1784
Vol. 64
NOTES I
I
I
I
I
listed as series B in Table I. Other materiab and solution8 were prepared as described earlier.6 During series B measurements the usual troubles with dioxane, due to development of acidity by oxidation, were experienced even though the solutions were protected rather carefully by a nitrogen atmosphere at all times. The solvent was repurified and the solutions were made up freshly several times during the study. Procedures.-The measurements have been described in detail previously.*~6 In sene8 B the pH-meter calibration and estimation of the ion product of the medium by titration of perchloric acid with potassium hydroxide were repeated before every determination of the acidity of dibenzoylmethane. Values of U H in theequation -log [H+]= meter reading log UH varied slightly from one series of experiments t o another, despite the fact that the electrodes were carefully conditioned and that the meter was always set to read H 7.00 in the appropriate, standard Beckman buffer. Smalfvariations also were noted in the apparent values of -log [H+][OH-1: The average value from Fix runs was 16.77 =k 0.16. The apparent values for the individual runs ?ere used in calculating the values of -log Q.4 reported in rable I, and the values reported in Table I1 were calculated using a single, average value of 16.9 obtained by a series of calibration measurements in the early phase of the study.
+
10.5
I
I
-0.30 -0.20 -0.10
I
I
0
+0.10
I
I
+0.20 +0.30 +0.40
d. Fig. 1.-Hammett plots of ionization constants of dibenzoylmethanes and benzoic acids in 75% dioxane-water a t 25.’: dibenzoylmethanes, w m , a = 0.00; benzoic a = 5.00. acids, 0-0-0,
plot of the data, and the values for the ionization of the corresponding benzoic acids are also plotted for purpose of comparison. The value of p for ionization of the diaroylmethanes is 2.63 (7 = 0.33) and that for the benzoic acids is 1.28 (r = 0.15). Since the diketones are each symmetrically di-substituted, the effect per substituent is the same for the two reactions with a precision better than those of the fits to the Hammett equation. The implication is that interactions between the substituents and the functional groups are similar in the two series of compounds. This, in turn, implies that internal interactions within the functional groups are probably much more important than are conjugated interactions with the adjacent benzene rings. If the effects of substituents in the diketones were cumulative, one would expect that a single substituent would have half the effect of two. Actually) a single p-methoxygroup decreases the acidity constant by 0.62 log unit and a second methoxy group gives an additional decrease of only 0.20unit. Mecke and Funkg have analyzed the infrared spectrum of the enol of acetylacetone and have concluded that the compound is symmetrical. This is aview with which we have agreed, but present data suggest that p-methoxydibenzoylmethane is stabilized to some extent by interactions such as those shown in I. This may induce significant asymmetry in the G O bonds and in the position of the bridge hydrogen. H -0
‘0
I
Experimental Materials.-Preparation and purification of the diaroylmethanes were described earlier.6 A new sample of dibenzoylmethane waa used and purified for the measurements (9) R. Mecke and E. Funk, 2. Elektrochcm., 60, 1124 (1956).
Acknowledgment.-The later part of this study was supported by a grant from the Film Department of the du Pont Company. T H E REACTION BETWEEN BzO3(1) AND C(S): HEAT OF FORMATION OF B202(g)* BY PETERRENTZEPIS,DAVIDWHITEAND PATRICK N. WALSH
Cmogente Laborafory, Department of Chemistry, The Ohio State Uniuersaty, Columbus 10 Ohio Recezved June 16, 1960
During the past few years it has been demonstrated that, under reducing conditions, boron suboxide, Bz02, is an important species in the vapor in the B-O system. Inghram, et d . , l and Scheer2 identified the species and measured its partial pressure in the system B(s) B2O8(1),the former by mass spectrometric and the latter by torsion effusion techniques. Searcy and Myers3 measured the effusion of B202and Mg vapors from a mixture of MgO and B at elevated temperature. Evans, et u Z . , ~ in a review of the properties of B-0-H vapor species, concluded that the results of these three investigations lead to the following values of the heat of formation (AH! (f)) of B202(g),-103.6 f 5, -109.6 ZJ= 6 and -105 i 5 kcal./mole, respectively. From the above figures a “best” value of -105 f 5 kcal./mole was chosen. Scheer2 has suggested that the attainment of equilibrium in the reaction of B203(1)with B(s) in his experiments and those of Inghram, et al., might have been impeded by the formation of a condensed polymer (BO),, with the result that the observed pressures of B202may be lower than the true equilibrium pressures. Similarly, Searcy and Myers3 suggest that poor contact between their solid react-
+
*This work was supported by the Office of Naval Research Washington, D. C. (1) M.G. Inghram, R. F. Porter and W. A. Chupka, J . Chem. Phys.. 25, 498 (1956). (2) M. D.Scheer, THISJOURNAL, 62, 490 (1958). (3) A. W. Searcy and C. E. Myers, ebid.. 61, 957 (1957). (4) W. H. Evans, E. J. Prosen and D. D. Wagman in “Thermodynamic and Transport Properties of Gases, Liquids, and Solids,” Y.5. Touloukian (Editor), RlcGraw-Hill Book Co., New York. N. Y., 1959.
Nov. , 1960
NOTES
ants may have resulted in a lowering of the partial pressures of the gaseous products. Hence, all the above investigations may have led to an upper limit for the heat of formation of B202(g). In order to avoid some of the difficultiesmentioned above, and establish possibly a lower limit for the heat of formation of B202(g),the reaction of BzOa(1) with carbon has been studied. B203-C Reaction.-Consider the vaporization of B2O3(1),to which carbon has been added, from a carbon Knudsen cell, in the temperature range 13501650°K. Only two reactions are of importance6
1785
Results It was found that, a t temperatures below 1550" K., only very small quantities of B202(g) were
evolved from the effusion cell for any appreciable rate of weight loss (5-20 mg./hr.). This was readily detectable because BzOzcondenses to a deep brown amorphous solid on the cold walls of the condenser. Within experimental error, therefore, the only gas effusing from the cell below 1550°K. is CO, and reaction 1is predominant. In Table I, the observed CO pressures are compared with those calculated for this reaction. The free energy functions used in the calculations were taken from these several BzOdl) 3C(s) +3CO(g) 2B(s) (1) sources: B(s) and C(s), the compilation of Stull BzOd1) C(S) +BdMg) CO(g) (2) and Sinkes; B203(1),reference 4; CO(g), free energy At the lowest temperatures, assuming a reason- of formatione combined with free energy functions able value for the heat of formation of B202(g), of C(s) and 02(g).* reaction 1 should predominate. As the equilibrium constant for this reaction can be calculated TABLE I from available thermodynamic data, one can de- COMPARISON OF EXPERIMENTAL AND CALCULATED CO termine, (a) under what conditions equilibrium PRESSURES FOR REACTION B20a(l) + 3C(s) + 3CO(g) can be established in the Knudsen cell, and (b) 2Ws) whether any new condensed phases (say (BO),), Pco, atm. X 10:. AFO, kcal. Pco, a b . X 101, calcd. exptl. calcd. T,OK. which would appreciably affect the equilibrium 58.1 0.836 1376 1.03 pressure of CO, are formed. If the reaction then is 57.5 0.964 1389 1.08 studied a t higher temperatures, where reaction 2 48.7 3.82 1468 5.24 cannot be neglected, it can be shown that 48.2 4.12 3.81 1473 RT In ( J ' c o / P B=~ flof(BzOz(g)) AFor(CO(g)) -
+ +
+ +
+
-
1/3AFodBz0d1)) (3)
where AFofrepresents the standard free energy of formation a t temperature T. The heat of formation of Bz02(g)thus can be obtained from the ratio of pressures, using the appropriate thermodynamic data. Experimental The effusion experiments were performed in an apparatus similar to that described by Skinner, Edwards and Johnston6 with a demountable silica gel trap cooled to 77'K. in the vacuum line when CO samples were collected. For rate of effusion measurements, the system was pumped continuously and no absorbent was used in the trap. The temperature in the induction-heated cells was measured by sighting through the effusion orifice with a calibrated Leeds and Northrup optical pyrometer. The boric oxide was prepared by vacuum dehydration, at approximately 1100", of reagent grade Matheson, Coleman and Bell boric acid. The graphite effusion cells and lids, and the graphite particles placed in the cell, were made from material of spectroscopic purity obtained from the National Carbon Company. The graphite was thoroughly degassed at 2000' in vacuo before being used in any of the experiments. The effusion cella were short cylinders, with internal diameter approximately 0.3 cm. and height approximately 1.1 cm. The orifice was drilled axially through the tight-fitting lid. When the ratio of orifice area to area of the evaporating surface was greater than 0.008, i t waa found that the CO pressurea (calculated by the method of Motzfeldt') were not independent of the orifice area, but decreased with increasing area. This probably came from a reduction in the effective area of the evaporating surface caused by floating graphite particles. The results in Table I and I1 were obtained using orifice-to-cell area ratios of about 0.005. The diameter of the orifice was, on the average, approximately 0.02 cm. and its length approximately 0.04 cm. (5) BlOl(p) comprises a maximum of 10% of the boron-containing species in the vapor at any temperature. X-Ray analysis of the solid phase disclosed no B4C or B-C solid solution. Even if these were present, their free energies of formation are so small that they would have a negligible effect on the observed pressures. (6) G. B. Skinner, J. W. Edwards and H. L. Johnston, J . Am. Chsm. Soc.. 73, 174 (1951). (7) K. Motzfeldt, THISJOUBNALI S . 139 (1965).
1476 1525 1563 1593 1603
8.60 23.9 41.4 71.1 73.4
47.8 42.4 38.1 34.9 33.8
4.36 9.43 16.8 25.4 29.0
It is evident from Table I that there is good agreement between the experimental and calculated CO pressures a t the lower temperatures. At higher temperatures, the apparent divergeme probably arises from the fact that the equation for molecular flow7 used in calculating the pressure is no longer applicable. The magnitude of the error arising from viscous flow in these experiments can be approximated from the relationship between a/L and F/Ft tabulated by Dushman.'o If F is the actual flow through the orifice, Ft, the calculated molecular flow, a, the orifice radius and L, the mean free path a t the average pressure across the orifice, then for a / L = 1.0, which corresponds approximately to the case when the CO pressure is 5 X atm., F/Ft = 1.00. For a / L = 10 (CO pressure approximately 5 X loy2atm.), F/Ft = 2.29. Considering the approximations involved in the Dushman table and the uncertainties inherent in Knudsen effusion studies, the experimental data appear consistent over the entire temperature range. It is therefore concluded that no significant lowering of tfheactivity of B203(1)takes place. It is only a t temperatures in excess of 1600°K. that the vapor contains enough B202to permit a reliable determination of the ratio of the partial pressure of CO to that of BzO2. It is evident from the foregoing discussion, however, that the classical (8) D. R. Stull and G. C. Sinke, "Thermodynamic Properties of the Elements," American Chemical Society, Washington, D. C., 1956. (9) J. P. Coughlin, Bur. Mines Bull., 542 (1954). (10) S. Duphman, "Vacuum Technique," John Wiley and Sons, Inc., New York. N. Y.,1949, p. 114.
KOTES
1786
l7Ol. 64
TABLE I1 HEATOF FORMATION OF BZOz(g) T,OK.
Initial BzOa, mmoles
Total CO evolved, mmoles
Total BaOn evolved, mmoles
Total wt. loss, mg. exptl.
Total wt. loss, mg. calcd.
Pea/
1603 1616 1626 1651 1656
1.763 2.868 2.191 5.270 3.471
5.14 8.41 6.40 16.19 9.86
0.075 .095 .085 .31 .27
148.0 240.6 183.9 442.0 290.8
148.3 241.3 184.2 441.7 290.6
69 89 75 49 37
AHaOf kcal./mole (BaOz),
PBaOc
-110.6 -109.6 -110.1 -111.3
-112.1
-4~.-110.7 i 1 . 5
Knudsen technique cannot be applied under these conditions; a different technique therefore was employed. A known weight of B203(1) was introduced into the cell, along with some graphite particles, and the reaction allowed to go to completion a t a constant temperature. The CO evolved was collected on a silica gel bed which previously had been cycled with CO. The amount of CO was determined by desorption into calibrated volumes. On the assumption that only reactions 1 and 2 were involved in the depletion of the B203, the amount of Bz02 evolved was computed. The experimental data and the derived heat of formation of B202(g)are given in Table 11. From a comparison of columns 5 and 6, it would appear that the assumption that reactions 1 and 2 are predominant in the depletion of the liquid boric oxide is justified. The ratio of P C O I P B ~ O was ~. calculated from columns 3 and 4 assuming viscous flow. I n calculating the heat of formation, the thermal functions cited above were used for B20a(l), B(s), C(s) and CO(g). For BzOz(g), those reported by White, et al., were emp1oyed.l' Discussion From the nature of the experiment, it is felt that the heat of formation of gaseous BzOz given in the last column of Table I1 probably represents a lower limit. The experimental uncertainties are unquestionably large in this type of determination, because of the large Pco/PB~o~ ratios and the small quantities of BzOz evolved. The lower limit is suggested because the largest uncertainties arise from (a) the extent of incomplete desorption of the CO, which is estimated as 0.002 mmole, and (b) the presence of B203 in the vapor, which, though small, may appreciably diminish the Pco/PB~o~ ratio. In order that a comparison of this work with that of the previous investigators be meaningful, it is essential that the same values for the thermal functions of each of the species should be used throughout. The heat of formation of B202(g) has therefore been recomputed from the experimental data of references 1, 2 and 3. The free energies of reaction 4 reported by Inghram, Porter and Chupkal have been used in conjunc2/3B20dg)
+ 2/3B(s) +BzOdg)
(4)
tion with the thermodynamic data listed to calculate AH: of formation of BzOz(g). Free energy functions of t,hevapor species were takenfromwhite, et uZ.,11 while the compilation of Stull and Sinke was used for boron. Together with the heat of for(11) D. White D. E. hfann, P. N. Walsh and A. Sommer, J . Chem. P h y s . 32 481 (1360).
mation of B203(g),AH; = -209.5 i 2 kcal./'mole,lz the data yield AHO(f) = -111.6 i 3 kcal./mole for BzOz(g). Scheer2 studied reaction 5 2/3BzOa(l)
+ 2/3B(s) +BzOdg)
(5)
by the torsion effusion method. The slope of a Clausius-Clapeyron plot of his data yields - 108.2 kcal./mole for AH;(f) of BzOz(g) while applications of the third law, using free energy functions already discussed, give a value of - 112.7 i 4 kcal./mole. Despite a small but definite variation of the "third law" heat with temperature, the latter value is probably more reliable. It is interesting to note the good agreement between the two sets of investigators, despite the factor of 6 to 8 difference in their reported pressures at a given temperature. The probable explanation is that the mass spectrometric technique' gives better relative than absolute pressures. Searcy and Myers3 studied the reaction 2I\IgO(s)
+ 2B(s) + 2Ng(g) + B2Odg)
(6)
by conventional Knudsen effusion procedures. Their reported pressures, which are not in accord with the stoichiometry of reaction 6 corrected for differences in rates of flow of the two vapor species, have been recomputed. The thermodynamic properties of B(s) and B202(g) were taken from the sources previously cited; those of Mg(g) were derived from Stull and Sinke.8 The free energy functions of MgO were calculated from the high temperature heat content data of Victor and Douglas.16 These differ little from those computed using Coughlin's table.9 The average of all experimental points yields - 105.7 f 5 kcal./mole for the heat of formation of BzOz(g). The authors, however, point out3 that poor contact between the solid reactants may be responsible for the rather large average deviation of their heats, and suggest that the single determination carried out with a compacted sample and very small effusion orifice probably represents a more reliable value than the average. From this point, we calculate a heat of formation of B202(g) of -111.9 kcal./mole a t 0°K. (12) Computed as follows: The vapor pressure data"814 yield 94.4 f 2 kcal./mole for the heat of Sublimation of crystalline BIO, a t O'K. when the free energy functions of ref. 11 are used for BaOsk). This value p a s combined with AH:(f) of crystalline BaOa'S to obtain AH;(f) of BzOa(g). (13) R. Speiser, 9. Naiditch and H. L. Johnston, J . Ani. Chcm. Soc., 72, 2578 (14) (a) M. D. Scheer, THISJOCRNAL,61, 1184 (1957); (b) J. R. Soulen, P . Sthapitanonda and J. L. Margrave, ibid.. 69, 132 (15) E. J. Prosen, W. H. Johnson and F. Y. Pergiel, J . Research, Natl. BUT.Standards, 62, 43 (16) Unpublished data by A. C. Victor and T. B.Douglas.
(1950).
(1955).
(1959).
NOTES
Nov., 1960
It is evident that all the values, Some of which should be lowerand upper limits, agree within experimental error. This strongly indicates that the average, Z I ~ Z . ,- 111.7 f 2 kcal./mole, is a reliable value for the heat of formation of BzOz(g) at 0°K. THERMOCHEMISTRY OF SULFUR TETRAFLUORIDE
1787
tion. HF was analyzed by thorium nitrate titration of fluoride ion. It was necessary to determine sulfite a t various time intervals because of the slow autooxidation to the initial sos- was taken that found.by extrapolation of the observed linear relation to zero time of exposure of the bomb contents to air. Sulfur was determined indirectly from the H F , SOS’ and/or HrS values; preliminary experiments showed that the weighed quantity of free sulfur was in good agreement with the amount e\pected from the weighed initial quantities of SF4 and H?.
Results and Discussion
BYJ. D. VAUGHNAND E. L. MUETTERTIES
The heat effects accompanying the bomb reaction included the heat Of hydrogenation Of SF4, Contribution h’o. 657 from the Central Research DepaTtment Expen’. mental Statzon, E. I . du Pant de Nemours and Company, Wilmzngton, the heat of condensation of the product HF, and Delaware the heat of formation of HzS (when H2 was in exReceived June 88,1960 cess). To compute the heat of condensation of HF, Sulfur tetrafluoride is strikingly reactive in com- the extent of condensation mas estimated. The parison to sulfur hexafluoride and we have sought saturation vapor pressure of liquid HF a t 298OK. a comparison of their thermochemical properties. is 908 mm.,2 corresponding to an average molecuTost and Claussenl have determined the heat of lar weight of 60.6 g. ; the vapor density of H F under formation of SFe by direct reaction of sulfur and these conditions is given by d = J f P / R T = 0,00295 fluorine, but nothing has been reported on the heat g./ml. The weight of H F in the gaseous state in Of formation of SF4. I n the investigation described the bomb of 360 m]. is, therefore, 1.062 g., and the herein, the heat of formation of SF4 was estimated fraction of H F that becomes liquid (1 - 1.062/W), calorimetrically by use of the hydrogenation of where is the total weight of HF found in the SFc bomb a t the completion of the hydrogenation. The SF,(g) 2Hz(g) +4HF(1) S(S) heat evolved due to condensation is given by the SFa(g) 3Hz(g) +4HF(1) HzS(g) product of the fraction of liquid HF and the heat of per gram HF*3 Only the former reaction occurred with SF4 in stoichiometric excess, but both occurred with Hz The heat of famation of HF of ~ o l e c u l aweight r 60.6 g. was required for computation of the heat of in excess. hydrogenation of SF4. Simons and Hildebrand2 Experimental indicated that, the apparent molecular weight is Purity of Sulfur Tetrafluoride.-SF4 was passed through an oxy-hydrogen flame and then fluorine was determined due to incomplete polymerization of HF to the by thorium nitrate titration of fluoride ion and sulfur by hexamer (HF)G. Long, Hildebrand and Morrel14 precipitation as barium sulfate. The results of two analy- gave the heat of polymerization of HF as -6800 ses are given in Table I. ca1./20 g. A molecular weight of 60.6 corresponds to 50.5%polymerization, such that the heat evolved TABLE I because of polymerization at 298’K. is -3434 cal. CHEMICAL ANALYSIS OF SF, The heat of formation of H F under the experimental Sample 8, % F,% %S+%F conditions then equals the heat of formation of 1 31.06 69.32 100.38 monomer plus the heat evolved in polymerization, 2 30.82 69.71 53 that is, -67.6 kcalJ20 g. The heat of formation Theoretical 29.60 70.40 100.00 of HF monomer, as well as the heat of formation of Mass spectrophotometric analysis gave the following HzS, were taken from the Xational Bureau of results: 95.6% SFI, 3.8% SOFz, 0.3% SiF4, 0.2% CS?, Standards Circular 500.3 and 0.270 N?. Preconditioning of the mass spectrometer The experimental results are summarized in led to higher values (>98%) for the SFa content. The presence of SOF? was assumed to he due to hydrolysis Table 11. Consideration was given t o the following within the mass s ectrometer in view of the implied ahsence uncertainties: solubility of Hz, H2S and SF4 in of other elementa! species in the rhemical analyses. condensed HF, polymerization of HF in the vapor Calorimeter.-A Parr double-valve isothermal oxygen bomb calorimeter was adapted for the hydrogenation phase, the effect of final reaction temperature other reaction by replacement of the rubber sealing ring and the than 298’K., calibration error, extent of automica electrode insulator with “Teflon.” A hot platinum oxidation of SO3- to so4-, and purity of SF4. wire was used for initiation of the calorimeter reaction both First approximation estimates of these uncertainfor the calibration combustion and the hydrogenation of SF,. The calorimeter was calibrated by combustion of ties indicated them to be smaller than the average benzoic acid; thirteen such combustions gave an average deviation from the mean value of the standard calorimeter heat capacity of 2488 cal./deg. heat of formation of SF,. Because of the difficulty General Procedure.-The bomb was flushed several of precise estimation of these uncertainties, and times with hydrogen before SF4 was added under pressure. In three of the experiments, SF4 was added in excess, and the uncertainties arising from the extreme nonin three others, Hz was added in excess. At completion ideal behavior of HF in vapor condensed phases, of the calorimeter reaction, 2.5 M NaOH was added to the the over-all uncertainty is taken as the sum of the calorimeter bomb contents under external hydrogen pressure. average deviation from the mean ( 2 . 3 kcal./mole) Evcess unreacted SF, was determined as the hydrolysis product, sulfite ion, by titration with standard iodine solu- and the calibration error (0.3 kcal. ’mole).
w
+ +
+ +
tion. HsS was determined by addition of excess iodine solution and back-titration with standard thiosulfate solu(1) D. M. Yost and W. H. Claussen, J . Am. Chsm. SOL, 56, 885 (1933).
(2) J. H. Simons and J. H. Hlldebrand, ibid., 46, 2185 (1924). (3) N.B.S. Circular 500, Washington, D. C., 1952. (4) R. W. Long, J. H. lhldebrand and W.E.Morrell, 3. Am. Chrm. Soe., 66, 182 (1943).
