1250
NOTES
Oxygen exchange between chemisorbed C 0 molecules on Ni surfaccs is somewhat analogous to experimental observations made with CO on catalytic Ice by Webb and 14>ischens.5 These workers observed that on Ve appreciable oxygen exchange occurred a t - 78' and that statistical equilibrium was achieved a t about 160°. They also showed that the presence of the gaseous phase at relatively high pressure (300 torr) increased the rate of approach to statistical equilibrium on the Ice surface. A cknowledgmenl. The author wishes to acknowledge the financial support given to this work by the Research Corporation and by Antioch College. Particular acknowledgment should be made to the Kettering Research Foundation for their gift of equipment and for their kindness in allowing the use of the mass spectrometer for this work. The author wishes to thank Mr. Howard Bales for his work in performing the mass spectral analyses. Finally, tho author wishes to acknowledge the encouragement provided by Prof. R. G. Yalman a t Aritioch College.
troduccd into a 5-crn. spectroscopic cell, and the conccntration of Iz present a t any temperature was determined from thc absorption at 4500 A. (see ref. 7 for further details). Iklow about 35 O , solid complex could be seen and since there was always an excess of dioxane, the concentration of I2 was too low to mctasure satisfactorily. After each change of ternporaturc, sufficient time was allowed for cquilibriurn to bo attained. In each expcrirncnt, iodine concentrations were measured at successively higher temperatures until dissociation was complete (50-60 O depending on relative concentrations) and the values were rerncasured (and found to agree) as the temperature was being lowered; the equilibrium constants werc then calculated from the initial pressure of dioxane and thc measured concentration of iodine at complete dissociation. Various forms of equilibrium constant expression were explored, but the most satisfactory one was found to be
K,
=
l / [ I z ][dioxane]
corresponding to the equilibrium 1. The heat of reaction 1 was determined from the slope of the plot of log K , against 1/T using
IIeats of Formation of
a In K J ~ ( ; ; ) =
Dioxane-Halogen Complexes
Similar measurements were made with Brz and 1,4dioxane, and with I2 and 1,3-dioxane1 where a liquid complex is formed: nicasurcments could not be made
by C. A. Goy and H. 0. Pritvhard Chemistry Department, (lnicersity of Manchester, Manchester I S , England (Rivxicrd ,Voaember I , 1,963)
The oxistcncc of a liquid complex between bromine and 1,3-dioxanc1 arid of solid 1 : I complcxes of both bromine and iodine with 1,4-dioxane2 has h e n known for a long tiinc. S o thcrniodynamic study of these substances has been reported, although both equilibrium3 and calorimetric4 studies havc becn made on the solutions of iodinc and 1,4-dioxanc. The structurc of the solids has rcccntly been dctcrmincd by X-ray methods, and it was found that the halogen molecules act as a bridge bctmccri the oxygcii atonis of different dioxane rnolccules5 in contrast to the theoretical description given by hIulliken.fi Spectroscopic Studies. I t was found that the cquilibriuni Iz(g)
+ 1,4-dioxanc(g) [ 1,4-dioxane -
121
(s)
(I.)
could wadily tw studied spectroscopically in the ten)pcraturo rang(: 35- 60". Known aniounts of 1,4dioxanc: (about 2 -I cni. prwsurc) and iodinc were inT h e .lo?rrnal of Physical Chemisl7y
- 2 ~ AH/R
Table I : Spectroscopic Study of the Equilibria between Halogens and Dioxanes
Complex
Temp. range, OC.
36-60 1,4-l)ioxane-Iz(s) 1,4-I>ioxane-Br2(s) 23-42 1,3-l)ioxane-I~(1) 30-60
Menn lo$
K, a t
No. of detn.
2.5". mole-? I.Zn
10
8.4 f0.2 6 . 3f0.2 7.6 f0.1
10 4
Mean AH for reaction 1. kcnl./mole"
- 27.4 f 1 . 5 -26.1 f 1 . 5 -14.4 f 1 . 5
The errors quoted are such t h a t all determinations lie within these limits.
(1) H. T. Clarke, .J. Chrm. Soc., 101, 1788 (1912).
H.Rheinholdt and It. Roy, J . prakt. Chem.. 129, 273 (1931). (3) J. A. A. Ketelaar, .I. Phys. R a d i u m , 15, 197 (1954); D. I,. Qlusker and 11. W. Thompson, . I . Chem. Soc., 471 (1955); It. S. Drago and h '. J. Rose. .I. A m . Chem. Soc., 81, 6141 (1959). (4) K. Hartley and H . A . Skinner, Trans. Faraday Soc., 46, 621 (2)
( 1 950).
(5) 0. Hassel and J . IIvoslef, Acta Chem. Scand.. 8, 873 (1954).
(6) 11. S. hlulliken, J . A m . Chem. Soc., 72, 600 (1950). (7) C. A . Goy and 11. 0. I'ritchard, .I. Mol. Sprctry., 12, 38 (1964).
NOTES
1251
~
~~
T a b l e I1 : Calorimetric Study of Halogen-Dioxane
Final concn.,
Systems AH(ao1n.) oomplex, kcal./mole
System
&/I.
AH(ao1n.) halogen, kcal./mole
l,4-Dioxane-I~ 1,CDioxane-Br~ 1,3-I>ioxane-Iz
8.0
+1.88 f 0.08
+6.62&0.08
4.5
-2.84A0.08 $2.35 =t0.15
+7.00f0.10
8.4
for Ur2 and 1,s dioxane because the bromine was rapidly decolorized slightly above room temperature. The results arc summarized in Table I; the difference in heats of formation of the two iodine complexes is consistent with the fact that one is solid and the other liquid. Calorimetric Studies. Heats of solution were measured at 25” in a 100-ml. capacity dewar calorimeter of the kind which has been described previously.* 1,4Dioxane was purified by refluxing with dilute HC1, dried by treatment with KOH, and fractionated from sodiumQ; 1,3-dioxane is not so easily purified and was simply fractionated. For the determination of the heats of solution of the solid complexes of 1,4-dioxane with 1, and Rrz, fresh samples were prepared before each run by the method described by Rheinholdt and Boy2; the iodine and bromine used were of analytical grade. The quantities were adjusted so that the final solution was the same, independent of whether the heat of solution of the halogen or its complex had been measured; three determinations were made for each heat of solution, and all determinations fell within the limits quoted. No satisfactory heat of solution of Br2 in 1,3-dioxanc could be determined because a slow reaction occurrctl. A summary of the results is given in Table 11. In the final column of Table 11, we give the calculated heat of reaction 1 based on the calorimetric measurements quotcd and the latent heats of vaporization of the halogenslo and 1,4-dioxane.” It is seen that the agrcemcnt with the figures obtained in Table I is quite satisfactory and strongly supports our choice of expressioii for the equilibriuin constant; it also gives some confidence in our estimate for the heat of formation of the liquid coiiiplex between iodine and 1,3dioxane. Acknowledgment. Wc wish to thank Dr. €1. A . Skinner and Dr. G. I’ilcher for the use of their calorimetric equipment. (8) H. 0. Pritchard and H . A. Bkinrrr, J . C h m i . 9oc.. 272 (1960). (9) E. Eigenberger, J . praltt. Chem., 130, 76 (1931). (10) National Bureau of Standards Circular 600.
Calcd. AH for reaction 1 at 25O, koal./mole
-28.6 f0.3 - 2 6 . 2 i0 . 3
(11) Much of the relevant work on 1.4dioxane in the literature has been done with dioxane which has been simply treated with sodium’? or calcium chloride’s and, unfortunately, could not be more than about 97% pure. However, in their vapor pressure determinations, Crenshaw, et al.,” did rigorously purify their dioxane. and for this reason we derive our latent heat of vaporization from their data, giving XvSp = 9.0 0.1 kcal. a t 26‘. (12) A. F. Gallhaugher and H . Hibbert, J . A m . Chem. SOC.,59, 2421 (1937). (13) C. G. Vinsen and J. J. Morton, J . Chem. Eng. Data, 8 , 74 (1963). (14) J. L. Crenshaw, A. C. Cope, N. Finkelstein, and R. Rogers, J . Am. Chem. SOC.,60, 2308 (1938).
*
Surface Areas from the V / n Ratio for Marine Sediments by W. H. Slabaugh and A. D. St,ump Department of Chemistry and Department of Oceanography, Oregon State University. Coruallis, Oregon (Receiued November 4! 1963)
We have recently made a series of determinations of the B.E.T. (Brunauer-Emmett-Teller) surface areas and pore-size distribution of marine sediments obtained from the continental terrace off the coast of Oregon by the research vessel hcona of Oregon State University. Our results showed good agreement bctween surface areas calculated by the T3.E.T. method and by the use of the ratio V l n , where V denotes the volume absorbed a t a certain relative pressure and n is the statistical number of molecular layers calculated from Frenkel-Halsey-Hill equation’s2 for nitrogen adsorption.
Experimental Apparatus. The adsorption isotherms were determined gravimetrically on a conventional gravimetric adsorption balance with quartz helical springs. The balance had a sensitivity of about 1 mm. extension (1) C. Pierce, J . Phya. Chem.. 63, 1076 (1969). (2) C. Pierce, ibid., 64, 1184 (1960).
V o l u m 68,Number 6
M a y , 19134
