Herbert R. Ellison Wheaton College Norton, Massachusetts 02766
I
I
A Simplified Apparatus for Gas Phase Reactions Decomposition of di-t-butyl peroxide
During the past few years several experiments have appeared in THIS JOURNAL describing apparatus suitable for the study of gas phase reactions and, in particular, for measuring the rate of decomposition of di-t-butyl peroxide.' The basic simplicity of the experimental measurements and the theoretical importance of the results make gas phase reactions most useful for consideration in both elementary and advanced courses in chemical kinetics. Unfortunately the apparatus described in these articles either requires the services of a skilled glass-blower for their fabrication or is expensive and difficult to purchase. This communication presents equipment which is easy to assemble and operate and which combines ruggedness and low cost xith the ability to yield reasonably accurate results.
is made from 4-mm tubing bent in a U and firmly attached to a. 50-cm segment of meter stick which is supported by a ring stand. The manometer is connected to the reaction vessel through a 'Jlj standard taper joint. The use of standard taper joints enables the apparatus to he quickly set up or disassembled for cleaning. I t is necessary to hest the side arm leading to the manomT O PUMP
Experimental
two stopcocks A and B. Stopcock A connects the system to the vacuum pump, not necessarily a high vacuum pump, and s t o p cock B leads to the di-1-hutyl peroxide reservoir, a 50-ml round bottom flask with a 14/20 standard taper joint. The manometer
' PRICE;A. H., AND BIKER, R. T. K., J. CHI:M. EDUC.,42, 614 (196.5); GUILLORY, W. A,, J. CAEM.EDUC.,44, 514 (1967); TROTMAN-DICKENSON, A. F.,J. CHEM.EDUC.,46,396 (1969).
Figure 1.
Apporatut for gar phore reactions.
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eter in order to prevent acetone, one of the reaction products, from distilling out of the reaction flask. This can be done either by using a heat lamp or, with more eese, by wrapping the side arm with heating tape. I t is not critical to rtccurately control the temperature of the side arm since its volume is made much smaller than the volume of the reaction flask. However it must be higher than the boiling point of acetone (56'C) and preferably within 10-2O0 of the temperatureof thereaction flask. Temperature of the reaction flask is controlled to +0.3' by immersing the vessel in an asbestos wrapped liter beaker containing silicone oil m d heated and stirred by a combination hot plate-stirrer. The experimental procedure is. quite simple. The peroxide, vacuum distilled before use if necessary, is placed in the rese~voir.~ Stopcock I3 is closed and A opened and the reaction vessel degassed while the oil bath is brought to constant temperature between 140 and 160°C. Stopcock B is then opened and the whole system pumped dawn as low as possible. Stopcock A is closed and a beaker of water a t 60-65'C placed under the resemoir and about 5 a m m of di-tbutyl peroxide vapor forced into the reaction flask. Stopcock B is then closed, the pressure read on the manometer, and the timer started. This is the initial pressure Pi. The total pressure P , is read at suitable intervals over 8. period of about thirty minutes.
Raley, Rust, and Vaughau3 and Batt and Benson4 have shown that the decomposition of di-t-butyl peroxide proceeds according to the following scheme
-+ -
(CHa)&OOC(CHa)a 2 (CHs)aCO.
+2
2 (CHa)aCO.
CHa.
2 CHJ.
2 (CHS)~CO
CIHs
(1) (2) (3)
The first reaction, the cleavage of the 0-0 bond is the rate determining step. The overall stoichiometry of this decomposition shows that three stable molecules are produced for each molecule of peroxide that reacts. Thus the first order rate equation can be shown to be1
where P , is the initial pressure of the peroxide and P , is the total pressure of reactant and products at time t. The first-order rate constant is kl. Equation (4) predicts that a plot of log [(3P, - P,)/2] versus t d l be linear with a slope equal to - k1/2.303. I n a 3-hr laboratory period it is possible for students to carry out t ~ runs o at the same temperature but with different initial pressures of peroxide and a third run a t a different temperature. From their results they can verify that the reaction is indeed first order, and they can also evaluate the activation energy for the reaction. Figure 2 presents an Arrhenius plot of the results ohtained with this apparatus by eight students over a 4-yr period. Least-squares analysis of the data yields the equation
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Figure 2. ~rheni"'; plot of results obtained by eightstudents during the port four yean.
log kt (sec-1) = 15.35 - 36,800/(2.303 RT)
(5)
The equation obtained by Raley, Rust, and Vaughan3 is log k, (sec-I) = 16.51 - 39,100/(2.303 RT)
(6)
and the result found by Batt and Benson4is log k, (sec-1) = 15.60 - 37,400/(2.303 RT)
(7)
The agreement between the student results and the literature equations is seen to be reasonably good. I n general student values of k1 have been within 10-2070 of the values calculated from either eqns. (6) or (7). Larger errors occur when students fail to heat the side arm leading to the manometer to a high enough temperature to prevent condensation of acetone. This then results in P , being smaller than it should be, thus increasing log [(?,Pi - P,)/2], and finally resulting in an experimental kl which is too small.
*
NO s~ecisilnremutions need be emuloved in handline di-t-
a t atmospheric pressure without decomposition and is insensitive to the usual organic peroxide tests. Di-t-hutyl peroxide is "exceedingly stable" in the presence of strong bases, is not affected by concentrated HC1 and dissolves in concentrated H2SO4 to form polymeric hydrocarbons. I t can he reduced to t-hutyl alwhol with finely divided Na or with Zn dust in xylene or other 4nlventq . ... . .... . a RALEY, J., RUST,R.
F., A N D VAUGHAN, W. E., J . Amer. Chem.
Soc., 70, 88 (1948).
S. W., J. Chem. Phys., 36, 895 (1962). BATT,L.,AND BENSON,
