NOTES
1518
ary alkyl radical. 3-5 The subsequent reactions have not been elucidated in detail, although radical disproportionation and recombination undoubtedly occur. The major difficulty in the study of these systems is that there are several low activation energy, highly exothermic reactions which may and probably do occur simultaneously. The result is a variety of products particularly in the case of the higher olefins. The reaction between hydrogen atoms and some condensed 1-olefins has been shown to occur a t - 19505s6 Under these conditions, atomic cracking reactions' were found to be absent (methane is not formed as a reaction product), so that the reaction system is considerably simplified. Further, large differences in reaction rate among the various olefins are found. 1-Butene a t -195" undergoes rapid reaction with hydrogen atoms while 2butene under the same conditions is inert. If in the reaction between 1-butene and hydrogen atoms the double bond is shifted, 2-butene should accumulate in the products. This does occur, as we have reported previously.6 Two possible disproportionation reactions resulting in a double bond shift are 2CH3CH2CHCH3----f CHs(CHz)zCH, H CH3CH2CHCHs +Hz
+
+ CHaCH=CHCHa + CHaCH=CHCHa
Vol. G3
radical are in the order tertiary > secondary > primary. The rate of double bond isomerization occurring by this mechanism is thus consistent with bond strength considerations.
THE RATE OF OXIDATION OF ITO HYPOIODITE ION BY HYPOCHLORITE ION BYYUAN-TBAN CHIAAND ROBERTE. CONNICK Received January SI, 1969
In studying the behavior of iodine in the +l oxidation state in alkaline solution,l it was convenient to prepare IO- solutions by the oxidation of iodide ion by hypochlorite ion. Under most of our conditions this reaction appeared to be instantaneous as judged by spectrophotometric observations; however, a t high hydroxide and low iodide concentrations at least part of the change from the hypochlorite spectrum to the hypoiodite spectrum occurred a t a detectable rate. A fast mixing technique was then applied to make the reaction more readily measurable. The net reaction was shown to be I-
(1) (2)
Reaction 2 may be established by using deuterium instead of hydrogen atoms. The accumulation of H D in the gas constitutes proof of (2). The exposure of solid isobutane a t -195' to deuterium atoms did not yield any HD. Since the tertiary hydrogen in isobutane was not abstracted under these conditions, the formation of HD by reactions other than (2) is unlikely.
+ c10- 1JIO- + c1-
(1)
by studying the stoichiometry spectrophotometrically a t varying ratios of I- to ClO-.' For interpreting the rate data the rate constant IC' WRS defined
where parentheses indicate concentrations in moles per liter. In Table I are given values of IC' a t a variety of experimental conditions, It is seen that the rate is first order in (I-) and (C10-), a TABLE I result also borne out by the time dependence of the TUNGSTEN RIBBON TEMPERATURE, 1600'; EXPOSURE TIME, individual runs. In the last column is the rate con30 SECONDS; INITIAL D,, 50~1.AT -195' stant IC which incorporates an inverse hydroxide Olefin HD/DP dependence Propylene 0.0 1-Butene .2 3-Methyl-1-butene .9 a Corrected for the Hz formed as a result of isotope equilibration on tungsten ribbon.
Table I gives the H D formed as a result of irradiating propylene, 1-butene and 3-methyl-lbutene with deuterium atoms a t -195". Hydrogen abstraction from propyl radicals by deuterium atoms does not take place a t - 195'. However, the warmed-up products of the reaction D CH&H= CH2 showed the presence of propane-dl.6 Whether the reaction
+
2CHaCHCH2D --.) CH2=CHCHzD
+ CHaCHzCHzD
occurred at - 195" or during warm-up has not been established. The relative rates of abstraction of the hydrogen alpha to the free spin site of the alkyl (8) W. J. Moore, Jr., and L. A. Wall, J. Chem. P h u s . , 17, 1925 (1949). (4) J. N. Bradley, H. W. Melville and J. C. Robb, Proc. Roy. Soc. (London), 336,339 (195b). (5) R. Klein and &I. D. Rcheer, J. Am. Chem. Sor.. 80, 1007 (1958). ( 8 ) R. Klein and M. D. Scheer, THIS JOURXAL,63, 1011 (1958). (7) B. de B. Dartvent and E. W. R. Steacie, J. Chem. P h y s . , 13,663 (1945).
d(I0-) - k(I-)(ClO-) ~ dt
(3 1
(OH-)
TABLE I RATEDATAFOR OXIDATION OF I- BY IOTemp. 25", ionic st,rength 1.00 M
4.00 2.00 2.00 2.00 2.00
2.00 4.00 2.00 2.00 2.00
1.00 1.00
1.00 0.500
0.250
0.004 0.002
.002 .502 .752
60.35 5 62.7* 5 60.62~ 5
110 234
f: 10 2Z20
60.315 62.7&5
60.655 58.05 5 58.5*5
The hypoiodite formed is itself unstable with respect to disproportionation to iodate and iodide. Under the conditions chosen for the present study this latter reaction was much slower than the Oxidation of iodide by hypochlorite and could be neglected in interpreting the data. Discussion.-Anbar and Taube2 have presented two possible mechanisms for the reaction of halide (1) Y.-t. Chia. Thesis, University of California, Berkeley, June, 1958; University of California Radiation Laboratory Report UCRL8311,June 2,1958. (2) M. Anbar and H. Taube, J . A n . Chem. Soc., 80,1073 (19581,
NOTES
Sept., 1959
+
+
ions with hypohalite ions: XYO- = XOY-, when the form of the rate law is that found in the present study (equation 2). In the first mechanism the halide ion attacks the oxygen of the hypohalous acid. For our case the steps would be
._
+ HzO HOCl + OH- rapid equil. + HOC1 --e+HOI + C1- rate determining IO- + HzO rapid HOI + OH-
C10-
I-
I n the second mechanism the halide ion attacks the halogen of the hypohalous acid to yield the halogen molecule, which subsequently hydrolyzes to yield HOX aiid Y-
+ H?O e HOCl + OH- rapid equil. + HOCl --+ TCI + OH- rat)e detn. IC1 + OH- r-f 1011 + CI- rapid IOH + OHTO- + HSO rapid
'210-
I-
Using the unpublished experimental results of Anbar aiid Rein they concluded that the exchange of bromine between Br- and HOBr proceeds by attack of the Br- on the oxygen of HOBr, corresponding to the first mechanism. No conclusion could be drawn about the exchange of chlorine between Cl- and HOC1. From a comparison of several rate constants they reached the tentative conclusion that the reaction between Br- and ClO-, with rate law analogous to equation 3,3goes by the first mechanism also. From the above results no definite extrapolation can be made to the I-C10- reaction, although the first mechanism would seem more likely.
1519
syringe through a No. 16 needle. The syringe was filled to the 2-ml. mark with care being taken to avoid bubbles. With the Gary spectrophotometer recording at 400.6 mp, the syringe was quickly emptied manually with the tip of the needle just below the surface. The measurement of absorbance was carried on as a function of time at the fixed wave length. Zero time was taken as the point at which the absorbance on the spectrophotometer chart began to rise. Other times were calculated from the running speed of the chart. A typical set of points (third experiment of Table I) read from the chart is shown in Fig. 1 where the absorbance is plotted against the time in seconds. Slopes read from such plots yielded values of d(I0-)/dt for calculation of k' by equation 2. The molar absorptivity of IO- and GlO-at 400.6mpare 38.5 and 0.510, respectively.' Since the exact time of mixing (-1 sec.) was not known, the zero time was somewhat uncertain. The use of derivatives eliminated the need for zero time in the treatment of the data. The syringe was lubricated carefully with a small amount of Kel-F No. 90 grease in the part above the 2.5-ml. mark. Too thin a layer gave rise to bubble formation. A qualitative test showed that the attack of the stainless steel needle by hypochlorite was unimportant. No effect was found on reversing the order of mixing by injecting hypochlorite into iodide. In the hydroxide dependence experiments the ionic strength was held constant by the addition of sodium chloride.
THE DIFFERENTIAL THERMAL ANALYSIS OF PERCHLORATES. 111. THE SYSTEM LiC104-NH4C104 BY MEYERM. R ~ A R K O W I T Z AND ROBERT F. HARRIS Poote AJzneral Coinpanil, Reseal c h and Development Laboratoraes Berwyn, Pennsyluanzn Recezued March 25, 1059
A previous study1 had shown the utility of anhydrous lithium perchlorate as a component in phase investigations of ansolvous perchlorate systems. As an extension of this earlier work, a substantial portion of the system LiC104-NH4C104 lias now been elucidated up t o liquidus temperatures consistent with the thermal stability of ammoniuin perchlorate. These results are of interest in that a phase diagram involving the thermally relatively unstable ammonium perchlorate has been constructed, and that a t 200" the rate of thermal decomposition of this salt in adrnixturc with lithium perchlorate may be qualitatively re' lated to the composition of the mixture.
0.160
u 120 2
50.080
eE
00.04
0
80 120 160 200 Time, sec. Pig. L-Tlie absorbance-time curve for the third cxperiment of Table I. Experimental The reaction was followed spectrophotometrically on a Gary recording spectrophotometer. The rapid mixing was mcomplished by injection with a hypodermic syringc, following the method developed by Stern and DuBois4 and M e r applied by Below.6 The 2.15 cm. absorption cell was that used by Below. Eight ml. of solution with thc (IeRired hydroxide and hypochlorite concentrations werc laced in the cell and the absorption recorded on the chart. yodide solution containing the same hydroxide concentration as that in the cell was introduced into a hypodermic 0
40
(3) L. Farkas, M. Lewin and R. Block, J . A m . Chem. Soc., 71, 1988 (1849).
(4) K . G . Stern and D. DuBois, J. B i d . Chem., 116, 575 (1936). ( 5 ) J. F. Below, Jr., University of California Radiation Laboratory Report UCRL-3011,June, 1955.
Experimental Procedures The eqmiinental arrangement for carrying out tlic differcntial thermal analyses (d.t.a.) at a lincar heating rate of 5' per minute has been described previously.1 All liquidus temperatures were obtained by visual observations and were measured with a calibrsted mercury-in-glass thermometer. Dry nitrogen gas was passed over the sample during each d.t.a. run and liquidus determination to avoid moisture absorption from the atmosphere. The preparation of anhydrous lithium perchlorate has nlready been detailed.' Analysis of product by precipitation of nitron perclrlorate: Clod-,94.1 (calcd. 03.5). Reagent grade ammonium perchlorate, dried at 105" for two hours was analyzed for ammonia content by distillation with sodium hydroxide solution. Analysis of product: "3, 14.37 (calcd. 14.49). The Thermal Behavior of the Components.-Lithium perchlorate showed but one reversible break during any d.t.a. run up to about 260' and this was attributed to fusion a t 247O.I Ammonium perchlorate on the other hand does not fuse but does undergo a reversible crystallographic transi(1)
M. hl. Markowita, T H IJOURNAL, ~ 62, 837 (1858).
