Sept. 5, 1955
-1517
KINETICSOF PERSULFATE OXIDATIONOF ISOPROPYL ALCOHOL
[CONTRIBUTION FROM
THE
DEPARTMENT OF CHEMISTRY, STEVENS INSTITUTE
O F TECHNOLOGY]
Mechanism of Organic Oxidation in Aqueous Solution. I. Kinetics of the Persulfate Oxidation of Isopropyl Alcohol BY LEONARD S. LEVITTAND EDMUXD R. MALINOWSKI' RECEIVED MARCH21, 1955 The kinetics of the persulfate oxidation of isopropyl alcohol were investigated a t 60 and 50" in buffered aqueous solution of PH 8.0 and ionic strength 0.5. The rate of disappearance of persulfate was followed by iodometric titration. The reactions all were first order with respect to persulfate. A limiting rate is approached a t higher initial alcohol concentrations, and thr rate becomes independent of initial alcohol concentration. At lower initial concentrations of alcohol the rate exhibits a marked decrease and becomes proportional to the alcohol concentration. Thus the reaction is over-all first order at higher alcohol concentrations and over-all second order a t lower concentrations. An empirical rate equation relating the observed rate constant to initial alcohol concentration is developed to fit the experimental data. On application of this equation to the data, good agreement between the calculated and experimental rate constants is obtained. An ionic mechanism involving a two-electron transfer, similar to that proposed for the persulfate oxidation of mercaptans and sulfoxides, is presented. From the proposed mechanism a rate equation of the required form is derived: -d lS,O,-] /dt = k& [SZO~-] [ROH]o/( [Sod-] ) ( K -2 [ROHIo). The common characteristics of many organic oxidations in polar media are discussed.
+
Introduction I n a previous paper2 we have shown that the rate of oxidation of diethyl and diphenyl sulfoxides by potassium persulfate a t 80" in aqueous solution of pH 8.0 and ionic strength 0.5 is first order with respect to persulfate but independent of the sulfoxide concentration over the range studied (0.010.02 M ) . Later work3 has shown that the rate falls off rather sharply a t somewhat lower sulfoxide concentrations, as was anticipated. In this low concentration range, then, the rate is proportional to the sulfoxide concentration and the reaction is therefore over-all second order. The same result has since been observed when thiodiglycol sulfoxide is oxidized by persulfate. I n the persulfate oxidation of mercaptans6 a limiting rate also is attained, and since the same appears to be the case when sulfides6 are oxidized by persulfate, i t was considered of interest to ascertain if this phenomenon extends to compounds other than those containing oxidizable sulfur. We have chosen isopropyl alcohol for this purpose, since the oxidation product, acetone, is not itself subject t o very rapid oxidation, and also because the kinetics of the chromic acid oxidation of this alcohol have been so thoroughly investigated.' A kinetic study of the persulfate oxidation of methanol under conditions quite similar to ours has previously been carried out; the reaction order was found to be three-halves with respect to persulfate and one-half with respect to methanol concentration.8 Experimental All kinetic runs were made a t 60.0", with an initial concentration of 0.02 M potassium persulfate, an ionic strength of 0.5, and pH 8.0. The samples were titrated iodometri(1) Done in partial fulfillment of t h e requirements for t h e M.S. degree a t Stevens I n s t i t u t e of Technology. ( 2 ) E. Howard, Jr., and L. S. Levitt, THISJ O U R N A I , , 76, 6170 (1953). (3) L. S . Levitt a n d E. Howard, J r . , unpublished results. (4) I,. S. Levitt and D. G. Ziebell, Abstracts of Papers (Organic Division), Cincinnati Meeting of A.C.S., April 7, 1955, p. 41N). ( 5 ) R . L. Eager and C. A . Winkler, C a z . J . Reseavch, B26, 527 (1948). (6) E. Larsson, Tvans. Chaimevs Uniu. Technol., Gothenberg, 87, 23 (1949). (7) F. H Westheimer and A . Novick, J . Chem. Phys., 11,506 (1943); F. Holloway. hf. Cohen a n d F. H. Westheimer, THIS J O U R N A L , 73, ( i i ( 1 9 j I ) ; A. Leo and F. H. Westheimer, ibid.. 74, 4383 (1H52). (8) P. D. n a r t l e t t a n d J . D. C o t m a n , J r . , ibid., 71,1410 ( 1 ~ ) .
cally. The isopropyl alcohol was Eastman Kodak Co. "spectro" grade. The buffer solution and solution for titration of the samples were prepared in the same manner as previously described,* and the experimental procedure also was identical.
Results I n Fig. 1 are presented logarithmic plots of the kinetic data obtained in the oxidation of isopropyl alcohol a t initial concentrations of 0.02, 0.01 and 0.005 M. Figure 2 gives similar plots for 0.20, O.l5,0.10,0.04,0.006and0.003Malcohol.
0.020 0.015
-.0.010 -
c
o"O.008
-
g 0.006 I
h
I\
0.004 \
I
t
0.002l
\. I
I
1
I
I
1
,
50
70 90 110 130 Time, minutes. Fig. 1.-Persulfate oxidation of isopropyl alcohol: A, 0.02 ,M (50"); B, 0.005 M (60'); C, 0.010 ( B O " ) ; D, 0.020 ;14(two runs, 60" j. 10
0.020 -0.01s c0.016 II: -0.014 0.012 0.010 0.0101
30
I , , , , , 'CJ 10 14 1s 22 26 Time, minutes. Fig. 2.-Persulfate oxidation of isopropyl alcohol a t BO": A, 0.003 M; B, 0.006 M; C, composite of four runs-0.20, 0.15, 0.10 and 0.04 M .
'
2
I
1
1
I
I
I
6
,411 the reactions are obviously and auite strictlv first order with respect to persulfate.' There are
LEOKARD S. LEVITT.IND EDMGND I> [SO,];then, since [XI cannot exceed [SO4],it is negligible compared to [ROHIo electron pair of the coordinate bond, resulting in a but not negligible in comparison to [Sod]. With transition state in which an incipiently forming positive organic species simultaneously expels a this simplification eq. 12 may be written second positive fragment or combines with a second Rp = [X]/[KOH]a([SO,]- [XI) (13) negative ion. from which [XI is given by -411 the common oxidizing agents are much more powerful in strongly acidic solutions principally, [XI = Kz[ROH]o[SOjI/(I;.,[ROH]~ 1) = [ROH]U[SOI] ;( [ROHlo R--2) (14) it seems, because of the ease with which the active Substituting in eq. 14 the equivalent of [SO4] oxidizing species are formed in the presence of a given by eq. 10, and setting the over-all rate equal proton donor by reactions typified by the following examples to ksX, one obtains the final rate expression
+
+
2HOSO. 2HOCl
It is seen that this equation is of the same form as the empirical expression (eq. l), and that therefore kmax
=
k,&j
(16)
i[SOA']
and
K-: = Ij& (1:) Thus the equilibrium constant K z for the formation of the ester X may be calculated from eq. 17. For isopropyl alcohol a t GO", Kz = l / T . O X = 1.4x 102. From eq. 16, the product kaKl can evidently be evaluated, but the rate constant k3 and the equilibrium constant K1 cannot be individually calculated from the data obtained in this investigation. The Common Characteristics of Organic Oxidations in Polar Solution.-The mechanism postulated here for the persulfate oxidation of a secondary alcohol seems quite possibly to be the general mechanism by which alcohols (and perhaps many other organic compounds) are oxidized in polar media by the common strong inorganic oxidizing agents. The mechanistic courses of these reactions appear almost always to possess in common the following characteristics. 1. The active oxidizing species is rarely, if ever, a negative ion ( e . g . , X O r , CrOT, S z O ~ C10-, , IO,, OOH-), but appears to be most usually an b =
NO.;
11
T
OC1-
1IOOIi
+ H26)SO2 4- H,hCI
-
H200I~I
+ I-lIpOhI1lO, ~I.C~C~O.,H :jIi+- + CrO,= 'ti
7
+ 502 I L O + bl IizO + O H . 1 1 2 0 + Mk>., H ~ O+ C ~ . ~ I I H 2 0
-=
11110,-
The four characteristics outlined above provide the common basis of many oxidation mechanisms previously proposed, such as the oxidation of merc a p t a n ~and ~ sulfoxides4 by persulfate ; the chromic acid oxidation of secondary alcohols'; the oxidation of alcohols'* and of formic acidIg by hypochlorous acid; and the oxidation of sulfides by acidic hydrogen peroxide.20 One example, the hypochlorous acid oxidation of primary alcohols,18will suffice for illustrative purposes. This reaction apparently proceeds through the intermediate formation of the hypochlorite ester in the following manner 13'
+ IIOCl
RCI12:O:li
H:&l
+ 21: - J_
HnO -.. .. RCI1:::O:CI: , . .. H
+ C1' HC
+
electron deficient cationic species (e.g., NOz,
+
+J-
-
Cr02H, C1, IOn, HOi or a highly polar molecule s+s-
sts-
s+a-
(e.g., XOZ, Cr03, 0.50,) capable of undergoing a
-
+
simple two-electron transfer" (e.g., N O n .--t XO,;
+
Cr03H -
+
CrOJ-;
SO4
Soh=; C1+
-+
Cl-;
IO3 + IOa-). These electrophilic substances are generally identical with those responsible for ordinary aromatic substitutions. 2. The positively charged atom of the active oxidizing species initially attacks a free electron electronegative pair of the ,reductant . ,. .a. t a relatively .. site (e.g., -0--,--S-, -X:, =9-)forming a coiirdinate .. ,. .. bond. The atom attackcd gains a positive charge. 3. The positive charge is removed by expulsion of a positive fraqinent (usually H+) or by conibination with a readily available anion. ( l i ) T w o electron onidation~reii,lctionrcnctions have been well established in m a n y instxnce.: siicli a s i n t h c oxidation of sulfite b y c h l o r a t e C J . Hallierin and 11. ' r ~ i i ! )l'irrs ?, J O I ' R N A T . , 74, .< ' :> f 19.52 A . 'i sI , , , I, :,,,
