1917
NOTES
Sept., 1963
TABLE I CHARACTERIZATION OF CATALYST Support size (microscope) b Sample a Above 149 149-74 Sieve size ( p ) Av. diameter ( p ) 280 76 Extreme dimensions (y ) from 10 measurements 440-180 116-70
74-62 66 80-60
d 6244 38 60-20
e Below 44 24 38-12
Platinum content (iodometric method) 0.53 0.58
0.58
0.61
0.68
3.32 1.89 0.80
3.57 1.95 0.80
Wt. % platinum”
C
Support density Skeleton density (g./cc.) ... 3.44 Bulk density (g./cc.) ... 1.87 ... Wt. sample/bulk volume (g./cc.) ... 0.76 ... Platinum-?-alumina Form Correlation distance a (small angle X-ray scattering) a (A.1 22.9 22.0 Estimated error (%) ... 5 3
...
...
23.2 4
19.7 4
Transmission measurements 28.9 31.5 29.7 40.7 Excess extinction coefficientb(cm.-1) ... 4180 4293 Wave length of max. transmission (A:) ... 4193 4180 Unsifted sample has a platinum content of 0.61 wt. %. Platinum content determined bly Esso with X-ray spectroscopy = 0.614 wt. %. Excesn extinction coefficient Text(excess) = yexo(Pt-alumina) - Text (alumina). Text(alumina) L% 1.4 em.-’ and remained unchanged when heated at 600” in air.
TABLE I1 GROWTH OF SUPPORTED PLATINUM AT 600’ (f15’) IN AIR Sample
b C
d b C
b C
b C
b C C
b C
C C
b
Time (hr.)
Dilution (wt. cat./wt. cat. blank)
Excesb extinction ooeBcientb (om.-*)
0.128 .242 .09 .0884 .1352 .128 .242 .128 .242 .128 .242 .1352 .0884 .1352 .242 .242
33,33.7 33,33.3 31,30.4 45 53 62,61 62 61.5 62 62.5 65 68.5 62 66 67 64
+
0 0 0 2.1 2.1 6.6 6.6 6.8 6.8 20.8 20.8 42.3 99.8 99.8 192.1 381.7 122.3 hr. plus 20 hr. a t 700”
.0884
61.5
results of Herrmann, et al., Table IV14and Johnson, aZ.,6 where tlie soluble platinum after heat treatment is about 30 w t , yo. Acknowledgment.-The measurements in this report were performed by Mr. R. R. Hampton at Cornel1 University. The author wishes t o thank the Esso Engineering and Research Company for financial support and Professor IP. Debye for many inspiring discussions.
et
ERRORS I N “DROP” CALORIMETRY DUE TO SAMPLE CONTAINER TRANSITIONS BYD. C. GINNINGS National BuTeau of Standards, Washingtot&,D. C.
Received March 9, 1963
I n using the “drop” method for measurements of enthalpy a t high temperatures, usually it is assumed that the sample container gives up the same heat to the
calorimeter when it is empty as when it contains a sample. If there are no transitions in the container material, this assumption is reasonably justified, because the effect of the sample on the heat lost from tlie container during the drop frequently is much smaller than can be detected. However, if there are transitions in the container material, this assumption may not be valid, because the time available for the transition in a certain “critical” temperature range below the transition temperature depends on the presence arid thermal properties of a sample. The transformation in the container material will be more complete when the sample has a high thermal conductivity and heat capacity and is in good thermal contact with the container. Using this logic, it is concluded that if there is an error of this kind, the “apparent” enthalpy of the sample above the transition temperature will be larger than the true enthalpy of the sample. The magnitude of this error depends on the difference in container material transformations for the empty and full experiments. With this in mind, an examination was made of thwe sets of published results, all using the “drop” method but each using a different container material having a transition in the experimental temperature range. In the experiments with aluminum oxide up to 90Oo,’ a Kichrome V container was used which has a small transition between 600 and 700°.2 The results of these experiments are shown in Fig. 1 as experimental deviations from an equation which was fitted by the method of least squares to the experimental enthalpy values. The observed values (heavy line) of enthalpy, H , show an abrupt increase in slope between 600 and 700’ as would be predicted if there is a transition error as described. The enthalpy values above 600’ were arbitrarily “corrected” for this change of about 1 j./g, (equivalent to 16 j . in the experiments), and the same form of equation was fitted by least squares to these (1) G. T. Furukawa, T. I). Douglas, R. E. MoCoskey, and D. C. Ginnings, J . Res. Nail. Bur. Std., 67, 67 (1956). (2) T. B. Douglas a n d J. L. Dever, ibid., 64, 15 (1955).
1918
Vol. 67
NOTES
y - ,
ALUMIA!UM OXIDE ENTHALPY EQUATION
1.01% AH/At
8 - \
-. x-.
LEAD ENTHALPY EQUATION
...
X----
L .
i
.-
‘x
MAGNESIUM ALLOY ENTHALPY EQUATION? f 5% AH/At
”’d *.‘x\*
300
400
500
600
\.x- - -x ,
700
800
TEMPERATURE (ti. “C,
Fig. 1.-Correlation of enthalpy experiments: 0, observed ehthalpy ( H ) deviations from equation; X, corrected observed H deviations from equation; 0, observed A H / A t deviations from equation; a, corrected observed A H / A t deviations from equation.
experiments4 which used the drop method with a sample container of stainless steel No. 430 which has a transition between 650 and 700’. I n this case, as with lead, the sample of magnesium alloy was liquid in the region of the transition. Figure 1 shows the deviations of the observed values from the published equation. Again, there is a significant increase in enthalpy above the transition temperature, this time about 7.7 j./g. (corresponding to about 75 j. in the experiment). The figure indicates the true values of heat capacity of the liquid are about 5% lower over most of the liquid range than those given by the published equation. I n the above experiments, the heat correction made arbitrarily to the observed enthalpy values above the transition temperature can represent only a fraction of the total heat of transition of the container material. It is believed that the three examples given should serve as a warning to experimenters who are using the drop method with sample containers having transitions. With precise calorimetry, it seems better to avoid these materials, if possible. (4) R. A. MoDonaldand D. R. Stull, J . Chem. Eng. Data, 6,609 (1961).
T H E SOLUBlLITY OF CARBON TETRAFLUORIDE IN
PERFLCOROMETHYLCYCLOHEXANE BY L. W. REEVESAND J. H. HILDEBRAND Department of Chemistry, University of California, Berkeley California
Received March 16, 1068
values. A “corrected” enthalpy equation is obtained which is believed to better represent the true enthalpy of aluminum oxide. This “correctedJ’ enthalpy equation is
Ht - Hp
=
1.4512422 - 2.0399 (10-5)t2 462.840 log [t
+ 273.15)/273.15]
where H is in joules per gram, referred to O’, and t is in degrees C. In the figure are shown the observed values of AH/At compared to the values calculated from this “corrected” equation for aluminum oxide. It is obvious that the value at 650” (representing the range 600-700’) is unique, departing from the equation by the equivalent of 1% as compared to a nzaxilnum deviation of all the other points of about 0.1%. I n fact, the figure indicates that the form of equation used does not fit the observed data within the precision of the data. The “corrected” equation gives values of heat capacity which are about 0.3y0 lower at 900’ than the uncorrected equation. In another series of experiments with lead,3 a stainless steel (KO.446) sample container was used which has a transition between 500 and 600°.2 Figure 1 shows a significant increase in apparent enthalpy of liquid lead above this transition temperature. This increase is about 0.2 j,,/g,, corresponding to 9 j . in the experiment. The figure shows the deviations of observed values of A H l A t from those calculated from the published equation. The five values above 400’ now are quite consistent, giving “corrected” heat capacities about 0.3% lower (0.00036 j./g, ‘C.) than those from the published equation. ,4n investigation was also made of a third series of (3) T,B . Douglas and J. L. Dever, J. A m , Chem. SOC.,76, 4824 (1964).
The results published herein were obtained immediately after those for the solubility of argonJLwhich were published by us in 1957. They presented unusual features and were withheld at the time pending the accumulation of further data on the solubility of this gas. It is appropriate to release them now in the issue of this journal that presents a comprehensive survey of the solubility of CF, and SFBin a number of nonpolar solvents by Archer and Hildebranda2 The apparatus, materials, and procedure are described fully in the argon paper. The results are given in Table I. SOLUBILITY O F
TABLE I CF, AT 1 ATM.PARTIAL P R E S S C R E FRACTION 22
Solvent
T,OK.
GHie
(312.86) 298.07 286.75 272.46 262.88 298.16
AS hf0LE
104~~
(93.1) 1C5.9 119.1 139.8 156.8 105.3(calcd.)
Plots of log x2 us. log T in the case of accurate measurements give straight lines whose slope, multiplied by R = 1.986 cal. deg.-l- mole-lJ give the partial molal entropy of solution, Sz - Szg. The point enclosed in parentheses falls slightly off a straight line through the other points, and has been neglected in getting the calculated value of the entropy, -6.2 e.u. Acknowledgment.-We gratefully acknowledge the support of the Atomic Energy Commission. (1) L. W. Reeves and J. H. Hildebrand, J . Am. Chem. Soc., 79, 1313 (1967). (2) G.Aroher and J. H. Hildebrand, J . Phya. Chem., 67,1830 (1983).
