Dec., 1961
HEATSOF DECOMPOSITION OF HIGHERBORON HYDRIDES
bond ruptures rather than as a simultaneous elongation of two opposite carbon-carbon bonds in the ring. Thus a primary reversible ring opening might be easier because of the extra-ring strain energy, but the resulting species H,C=C-CHz / H~c=CH~
might be more stable than the tetramethylene analog and would require greater energy for complete dissociation to allene and ethylene, giving a larger net activation energy. A check of this point of view might he provided by studies to see if a labeled exo-carbon atom moved into the ring a t a rate comparable to or faster than the decomposition
2173
reaction. Examination of l,&dideuterio or dimethylcyclobutane for cis-trans isomerization reaction also would be valuable. On the other hand, increased energy for a single transition state involving ring elongation could be rationalized by the non-linear nature of the allene fragment in the transition state. It would less resemble the linear product molecule than would the non-linear partially formed propylene resemble the non-linear product in the methylcyclobutane decomposition. Thus the activation energy for methylenecyclobutane decomposition could be greater even though the ring strain is also slightly larger than for the alkyl cyclobutanes.
THE HEATS OF DECOMPOSITION OF SOME HIGHER BORON HYDRIDES1 BY STUART R. GUNNAND LEROYG. GREEN University of California, Lawrence Radiation Laboratory, Livermore, California Receited J u n e 2.9, 1961
Heats of decomposition of B4Hloj BbHs, B5fi.11 and B6Hlo have been determined by explosion of mixtures with SbH3. Bond energies of boron hydrides are derived and diecussed.
Introduction The heats of decomposition of B2H6,2 B5H2 and BlaH143 have been determined by a method involving pyrolysis of the compound in a furnace enclosed in a calorimeter. McCoy and Bauer4 have derived the heat of formation of BH3 from indirect measurements of the heat of dissociation of BzH6. There appear to be no determinations of heats of formation of other boron hydrides; these data are of interest for an understanding of the energetics of interconversion of boron hydrides and the nature of bonding in these unusual compounds. During a recent study of the heat of decomposition of several gaseous hydrides,s we detjermined the heat of decomposition of diborane, obtaining a value in reasonable agreement with that of Prosen. In the present work, we have extended this technique to B4Hl0,B5H9, &Ell and B6Hlo. Experimental BpHlO and B5H9 were taken from stocks available in the laboratory. B,H11 was prepared by the reaction of B2Hs and B4Hlo at 100" for two minutes. B6H10 was prepared by decomposition of the methyl ether adduct of BsHl1.6 Purifications were performed by bulb-to-bulb distillation on the vacuum line. Infrared spectra were used to follow the course of purification; final spectra of B4Hlo,BjHBand B5Hll were in good agreement with those of McCarty, et aZ.7 Other boron hydrides were not detectable. The final spec(1) This work was performed under the auspices of the U. R. Atomic Energy Commission. (2) E. J. Prosen, W. €1.Johnson and F. Y. Pergiel, J . Research Xatl. Bur. Standards, 61, 247 (1958). (3) W. H. Johnson, M. K. Kilday and E. J. Prosen, ibid., 64A, 521 (1960). (4) R. E. McCoy ar.d S. H. Bauer, J . Am. Chem. Soc., 28, 2061 (1956). ( 5 ) S. R. Gunn and L. G. Green, J . Phgs. Chem., 66, 779 (1991). (6) XI. W. Forsyth, W. V. Hough, MI.D. Ford, G. T. Hefferan and L. J. Edwards, National Meeting American Chemical Society, Spring, 1959. We are indebted to Dr. Hough for providing further experiniental details. (7) L. V. McCarty, G . C. Smith and R. 9. McDonald, A n a l . Chenr., 26, 1024 (1954).
trum of B6H:o agreed with an available unpublished spectrum except for the presence of a few per cent. of B51L. Purities of B4H10, BjH, and B5Hll were checked using a melting-point apparatus similar to that of Skau.8 A technique was used which permitted approximate evaluation of the heat of fusion and the purity from variation of the melting temperature with fraction melted. The copper-constantan thermocouple was calibrated a t the melting point of ammonia and a small linear correction t o standard tables applied. The indicated purities and melting point,s (obtained by extrapolation of the curve of temperature US. reciprocal of fraction melted to zero) were: B4H10, 99.0%, -121.0" (lit. --119.8)9; B5Hs,99.373, -47.2,' (lit. -46.8)'o; B5Hl1, 98.3%, -123.5: (lit. -123.!").11 It is probable that in all cases the principal contaminants were other boron hydrides, which would have a negligible effect upon the calorimetric results. The vapor pressure of the B6H1o a t 0.0" was 7.3 f 0.1 mm. in a volume such that 3% was vaporized and 7.2 f 0.1 mm. with 3073 vaporized (lit. T.z9; 6.8, 7.212). At 25.0' it was 28.3 f 0.2 mm. The apparatus and techniques were similar to those used for d i b ~ r a n e with , ~ modifications to decrease or eliminate (through use of break-seals) exposure of the gases t,o stopcock grease. Amounts of all reactants were determined by weighing in auxiliary bulbs. B4H10 was transferred from the external bulb to the reaction tube; in all other runs the boron hydride, because of its lower volatility, wae initially in the reaction tube and the stibine was in the external bulb; After the run, the hydrogen was transferred through -196 traps to a buret and measured; the traps then were warmed and the condensed gas measured. The lower two-thirds of the reaction tube then was flamed to the softening point of Pyrex to decompose solid hydrides, and the hydrogen was transferred through -196' traps and measured.
Results.-Results
of the runs are given in Table
I. (8) E. L. Sksu, Proc. Am. Acad. Arts Sei., 67, 551 (1932): J . Phys. Chem., 37, 609 (1933); see also J. M. Sturtevant, Calorimetry in "Techniques of Organic Chemistry," A. Weissberger, Editor, Vol. 1, "Physical Methods," Part 1, 3rd ed., Interscience Publishers, Inc.. New York. N. Y., 1959, p. 608. (9) A. Stock, "Hydrides of Boron and Silicon," Cornel1 University Press, Ithaca, N. Y., 1933. (lo! I. Shapiro and J. F. Ditter, J . Chem. Phys., 26, 798 (1967). (11) A. Burg and H. I . Schlesinger, J . Am. Chem. Soc., 66, 4009 (1933). (12) W. V. Kotlensky and R. Sehaeffer, ibid., 80,4517 (1958).
STUARTR. GUNNAND LEROYG. GREEN
2174 TABLE
1
HEATSOF DECOMPOSITION OF BORON HYDRIDES Boron hydride, mmoles
SbHi
0.409 .498 .573 .731 .705
1.583 1.988 1.831 2.194 1.423
Decompo- Condene nition able
%
%
&Hio 97.0 95.8 95.1 96.3 88.1
1.2 3.2 1.4 1.6 1.4
Hzin Solids
%
...
... 2.6 1.1 4.5
-AE,
kcal.
mole-1
10.9 11.0 12.1 11.8 13.2
11.8
0 ,423 .426 .475 .5m .678 .836
2.002 2.008 2.733
%He 89.9 3.6
...
...
96.0
0.2
1.881
...
1.725 2.047
93.2 93.1
1.0
...
1.2
. . . . . . 5.3 1.1
0.6
...
12.5 12.6 11.4 12.4 11.5 11.6 12.0
&Hi1
0.406 .520 .520 .709
2.321 2.032 2.017 1.989
07.4 93.7 96.9 85.3
0.8 3.3 0.9 13.0
0.9 1.5 1.6 3.1
18.7 19.3 19.2 20.4 19.4
0.096 .099 .lo7 .095
1.426 1.019 0.613 0.500
94 101 93 92
BBHIO 4 1 5 4
... 1 5
...
16.2 15.5 14.5 15.2
15.4
Observed heats have been corrected for the heat of compression, the fuse energy, and the stibine contribution (assuming -34.98 kcal. mole-' for a).Decomposition is calculated from the hydrogen found assuming quantitative decomposition of the stibine, and expressed as per cent. of theoretical. Condensable is assumed to be the unchanged boron hydride and is expressed as per cent. of sample. Hydrogen-in-solids is the non-condensable gas released by flaming the reaction tube, expressed as per cent. of theoretical hydrogen from the boron hydride sample. AE is calculated from the corrected heat and hydrogen produced, i t being assumed that no heat was contributed by the boron hydride not decomposed to the elements. There appears to be a slight trend toward larger - A E with smaller decomposition yield, but since the various boron hydrides were run over a similar range of mixture compositions and the results are treated in a relative manner, all runs are weighted equally in the average. The gas from BJ& runs 2 and 4 was lost and a yield of 95% assumed. We estimate the averages to be reliable within z k l kcal. mole-' for B4Hto,B5H, and B6Hll, and -+2 kcal. mole-' for BeHio. Correcting to constant pressure, the heats of decomposition are B4H10-9.4; B6H9, -9.9; B6Hll, -16.7; BBHIo, -13.0. Prosen, etaZ.,Zfound -13.0 for the heat of decomposition of B6Hgto amorphous boron, more negative than our result by 3.1 kcal.,
Vol. 65
or 0.62 kcal. (g. atom B)-l. Similarly, Prosen's value for BzHE,-6.73, is more negative than ours,4 -5.0, by 0.86 kcal. (g. atom B)-l. The agreement would seem to be more than coincidental; it is highly probable that the boron produced by our method of explosive decomposition is more finely divided and has a higher energy than the form produced by passing the gases through a heated tube. Hence we shall adjust our results to Prosen's, in essence taking his form of amorphous boron as the reference state, using 0.7 kcal. (g. atom B)-l for the heat of transition. Applying the usual estimated 0.4 kcal. (g. atom)-' value13 for the heat of transition from amorphous to crystalline boron, we calculate standard heats of formation of the gases: B4H10, f13.8; B5H11, $22.2; BaHIo, +19.6. Discussion Adler and Stewart14 have made measurements of the apparent equilibrium in the system B6H11112-B2HrB4Hlo a t various temperatures, obtaining a value of -7.56 kcal. mole-' for AH in the range 100-140" and about 1 cal. mole-' deg.-' for ACP of the reaction 2B6Hii(g)
+ 2H2(g) = 2B4Hidg) 3- M I s ( g )
While this is not a true equilibrium system, since components are being continually removed by competing reactions, it appears from the kinetic data that the equilibrium reaction is rapid compared with others. Our results give -9.3 for A H , which is in satisfactory agreement. Bauer'6 has predicted heats of formation of boron hydrides from a plot of AHfo/nus. 1 p/n where the hydride is described as B,H,+,; the known points for B, B2H6, B5H9 and B1oH14 were found to lie on a smooth curve. However, a non- obsolete value of +27 for AHro (BIOH4, g) was used; substituting the presently accepted value of +2.S3,18 the point lies far off the curve and it would appear that the relationship was largely fortuitous. Prosen" used the same input data to estimate heats of formation by a different method. He considered four bond types: B-H, B-B, B-H-B and BB-B, the first two being normal covalent bonds and the last two three-center bonds as formulated by Eberhardt, Crawford and Lipscomh, l8 and solved for the four unknown energies from the four known heats of formation. The procedure neglects resonance energies and changes in single-bond energies indicated by varying lengths and angles, but these may be expected to cancel to some degree in backcalculation of unknown heats of formation. We shall follow the appronch of Prosen and calculate thermochemical bond enrrgies a t 298"IC using 135.22 for aHfo(B,g)'" and 52.09 for AHtO (H, g).13 From the structural information of Lips-
+
(13) F. D. Rossini, el al., Circclar of the National Bureau of Standards, 500, 1952. (14) R. G. Adler and R. D Stewart, J . PhUs. Chem., 65, I72 (1961). (15) S. 1%.Bauer, J . Am. Chem SOC..80, 294 (1958). (16) W. H. Evans, E. J . Prosen and D. D. Wagman. "Thermodynamic and Transport Properties of Gases, Liquids, and Solida." American Society of Mechanical Engineers. McGraw-Hill Book Co., New York, N. Y., 1959, p. 226. (17) E. J. Prosen, American Chemical Society. Spriug. 1955: see ref. 15
(18) W. H. Eberhardt. B. Crawford, Jr., and IV. N. Lipacomb. J . Ckem. Phys., 22, 989 (1954).
THALLO US-THALLIC EXCHANGE AT VARIOUSACIDITIES IN PERCHLORATE 2 175
Dec., 1961
~ ~ m bwe, ascertain ~ ~ . ~ the ~ number of bonds of different types; these are listed in Table I1 together with standard heats of formation and heats of atomization, AH,. TABLE I1 ENERGIES AND BONDTYPES BH3 B2He BgHio BgHg BjHil BsHio BioHii
3Hf0
AHa
B-H
B-B
18f 1 7.53Z0.5 13 8 f 1 15 O f 0 . 4 22 2 % 1 19.63~2 2 83Z 1 5
273.49 575.48 1047.98 1129.91 1226.89 1312.62 2078.66
3 4 6 5 8 6 10
0 0 1 2 0 2 2
B-H- B-BB E
0 2 4 4 3 4 4
0 0 0 1 2 2 G
TABLE I11 B-B-B BONDENERGIES I
E(I3-B-B)
93 * 75 90.68 92.65
BsH,
AHa (calcd.) BaHii BaHio
BioH14
97.78
1233.04 1314.82 2054.46 x 1126.84 x 1308.68 2036.04 1128.81 1230.84 x 2047.86 x 1133.94 1241.10 1322.88
Experimental
1129.91 1226.89 1312.62 2 0 2 . 6 6
bonds and hence the B-B-B bonds may be expected to be less perturbed by resonance effects. B5Hs, BeHlo and B10H14 then have resonance stabilization energies of 3, 4 and 42 kcal., respectively. If for Bl0HI4one assumes a 10-04-8 bond conTo calculate the bond energies, we now have figuration instead of 10-24-6, the calculated value seven equations in four unknowns, an over-deter- of AHa in set 2 is 2058.72, deviating only half as mined set. The simplest approach is to calculate much from the experimental. The B-B distance to E(B-H), E(B-H-B) and E(B-B) from BHI, BzHs which the single B-B bonds are assigned in Lipscomb's formulation are only slightly shorter than and B4H10successively other B-B distances in the molecule; it may be E(B-H) = 1/3 LW.(BHJ 91.16 that the actual electron distribution is such as to E(B-H-B) 1/2 [AH,(B*Hs) - 4E(B-H)] = 105.42 make the bonding more uniform throughout the molecule. E(B-B) = AHJBIHlo) - 6E(B-H) - 4E(B-H-B) = The value of 79.3 for E(B-B) is in excellent 79.34 E(B-B-B) then can be calculated from any higher agreement with the 79.0 calculated from B2CL2'; boron hydride. These solutions are given in tge bond distances are also the same.19122 Acknowledgment.-We wish to thank Edward Table I11 where x denotes an input datum. It is perhaps most reasonable to accept the sec- J. Prosen of the National Bureau of Standards for ond set for general use, since BSHI1contains no B-B reading and commenting upon the manuscript. i=
(19) W. N. Lipscomb, J . Phys. Chem., 22,985 (1954). (20) F. L. Hirsrhfelcl, li. Erika. R. E. Dickerson, E. L. Tippert, Jr., and W. N. Lipscomb, ibzd., 28, 56 (1958).
(21) S. R. Gunn and L. G. Green, J . Phyd. Chum., 63,1787 (1959). (22) M. Atoji, P. J. Wheatley and W. N. Lipscomb, J . Chem. Phyr., 27, 196 (1957).
THE THA1,LOUS-THALLIC EXCHANGE AT VARIOUS ACIDITIES IN PERCHLORATE MEDIA' BY EDWIN ROIGAND RICHARD W. DODSON Radioisotope Applications Dioision, Puerto Riw Nuclear Center, Rio Pedras, Puerto Rico, and Chemistry Department, Brookhaven National Laboratory, Upton, Netu York Receiaed June 26, 1061
The Tl(1)-Tl(II1) electron exchange rate was measured a t 25" in 3 f NaC1O4-HClO4 media at various acid concentrations down to 0.1f. Since the hydrolysis equilibria of Tl(II1) have been measured under the same conditions, the result6 permit a more accurate comparison of the reactivities of T1+++and TlOH++than has been possible previously. In contrast to results a t higher ionic strengths, the rate is found to decrease as the acidity is decreased. The rate law can be written as R = Lo(Tl+)(Tl+++) kl(T1+)(TIOH++),with ko = 0.253 f 0.005f-1 h.-1 and kl = 0.089 i 0.012 f-1 h.-1. No evidence is found for an exchange of T1+ with Tl(OH),+; an upper limit of 0.l.f-l h.-1 is assigned to the specific rate of this reaction path. Measur'ements also were made a t 15' and 35". The activation energy for ko is 17.4 kcal./mole.
+-
Introduction In a number of cases it has been found that the rates of electron exchange reactions in aqueous perchlorate media decrease with increasing concentration of acid. This effect has been ascribed to reaction paths involving hydrolyzed ions, which were concluded to react more rapidly than unhydrolyzed ions. Such indication of hydroxide catalysis has been reported2-4 for the thallous(1) Researoh performed under the auspices of the U. S. .4tornio Energv Commission. (2) R. J. Prestwood and A. C. Wahl, J. A m . Chem. SOC.,71, 3137 (1949). (3) G. Hsrbottle and R. W. Dodson, ibid., 73,2442 (1951). (4) R. W, Dodson, ibid., 75, 1795 (1953).
thallic exchange, and the rate law has been formulated as R = ko(T1+)(Tl+++) kl(T1+)(TlOH++). From data obtained at ionic strength 6 it was concluded3.* that ko is negligible compared to k ~ . On the other hand, data obtained2 at ionic strength 3.68 indicated that both terms in the rate law are important. These studies were made before an accurate value for the hydrolysis constant of trivalent thallium was known, and the interpretations were subject to the uncertainty introduced by assuming that activity coefficient ratios are not seriously affected by the substitution of H + for Na+ at constant ionic strength. Subsequently, Biedermann's careful study6of the
+
