Nov., 1963
KINETICS OF REACTIOX BETWEES AKTIMOXY (111) AND FERRICYANIDE
7600 8. with the strongest band being the zero-zero band at 7590 The strongest CH3Oz band might absorb at wave lengths longer than 6500 k.,and, therefore, not be detected. I n addition, broad, weak bands could easily be missed. For example, the 35702750 8. formaldehyde bands were not observed even though, from the infrared evidence, formaldehyde must have been present in fairly large concentration. All of these objections make it seem unlikely that the 11.10-p species is the methylperoxyl radical, but they do not conclusively rule it out. It is probably fruitless to speculate further concerning the identity of the 11.10-p species. Additional evidence is needed. Use of would determine if the 11.10-p band is indeed due to a C-0 vibration. A grating instrument would have sufficient resolution to see if four bands ate produced in an 018-enriched matrix. I n addition, the low-lying optical transition predicted for CH302should be looked for more thoroughly. (42) F. W. Dalby, Can. J . Phys., 36, 1336 (1958).
2397
Conclusions The large ethane yield shows the importance of the cage effect even in solid oxygen. Ozone is formed by the photolysis of some azomethane photooxidation product. The light used to photolyze azomethane may decompose the 11.10-p species. The isotope effect shown by both the 10.22 and 11.10-p species proves that they contain oxygen but also that neither band is due to an 0-0 stretching vibrahion. It seems unlikely that the 11.10-p band can be assigned to the methoxyl or methylperoxyl radical. Further experiments are needed before positive 'assignment can be made. While this work illustrates the difficulties inherent in the identification of a product based on one infrared band, it also shows the potentialities of the matrix isolation technique for shedding new light on old problems. Acknowledgments.-The author is grateful to Dr. L. C. Roess for his interest and encouragement, to Dr. S. A. Francis and Professor R. M. Hammaker for many helpful discussions, and to Mr. F. J. Lopez for his help with the experimental tvork.
KINETICS OF THE REACTION BETWEEN ANTIMONY(111)RND FERRICYANIDE I N ALKALINE MEDIA BY LOUISMEITESAND RICHARD H. SCHLOSSEL~ Department of Chemistry of the Polytechnic Institute of Brooklyn, Brooklyn, New York Received M a y 18, 1963 The rates of oxidation of antimony(II1) by ferricyanide in sodium hydroxide media have been measured by following the polarographic diffusion current of ferricyanide in solutions containing excess antimony(II1) and also, with the aid of a simple device that causes the potential applied by a polarograph t o alternate between two preset values, by following the diffusion current of antimony( 111) in solutions containing excess ferricyanide. Over wide ranges of the concentrations of the reactants, the data can be described by rate laws deduced from a mechanism involving the reactions of the species (HO)zSb-O-Sb(OH)z and Sb(OH)e-$with ferricyanide and the existence of the former as the predominating species of antimony(II1) under the conditions employed.
Introduction I n developing a new technique of amperonietric titration intended to permit the location of the end point when the reaction is slow or incomplete, the reaction between antimony(II1) and ferricyanide was selected for study because of the farorable polarographic characteristics of the reactants and products. The results of a detailed investigation of its kinetics are described here. Experimental Ferrocyanide ion and antimony (V) are polarographically inert in strongly alkaline solutions; ferricyanide ion gives a cathodic wave rising from zero applied e.m.f.; and antimony(II1) gives both an anodic wave (€?L/~ = ca. -0.40 v. vs. 8.c.e. in 0.26 M hydroxide) reflecting its oxidation to the + 5 state and a cathodic wave (El/%= ca. -1.1 v. us. s.c.e. in 0.25 M hydroxide) reflecting its reduction to the elemental From these data it is apparent that, in a reaction mixture containing both ferricyanide and antimony(III), the diffusion current of ferricyanide can readily be measured a t an appropriate potential in the vicinity of -0.7 v. os. s.c.e., whereas the anodic diffusion current (1) This paper i s based on a thesis submitted b y Richard H. Sohlossel to the faculty of the Polytechiiic Institute of Brooklyn in partial fulfillment of the requirements for the B.S. in Chemistry degree, June, 1962. (2) J. J. Lingane, Ind. Eng. Chem., Anal. Ed., 16,583 (1943). (3) I. & Kolthoff !I. and R. L. Probst, Anal. Chem., 21, 753 (1949). (4) S. Vivarelli, Anal. Chim. Acta, 6,379 (1952).
of the antimony cannot be measured directly because it occurs only at potentials where ferricyanide is reduced. The rate of the reaction in the presence of excess antimony(II1) was followed by recording a current-time curve in the conventional manner a t -0.66 v. v8. a silver-silver chloride-saturated potassium chloride electrode.5 Appropriate volumes of water and standard solutions of sodium hydroxide, potassium ferricyanide, and antimony(II1) chloride (in a hydrochloric acid sohtion whose acidity was taken into account in computing the hydroxyl ion concentration of the reaction mixture) were mixed in the solution compartment of a modified H-celP that contained a little chloroform to prevent reduction of ferricyanide by mercury accumulated from the dropping electrode. Due precautions were taken t o avoid air-oxidation of the antimony(III), whose half-wave potential shows it to be a powerful reducing agent under these conditions. Reaction times were obtained from the recorder chart, whose rate of motion had been accurately measured. The sensitivity of the polarograph was readjusted periodically so that the pen deflection never fell below 507, of full scale. The recording was continued until no further decrease of current could be detected, and the residual current was taken to be the current that flowed after the reaction was complete. Data obtained with the same capillary on the diffusion currents of ferricyanide at various concentrations in antimony-free 0.25 F sodium hydroxide were used in computing ferricyanide concentrations from the recorded difi'usion currents. To permit the measurement of the diffusion current of antimony(II1) in the presence of a large excess of ferricyanide, a ( 5 ) L. Meites and S. A. Moros, Anal. Chem., 31, 23 (1959).
(6) L. Meites and T. Meites, ibid., 23, 1194 (1951).
LOUISMEITEXAND RICIIARD H. SCHLOSSEL
2398
synchronous motor-driven cam was so arranged as alternately to actuate and release a microswitch a t intervals of approximately 30 sec. The microswitch was connected to the initial-potential circuit of the polarograph. When it was open, a potential of -0.23 v. us. the silver-silver chloride electrode was applied to the dropping electrode by the span-potential circuit of the polarograph; when it was closed, a potential of -0.66 v. was applied by the span- and initial-potential circuits together. During the first interval, the sum of the diffusion currents of antimony(II1) and ferricyanide was presented to the recorder; during the second, only that of ferricyanide was recorded. By employing suitable compensation and damping, the difference between these two currents could be made to occupyalarge fraction of thewidth of the recorder chart. This difference is equal to the diffusion current of antimony after correction for the difference between the residual currents a t the two potentials employed, which was easily measured when the reaction was complete. I n this simple fashion useful data were obtained even when the diffusion current of antimony(II1) was only 0 1-0.2% of that of ferricyanide. This technique might be calied "low frequency (or large-amplitude) square-wave polarography." All measurements were made a t 25.00 =t0.02". Calibrated volumetric apparatus was employed throughout. The initial concentrations of antimony(II1) and ferricyanide were varied from about 0.02 to 10 mM; the diffusion-current data would have been afflicted by increasing uncertainties had smaller concentrations been used, while precipitation always occurred during the reaction when higher concentrations were used. One reactant was always present in a t least a IO-fold, and usually a 20- to 40-fold, excess, so that its concentration could be taken as constant. The hydroxyl-ion concentration was varied from 0.14 to 0.46 M .
Data and Discussion For every experiment a first-order pIot of the conceiitration of the reactant followed was coiicave downwards, indicating that the apparent order with respect to that reactant was smaller than 1. As the concentrations of the reactants were varied, the order with respect to ferricyanide appeared to vary from nearly zero to nearly 1, while the order with respect to antimony(II1) appeared to vary from about 1 / 2 to 1. A mechanism that serves to describe all of the experimental observations involves the assumption that the predominating species of antimony(II1) under these conditions is binuclear. No information bearing on this assumption is available in the literature; for whatever the analogy may be worth, however, bisniuth(II1) has been reported7-12 to exist predominantly in the form of a polynuclear hydroxo complex in alkaline solutions. The binuclear hydroxo complex is assumed to be in equilibrium with a monomeric species13
and also to be capable of reacting with ferricyanide. In this equation and hereafter, the symbols "Feic" and "Feoc" are used to denote ferri- and ferrocyanide ions, respectively. kz'
SbzlI1 3. Feic -+ SbIV
+ SblI1 + Feoc
(2)
To account for the fact that the apparent order with (7) F. GranBr and L. G. SillBn, Acta Chem. Scand., 1, 631 (1947). (8) J. Faucherre, Bull. SOC. c h i n . Prance. 253 (1954). (9) F. GranBr, A. Olin, and L. G. Sill&, Acta Chem. Scand., 10, 476 (1956). (10) R. W. Holmberg, K. A. Kraus, and J. S. Johnson, J . A m . Chem. Soc., 78, 5506 (1956). (11) A. Olin, Acta Chem. Scand., 11, 1445 (1957). (12) R. S. Tobias, J . A m . Chem. Soc., 8 2 , 1070 (1960). (13) Primes accompanying rate and equilibrium "constants" here signify that they actually vary with hydroxyl ion conoentration.
Vol. 67
respect to ferricyanide is always less than 1 and may approach 0, oiie must assume a further transition
yielding a complex that can react with ferricyanid~l~
+ Feic S SbIV + Feoc k4
(Sb'")'
(Id-
4)
(4)
while, finally
+ Feic +SbV + Feoc Id6
SbIV
(5)
This yields the rat,e law d_ [Feic _ ]_ - -2kz'[Sb2I1Ij dt
ka' [Sbllll) k,' [Sb211']
whence k3'[sb11'1) k,' [Sb;" ]
2]~~'[Sb:"]t
+c
(7)
If it is assumed that essentially all of the antimony(II1) in the solution is present in the binuclear form, then [ S b P ] = C 3 / 2 and [SbIII] = t'K'C3/2, where C3 is the number of moles per liter of antimony(III), and eq. 7 may conveniently be written In ([Feic]
+ a)
=
-bt
+c
(8)
where c is the constant of integration and
b
= k2'
C3
(9)
ab = k 3 ' d 2 K ' C 3 (10) For solutioiis containing a constant excess of ferricyanide, however, it is more convenient to write
whence, as above where f is the constant of integration and 1
e = - kz' [Feic]
4
(13)
Typical plots of the form of eq. 8 and 12 are shown in Fig. 1 to illustrate the conformity of the data to these rate equations. In every case the appropriate choice of numerical parameters served to reproduce the experimental values, within a reasonable estimate of the uncertainty of measurement, until the reaction was a t least 90% complete. The experimental conditions (14) The essential assumption is t h a t the product of the reaction described b y eq. l cannot react directly with ferricyanide. In eq. 2 we have indicated t h a t the antimony(II1) species formed when the dimer is oxidized is identical with the one formed when i t dissociates, With equal justice t o the rate data, however, one could write
Sbp
+ Feic +SbIV+ (Sb*II)' + Feoc 62*'
for this yields the same rate law with kz*' cussed further in footnote 15.
(2*)
= k ~ ' / 2 . This possibilityis dig-
TABLE I KINETICI’ARAMETEIZS 10’CsOR
6.86 10.72 15.87 19.85 51 .ti 104.3 0.22!) 0.!)61 1.019 44.8 48.5 47.5 44.9
FOR THE
REACTION BETWEEN ANTI>iONY(111) AND FERRICYANIDE I N SODIUM
~o~[Foic]~~
0.448 ,522 ,774 1.206 5.10 5.10 9.56 18.99 41.97 3.53 2.52 2.46 1.!)7
[NaOEII5
107k~~.\/~~~+
kl’bsO
0.243 .275 .253 ,235 .I335 .214 ,239 .229 .243 .461 .351 .145 .131d
0.175 ,380 .290 .217 .189 .l53 .195 .130 ,126 ,369 .348 .142 .140
IIYDROXlDE
SOI.UTIONS l@ks‘.\/K/ [OH -1’
ks’/[OH-]
3.58 6.10 6.80 5.58 4.31 7.40 1.49 5.91 4.24 36.8 14.25 0.82 0.83
0.72 2.5 1.38 2.9 1.15 4.2 0.92 4.3 .80 3.3 .72 (7.6) .82 (1.1) .57 4.9 .52 3.0 .80 3.8 .99 3.3 .98 2.7 1.07 3.7 Mean: 0.87 rt 0.19 3.5 f 0.6 a The first three coliimns give the initial concentrations, in moles/l., of antimony(III), ferricyanide, and sodium hydroxide in the reaction mixture. * When antimony(II1) was initially present in excess, values of b and ab were obtained from a plot according t o eq. When 8; then the value of kz‘ in the fourth column is equal to b/Cs and the value of k3‘v’j
