- C204- k d - s Cr(Cz04)z.0CYO~H.H20-HzO C ~ ( C Z O ~ ) Z ( H ~ O )HCa04Zk4,k-A Cr(Cz04)2~0C~Oa.H~0" HCa*O4Cr( C204)z.0C~*03.H~0' HClOn- k&s Cr(C2Or)2.OCa03H.H20' HC,*O4Cr(Cz04)z~OC2*OsH.H~O"HC204- k8,k-a _._____
.
I
10-8
A mechanism of exchange which is consistent in the range 2 < pH < 4 with this rate law is Cr(CzOa)a' + HzO J_ Cr(C2O4)2~OCZOa~H20" K I (1) Cr(C204)3e &of 1_
+ +
I
1
7 . 150 -
6-
k. = 1.1 X sec.-l kb = 1.1 X IOd4 1. mol --I sec.-l kc = 1.8 X 1. mole' sec.-l k d = 2.6 X 10-1 1.2 molea sec.-1
+ +
I
h
p
The curves shown in Figs. 1, 2 and 3 are calculated from this law, using the following "best-fit" rateconstant values
+ + + +
I
I
400
moles I.-' sec.-1
It is apparent that the reaction is first order in (complex) under all conditions. However, the dependence of the rate in the range 2 < pH < 6 on (free oxalate) and (H+) is of mixed order, partially zero and partially first.lg These data can be fitted by a rate law of the form R
300 -
R,
333
(2) (3) (4) (5)
(G)
(19) The apparent acceleration of the exchange above pH 6 most be due to new processes dependent on (OH-). Insufficient study of this effect has been made to warrant further discussion at this time.
1
2 3 4 5 (Complex), M X 102. Fig. 2.-Dependence of rate of exchang: on (complex); (oxalate) = 0.01581 M ; temperature, 75 ; p = 0.300; 0,R x 107 (PH2.15); A, R x 10s (PH 4.45:.
Assuming that each complex species other than Cr(C204)a31 is present in relatively small concentration,20 and that equilibria 1 and 2 are rapidly achieved and continuously maintained, it is seen that ka = K I ~ ~ ( H ~k bO=) ~Klk6(HzO), , k c = Kzk4(H%O)' and k d = K2ka(HzO). Reactions 1 and 2 permit the observed rapid racemization2 and oxygen exchanges of the complex, and the acid catalysis of these processes.* Reactions 2 and 4 account for the chemical equilibration in acid solution with slight liberation of oxalate, as mentioned above. The ionic strength data of Fig. 4 enable some (20) Support for this assumption comes from work of R. E. Hamm and R. H. Perkins ( J . Am. Chem. Soc., 77, 2083 (1955)), who showed CrO4' -+ Cr(CnO4)s. 2Hr0 t h a t the reaction Cr(C10SdHn0)lwas rapid and apparently complete a t 40°. The reaction is independent of ( H + ) in the range 4.0 < pH < 9.3 and also of (oxalate). This latter feature they explain by the suggestion t h a t chelation with waterelimination (reverse of reaction 1) is rate determining.
+
+
F. D. GRAZIANO AND G. M. HARRIS
334
I
I
I
P
1
I
300
25
Vol. 63
b
j
200
--. 7 20
100
d
90
ra * I
s 15
-i
8
.
W
5
30 20
10
X
a:
d *
5
II
10 9
e
$ 7 6 5 5 10 15 20 25 4 (Oxalate), M X 108. 3 Fig. 3.-Dependence of rate of exchange on (free oxalate): 2 (complex) = 0.00527 M ; tern erature, 75’; p = 8.300; 0,pH 2.15; A, pH 7.13; 0,~ 2 4 . 6 2 . 3.4
I
1
I
I
0.1
0.2
0.3
0.4
’
I
I
0.5
0.6
6.
3.2 A 3.0 2.8 2.6 2.4
Fig. 5.-Variation of rate constants of exchange with ionic strength a t pH 4.2: 0,Jcb X lo6; A, k, X 10’.
-’ -
I
#
I
I
e,
3 2.2
I
u
& 2.0 A 1.8 X 1.6 1.4 1.2 I
0.1
I
I
0.2
0.3
I
0.4
I
I
0.5
0.6
1
P.
Fig. 4.-Dependence of rate of exchange on ionic strength a t pH 4.2: top curve, (oxalate) = 0.01976 N ; bottom curve, (oxalate) = 0.00395 M .
conclusions concerning the mechanism of the reaction in the middle pH range (4< pH < 6). Since the (H+) is very low in these runs, the empirical rate law is approximated by R = (complex)[Ic, kb(oxalate)]. The two sets of data at differing (oxalate) enable calculation of k a and k b a t the various ionic strength levels. The logarithms of these calculated k’s have been plotted us.
