1920
Singh et al.
zene and toluene by SOC-. This reversible water elimination however was not observed in the hydroxylation of phenol, chlorobenzene, and nitrobenzene.I2 The increase in G(meta) with increasing metal salt concentration, may, however be solely due to a competition between oxidation and dehydration, assuming a greater rate constant for dehydration of the meta radical (11) than of the ortho and para isomer (I and 111). Contrary to the hydroxynitrocyclohexadienyl radicals where the ortho and para isomers are selectively reduced by Fe(CN)c4- no reduction of the three isomeric hydroxymethylcyclohexadienyl radicals by F e ( c N 1 ~ ~ was - observed (experiments 21 and 27). Only those radicals which have an electron-withdrawing substituent a to the radical site undergo reduction. This observation is in agreement with results of Walling and coworker^'^ on the redox reactions of a-hydroxyalkyl radicals with Fe3+-Fe2+. With KsFe(CN)s in neutral solutions we do not find any change in G(creso1s) nor in the isomer distribution with increasing concentration (experiments 18-21). This observation together with the minor change a t low pH indicates that the oxidation by KsFe(CN16 must be a fast process. This fast oxidation by KsFe(CN16 was also observed in the hydroxylation of nitrobenzene at low pH. We, therefore, suggest that the cresol isomer distribution obtained with K2Fe(CN)6 represents the relative reactivity of -OHradicals toward the different positions in the toluene molecule. At the higher metal ion concentration (5 X loh2M ) of the Cu2+ salts and K2Cr207 we observed almost the same isomer distribution as with KsFe(CN)s. From these results the
relative reactivity appears to be approximately 2:l:l. This preference for ortho-para substitution is to be expected on the basis of the well-known electrophilic character of the .OHr a d i ~ a l . ~From J ~ the results in Table I we can see that Cu(ClO& at a concentration of 5 X low3M and a pH of 2.6 gives more G(creso1s) and less G(benzy1 alcohol) than Fe(C104)s at the same concentration and pH. From this we conclude that Cu2+ oxidizes the hydroxymethylcyclohexadienyl radicals faster than does Fe3+. We have, therefore, the following reactivity sequence for oxidation: Fe(CN)s3> Cu2+ > Fe3+.
References and Notes (1) This paper was prepared in connection with work under Contract No. AT(40-1)-1833 with The U.S. Atomic Energy Commission. (2) E. J. Fendier and J. H. Fendler, Prog. Phys. Org. Chem., 8, 229 (1970). (3) M. K. Eberhardt and M. Yoshida, J. phys. Chem., 77, 589 (1973). (4) H. C. Christensen and R. Gustafsson, Acta Chem. Scand.,26, 937 (1972). (5) M. K. Eberhardt, J. Phys. Chem.. preceding article in this issue. (6) B. H. J. Bielski and A. 0. Allen, ht. J. Radiat. Phys. Chem., 1(2), 153 (1969). (7) H. C. Christensen, K. Sehested, and E. J. Hart, J. Phys. Chem., 77, 983 (1973). (8) C. R. E. Jefcoate, J. R. Lindsay Smlth, and R . 0. C. Norman, J. Chem. SOC.B, 1013 (1969). (9) C. Walling and R. A. Johnson, J. Am. Chem. Soc.,97,363 (1975). (10) 0. Volkert and D. Schuite-Frohlinde, Tetrahedron Left., No. 17, 2151 (1968). (1 1) C. Wailing and D.M. Camaioni, J. Am. Chem. SOC.,97, 1603 (1975). (12) L. 0.Shevchuck and N. A. Vysotskaya, Zh. Org. Khim., 4, 1936 (1968). (13) C. Walling and 0.M. ECTaliiwi, J. Am. Chem. SOC.,95, 844 (1973); C. Walling, G. M. ECTaliawi, and R. A. Johnson, ibM., 96, 133 (1974). (14) M. Anbar, K. Meyerstein, and P.Neta, J. Phys. Chem., 70, 2660 (1966).
Kinetics and Mechanism of the Osmium Tetroxide Catalyzed Oxidation of 2-Propanol and 1-Propanol by the Hexacyanoferrate( 111) Ion in Aqueous Alkaline Medium H. S. Singh,' S. P. Singh, S. M. Singh, R. K. Slngh, and A. K. Sisodia D8parfmeflt of chemjstry, University of Allehabad, Allahabad, M i a (Received April 25, 7974; Revised Manuscript Received June 4, 1975)
A study is reported of the hexacyanoferrate(II1) oxidations of 2-propanol and 1-propanol in the presence of osmium tetroxide as catalyst. The kinetic data suggest that the oxidation of these alcohols proceeds via the formation of an activated complex between the alcohol molecule and osmium tetroxide which rapidly decomposes to an intermediate product and osmium(V1) species. The osmium(V1) thus produced is rapidly oxidized to osmium(VII1) with hexacyanoferrate(II1) ion. The oxidation products are determined and a possible set of reactions for their formation is presented.
Introduction Sussela' and Solymosi2 have studied, from an analytical point of view, the oxidation of a number of organic compounds with hexacyanoferrate(II1) in aqueous alkaline medium using osmium tetroxide as a homogeneous catalyst. We were the first to examine the kinetic features of the OSmium tetroxide catalyzed oxidation of methanol and ethanol,'{by hexacyanoferrate( 111) ion in aqueous alkaline mediThe Journal of Physical Chemistry, Vol. 79, No. 78. 1975
um. Recently, we have also studied the kinetic features of osmium tetroxide catalyzed oxidation of ketones: ald e h y d e ~ a-hydroxy ,~ acids,6 and diols' with hexacyanoferrate(II1) in aqueous alkaline medium. However, the mechanism of oxidation of monohydric alcohols has yet not been studied and in the present study we have carried out the oxidation of 2-propanol and 1-propanol by aqueous alkaline hexacyanoferrate(111) ion using osmium tetroxide as a
1921
Oxidation of 1- and 2-Propanol by Hexacyanoferrate(ll1) homogeneous catalyst. The details of the results are presented and accordingly a reaction mechanism has been proposed.
Experimental Section Materials. (i) An aqueous solution of potassium hexacyanoferrate(II1) was always prepared from an Analar (BDH) grade sample. (ii) An aqueous solution of 1-propanol and 2-propanol was prepared daily from the Merck sample of G.R. grade. The alcohols were redistilled and the fractions boiling at 97.5 and 82.0' for 1-propanol and 2-propanol, respectively, were collected for the kinetic measurements. The solution of osmium tetroxide was prepared by dissolving the sample (Johnson Mattey & Co. Ltd.) in a solution of potassium hydroxide. The final strengths of KOH and that of o s 0 4 were kept 5.00 X and 3.93 X M, respectively. The samples of sodium hydroxide and potassium chloride were of G.R. (S. Merck) grade. The kinetic experiments were followed by adding the required quantity of the alcohol solutions maintained at a constant temperature to the solution of KsFe(CN16, NaOH, and Os04 kept in a reaction bottle a t the same temperature. The temperature of the reaction mixture was kept constant with an electrically operated thermostat with an accuracy of f0.1'. The progress of the reaction was measured by estimating the amount of hexacyanoferrate(I1) ion produced after a definite time interval with a standard solution of ceric(1V) sulfate, using ferroin as a redox indicator. This method always gave reproducible results. The final oxidation products were confirmed with paper chromatographic* studies. Results and Discussion Details of the kinetic data for the rate of osmium tetroxide catalyzed oxidation of 2-propanol and 1-propanal are presented in the Table I and Figures 1-5. The ionic strength of the medium was kept constant with potassium chloride. Figure 1 shows a zero-order plot for the rate of oxidation of 2-propanol and 1-propanol. It is obvious from these plots that the zero-order velocity constants begin to increase after a certain stage of the reaction. From these data we may conclude that these increases in the zero-order rate constant values are due to further oxidation of the intermediate product. In order to avoid the possible intervention of the products, the initial reaction velocity -(dc/dt) was calculated by plotting the remaining hexacyanoferrate(II1) vs. time. The initial -dc/dt values a t different initial hexacyanoferrate(II1) ion concentrations are listed in Table I and show the reaction rate to be independent of hexacyanoferrate(II1) ion concentration. The exact dependence of the reaction rate on hydroxyl ion concentration is shown graphically in Figure 2, which levels the reaction rate follows shows that a t low [OH-] first-order dependence and become independent a t higher concentration. I t has been observed that the order of the reaction is one with respect to both alcohol and osmium tetroxide (Figures 3 and 4). This gives rise to a rate law equation a t very low hydroxyl ion concentration as -d[Fey]/dt
= k[S][OH-][OsO,]
TABLE Ia A [1-Propanol] = 0.05 M [NaOH] = 0.05 M [ O S O ~=] 1.94 x
IM
[ -dc/dt 1 x [K3Fe(CN)G1, IO5, M x lo3 M min"
[-dc/dt] x 105, M min-'
[K3Fe(CNh], M x 103
_______________-____---
1.0 1.5 2.0 3.0
1.0 2.0 3.0 4.0 5.0 6.0
2.20 2.24 2.15 2.20 2.20 2.20 2.20 2.20 2.20
4.0
5.0 6.0 8.0
10.0
1.60 1.70
8.0
2.00 2.00 1.94 1.94 1.90
10.0
1.94
Temp = 30", p = 0.5 M (BJ 30
T i m e (mm)
25
20
IS
10
5
I
I
I
I
I
0.0
X
m 2
0.0
20
40
(A)
80
60
100
120
Time ( m i n )
Figure 1. Reaction feature with respect to ferricyanide: (A) K3Fe(CN)62.00 X M, 1-propanol 0.10 M, NaOH 1.0 x IO-* M, Os04 1.04 X M, p = 0.5 M, temp = 30'; (b) K3Fe(CN)6 2.00 M, 2-propanol 0.3 M, NaOH 0.11 M. OsOd 1.830 X M X p = 0.5 M, temp = 3 0'.
The osmium tetroxide in alkaline medium has been reported to exist as octahedral complexes of the formg trans [ 0 ~ 0 4 ( O H ) ( H 2 0 ) ] and ~ - [ 0 ~ 0 4 ( O H ) 2 ] ~The - . existence of these species might be considered via equilibrium step I of the proposed reaction scheme. It is also observed that the O S O ~ ( O H ) is ~ ~the - only reacting species of the osmium tetroxide.1° On the basis of the above-mentioned experimental results, the following probable oxidation scheme of these alcohols can be proposed: OSO~(H,O)(OH)" + [OH'] ( C)
-
* OSO~(OH)?'-+ K1
H?O (1)
(CJ
= complex (C,)
(ID
intermediate products
(110
OSO~(OH),~+ S
K2
(C,)
( 1)
where [Fey] = [Fe(CN)$- and [SI = [alcohol]. The value of k is calculated as 4.36 X IO2 and 5.15 x IO2 M - 2 min-I for 2-propanol and 1-propanol, respectively.
B [2-Propanol) = 0.10 M [NaOH] = 0.05 M [Osob]= 1.83 x 10-~11rl
C,
kl
(slow)
Os(Vl)
+
This scheme shows that osmium tetroxide forms an activated complex (Cz) which disproportionates slowly into a The Journal Of phvsiCel Chemistry, Vol. 79, No. 18, 1975
1922
Singh et al. [oK]xio2 11.0
10.0
9.00 I
1
8.00
700
I
I
(a)
5.00
4.00
I
I
I
3.00 I
2.00
1.00
I
I
0
-- -_*--.--.-_ --. ..
- 0.5s '1
6.00
- 0.6
3
0
-
0.4
I
>
0.3 -
0
0.2-
n
6 > 0.1 .t
LL
1
I
0
1.00
2.00
I
3.00
I
4.00
I
500
I
6.00
I
I
8.00
9.00
I
7.00
I
10.0
0
M, 1-propanol 0.10 M, os04 1.94 X lo-' M, Flgure 2. Effectof hydroxMe ion concentration on the reaction rate: (A) K3Fe(CN)e 2.00 X = 0.5M, temp = 30'; (B)K3Fe(CN)e2.00 X M, 2-propanol 0.3 M, Os04 1.83 X lo-' M, p = 0.5 M, temp = 30'.
p
Lo3 O
(BJ 12
14 I
8
IO
1
I
I
~i J io7 M 6 I
4 I
2 I
1.0
0.0
2.0
(a)
[os ]o,
x
105 M
Figure 4. Effect of variation of Os04 on the reaction rate: (A) K3Fe(CN)(, 2.00 X M, 1-propanol 0.05 M, NaOH 0.02 M, p = 0.5 M, temp = 30': (E) K3Fe(CN)6 2.00 X 2-propanol 0.30 M, M, p = 0.5 M, temp = 30'. NaOH 7.0 X 64)
Mrsubstrote]
-d[-Fey] -
Flgure 3. Effect of variation of substrate on the reaction rate: (A) (1propanol) K Fe(C& 2.00 X M, NaOH 5.0 X lo-' M, Os04 1.94 X 10- M, p = 0.5 M, temp = 3 0'; (B) (2-propanol)K3Fe(CN)6 2.00 X M, NaOH 5.0 X low3 M, Os04 1.83 X lo-' M, p = 0.5 M, temp = 30'.
a
first intermediate product and the hexavalent osmium, which in turn is reoxidized to its original oxidation state in two fast steps by the hexacyanoferrate(II1) ion. The first product undergoes fast oxidation taking more of Os(VII1) and F ~ ( C N ) G ~Under -. such conditions, it is better to express the final rate law in terms of the total osmium(VII1) concentration. [os(vnI)]T =
c + c1 + cz
( 2)
Now considering eq 2 the final rate law in terms of the decreasing concentration of the hexacyanoferrate(II1) would be The Journal of Physical Chemistry, Vol. 79, No. 18, 1975
dt
2k,K1Kz[s] [OH-][OS(VI10 ]T 1
+
Kl[OH'][l -t K*[S]]
( 3)
Since first-order kinetics was obtained with respect to alcohol concentration, the inequality 1 >> Kz[S]will evidently exist and rate eq 3 is reduced to
The derived rate eq 4 exhibits the observed kinetics. At very low hydroxide ion concentration the inequality 1 >> KI[OH-] is evident and rate eq 4 is reduced to
--d[FeyJ dl
~~,K,K,[S][OH'][OS(VIII)]T (5)
This explains the first-order kinetics of the reaction rate with respect to alcohol, hydroxyl ion, and osmium tetroxide concentration as well as the zero-order kinetics of the reac-
1923
Oxidation of 1- and 2-Propanol by Hexacyanoferrate(ll1) co"]-'x
I
r
a0
14.0 I
I
2.0
12.0 I
I 4.0
1
(8)
2
10.0
8.0
6.0
4.0
I
1
I
I
2.0
0.0
I
kiKz
kiKiKz 4.24 x lo2 5.72 x lo2
6.49 X 101 1.51 x 101
2-propanol 1-propanol
The close agreement in k l K l K z values obtained by the different methods (from eq 5 and 7) substantiates the validity of rate eq 4. Similarly, the values of klK2 obtained by two different methods (from eq 6 and 7) further indicate the validity of rate eq 4, and hence confirm the proposed reaction mechanism. In the oxidation of these alcohols the chromatographic study confirms the formation of oxidation products of oxalic acid and acetic acid in 2-propanol and propionic acid in 1-propanol. On the basis of these experimental results the full sequence of the steps would be I
I
60
BO [oH]-lx 10-2
10.0
I
I
12.0
14.0
(11
Figure 5. Plot of [-dc/dt]-' vs. [OM-]-' for (a) K3Fe(CN)a 2.00 X M, 1-propanol 0.10 M,Os04 1.94 X M,p = 0.5 M, temp M, 2-propanol 0.3 M, Os04 1.83 X 30'; (e) K3Fe(CN)6 2.00 X M, p = 0.5 M,temp 30'.
OS04(OH)22'
+
CH3CH,CHzOH
KZ
1-propanol [CH~CHZCH,-~OSO~(OH), 1'- (In (C,) k1
nfi
[CH3CH,CHz-O-Os03(0H),12'slow (C,)
CHsCH2CHO OSO~(OH)~ +~20" '
+
ZFe(CN):-
+
Os0,(OH)~2' (IIQ
last
OSO~(OH),~+ 2Fe(CN),'CH,CH,CHO propionaldehyde
Flgure 6. Effect of low hydroxide ion concentration dn the reaction rate. Caption same as Figure 2.
+
2H20
(0)
7 CH,CH,COOH propionic acid
Similarly with 2-propanol the following steps might again be proposed
tion with respect to the concentration of the hexacyanoferrate(II1) ion. The values of klKIK2 were calculated and these are 4.36 X lo2 and 5.15 X lo2 at 30' for 2-propanol and 1-propanol, respectively. At higher concentration of hydroxide ion the value of K1 will be quite large and the inequality Kl[OH-] >> 1 would dominate and rate eq 4 is reduced to
complex (C,)
-~ d[Feyl = 2k,Kz[S][Os(VII~], d2
This clearly explains the zero-order kinetics with respect to hydroxide ion concentration. Under such conditions the values of klK2 obtained are 5.90 and 1.3 X 10' for 2-propano1 and 1-propanol, respectively, a t 30'. Again rate eq 4 might be written as
h
V
CH,COO-
CH3-8-COO1 1 ( 7) ~ ~ l ~ , ~ , I S I [ O " IT l ~ O 2k,K,[Sl[Os(VIm ~ ~ ~ ~ ~ d IT
+
which shows that a plot of l/(reaction velocity] vs. l/[hydroxide ion], should give a straight line with a positive intercept a t t h e y axis. Such plots are shown in Figure5. This again shows the validity of the rate eq 4. The values of klK1K2 and klK2 obtained from the slope and intercept of the plots (Figure 5) are as follows:
yo-
(acetic acid) (oxalic acid)
COOThe oxidation products for acetone are reported to be oxalic acid (80%)and acetic acid (20%) with alkaline permanganate." With hexacyanoferrate(II1) in the carbonatebicarbonate buffer these oxidation products were also confirmed chr~matographically.~ Therefore, in the osmium tetroxide catalyzed oxidation of 2-propanol by hexacyanoferrate(II1) the oxalic acid and acetic acid will be approximately in the same ratio. The Journal of Physicel Chemistry. Vol. 79, No. 18. 1975
1024
Thomas J. Weeks and A. P. Bolton
In the proposed complex (Cz), the osmium will be in the eight coordination number, but the water molecule is not included in the complex (C2) only due to the simple structure of the molecule. The osmium(V1) species,lob O S ~ ~ ( O H ) obtained ~~-, after disproportionation of the complex (CZ)is shown to be oxidized fast with the hexacyanoferrate(II1) in the next step in order to regenerate the osmium(VII1) species. If we compare the k l K 1 K ~and klK2 values in both cases, it can be shown that the rate of oxidation is in the order 1propanol > 2-propanol, due to steric and inductive effect the rate of complex formation in 2-propanol and its decomposition will be slow when compared to 1-propanol. The abnormal increase in the velocity constant after a certain fraction of the reaction might be explained as due to fast oxidation of acetone and propionaldehyde obtained as the intermediate product from 2-propanol and 1-propanol, respectively (Figure 1).Thus, in order to determine the
actual nature of the reaction the initial reaction velocity was determined and the results are explained accordingly. Acknowledgment. Authors are thankful to the S.C.S.I.R., Lucknow (India) for financial assistance. References and Notes (1) (2) (3) (4)
N. Sussela, 2.Anal. Chem., 145, 175 (1955). F. A. Solymosi, et al, Magy. Kem. Foly., 62, 316 (1957). B. Krishna and H. S. Singh. 2.Phys. Chem., 231,399 (1966).
V. N. Singh. H. S. Singh, and 6 . B. L. Saxena. J. Am. Chem. Soc..91, 2643 (1969). (5) P. C. Pandey, V. N. Singh, and M. P. Singh, / M a n J. Chem., 9, 430 119711. N. P. 'Singh, V. N. Singh, and M. P. Singh, Aust. J. Chem., 21, 2913 (1966); N. P. Singh, V. N. Singh, H. S. Singh, and M. P. Singh, Aust. J. Chem.. 23. 921 (1970). H. S. Singh and V. P. Singh, unpublished work. R. D. Hartly and G. J. Lawson, J. Chromatogr., 4, 410 (1960). J. S . Mayell. lnd. Eng. Chem.. 7 , 129 (1968). (a) W. P. Griffith, Quart. Rev., 19, 254 (1965). (b) F. A. Cotton and G. Wikinson, "Advanced Inorganic Chemistry". 2nd ed, W h y , New York, N.Y., pp 1007 and 993. W. L. EvansandE. J. Wilzemann. J. Am. Chem. Soc., 34, 1096 (1912).
Thermochemical Properties of Ammonium Exchanged Type L Zeolite Thomas J. Weeks, Jr.;
and A. P. Bolton
Linde Division. Union Carbide Corporation, Tanytown Technical Center, Tarrytown, New York 1059 1 (Received March 7, 1975) Publicatbn costs assisted by Union Carbida CorporaNon
The thermochemical properties of ammonium exchanged type L zeolite (NH4KL) have been examined by infrared spectroscopy, thermal analysis, and high-temperature X-ray diffractometry. Differential thermal analysis of NH4KL in air shows an ammonia oxidation exotherm at 515'. Thermogravimetric analysis indicates that deammoniation and dehydroxylation overlap between 300 and 700'. Infrared spectra show that water can be removed by evacuation at room temperature and that a single weak zeolitic hydroxyl band occurs a t 3630 cm-'; major changes occur in the framework region upon deammoniation. High-temperature X-ray diffractometry indicates that thermal stability of NHdKL is very dependent on bed configuration during firing. A sample fired in a thin bed loses crystallinity a t 500' while one fired in a thick bed retains crystallinity until 8 0 0 O . Cell constants are relatively invariant with temperature.
Introduction Type L is a large pore (-8 A) synthetic zeolite with a high silica content (Si02/A1203 6).l I t possesses an hexagonal crystal structure and is composed of alternating cancrinite cages and double six rings arranged in connected columns. This gives rise to a large channel along the crystallographic c axis and a pore volume of 0.21 g of H2O/g of zeolite. Typical unit cell composition is Kg[(A10&( SiO&]22H20; approximately 2Wo of the potassium is not exchangeable. There have been a few investigations on the adsorption properties of cation exchanged2 and acid extracted type L,394 but little has been published on its catalytic activity. This may be due partly to the fact that the thermochemical properties of the ammonium exchanged form (a common catalyst precursor) have not been investigated. This is a study of those changes which occur upon thermal treatment of NH4KL as determined by thermal
-
The Journal of Physical Chemistry, Vol. 79, No. 18, 1975
analysis, infrared spectroscopy, and X-ray diffractometry. The effect of bed configuration during calcination is also discussed.
Experimental Section Materials. Synthetic KNaL (obtained as SK-45 from the Linde Division, Union Carbide Corp.) was exchanged four times with 10% aqueous NH4CI at reflux temperature. Deep bed calcined material was prepared by placing NH4KL in a preheated nonpurged muffle furnace a t 800' for 1 hr; the zeolite was in an uncovered porcelain crucible and had a depth to diameter ratio of 3 to 1. This material was subjected to a single exchange with 10% aqueous NH4Cl at reflux. Analyses of the starting KNaL (11, NHdKL (2), and the ammonium exchanged deep bed calcined NH4KL (3) are presented below. Oxygen capacity was determined a t -183O and 100 mmHg following vacuum activation of 5-g pellets at 400' for 4 hr.
