Vol. 67
SOTES
940
Another significant effect is the increase of CHzO concentration when nitrogen oxides are present. According to the mechanism proposed by EnikolopyanQ to explain his studies of methane oxidation, the maximum CH20 concentration should be independent of the surface. It has been shown, however, that PbO coatings alone reduce the concentration of CH20. Results similar to these have been found in the oxidation of methane over potassium tetraborate surfaces using oxides of nitrogen as homogeneous catalysts (see for example ref. 10). These workers observed that in quartz and steel vessels without the coating the reaction is not reproducible when oxides of nitrogen are present. Furthermore, they find that the amount of formaldehyde produced is increased by coating the surfaces with K2B407. The increase of CH,O when nitrogen oxides and PbO are present can be explained in a t least two ways. (1) The PbO alone allom CHzO to decompose or to be oxidized without chain branching. The nitrogen oxides then simply increase the rate of formation of CH20. ( 2 ) The PbO, as has been suggested previously,ll removes species such as KO2or H:O, which lead to chain branching. The nitrogen oxides then furnish a route for the oxidation of methane without the usual necessary destruction of CH,O. The latter possibility seems very probable. As pointed out above, the CH20yield begins to drop when the surface-to-volume ratio becomes less than about 8. A previous report3 indicates that marked diff ereiices in the rate of reaction appear when the surface-to-volume ratio is increased to 8 cm.-' in uncoated vessels. The fact that the same critical ratio appears in the present work indicates that the same radical species is involved. It is felt that studies of the methane oxidation using PbO coatings and nitrogen oxides offer a valuable tool for obtaining further insight into the role played by surfaces. Acknowledgment.-The author is grateful to Mr. G. C. McCollum and Mr. R. NT.Thomas for assistance in carrying out these experiments. (9) N. S.Enikolopyan, "Sepenth Symposium (International) on Coiiibuslion," Butterworths Scientific Publications, London, 1959, p. 157. (10) N. S. Enikolopyan, et a1 , Zh. Prrklad Khsm., 32, 913 (1959) (11) D. E Cheaney, et al , "Seventh Symposium (International) uu Cotubustion," Butterm ortha Scientific Publications, London, 1959, p. 183
LOKG RANGE SPIN-SPIN SPLITTlNGS IS 4-VINYL I D E S E CYCLOPEKTESE BY MELVINW.HANNA AND J. KENKETH HARRIAGTON
where aHand aH' are the hyperfine coupliiig constants of protons H and H', respectively, and Aa(T) is the a-electron singlet-triplet transition energy. For the methyl-substituted allenes studied by Snyder and Roberts eq. 1predicts J14= 2.9 c.p.s. 4-Vinylidenecyclopentene (I) has recently been isolated in these Lab~rat~ories,~ and its n.m.r. spectrum is of
I L)=c=cIE2 .-
I interest in checking a more general application of eq. 1. The proton resonance peaks of interest are a triplet due to the allylic protons a t 6.877 and a quintet due to the vinylic protons at 5.407. The spin-spin splitting if 4.33 f 0.08 c.P.s., substantially larger than the long range splitting in the methyl a1lenes.l This larger splitting is of interest because of the basic structural difference between allene I and the allenes studied by Snyder and Roberts. I n the latter case the methyl groups are freely rotating and the average value of a E ' = 7 5 X lo6 C.P.S. was used in eq. 1 to calculate 14"'. I n alleiie 1the methylene protons have a fixed spatial orientation with respect to the 2-p orbital on the adjacent carbon. This allows a test of the more ge11,era1 equation for the hyperfine coupling of an H-C-C fragment. I n this case
a"
(S155 cos2 4 ) X IO6 c.p.s.
(2) where # is the angle between the H-C-c plane and the &-orbital axis.24 Assuming an HCH bond angle of logo, aH' is 99.4 X 106 C.P.S. Using this value in eq. 1 gives AHH'(a) = +5.4 c.p.s., ingood agreement with the observed value of 4.6 c.p.s. The important thing is that a larger splitting is predicted for allene I than for met hylallene. Acknowledgment.-Acknowledgment is made to the donors of the Petroleum Research Fund administered by the American Chemical Society and to the National Institutes of Health for partial support of this research. (d) S. J. Crista1 and J. K. Harnnxton, t o be published. (4) C. Heller and H. &I. McConnell, J . Chem. P I y s . , 32, 15.33 ( L W O ) , D. Pooley and D. H. Whiffen, M o l . Phys., 4, 81 (1961).
THERMODYNAMICS OF SILVER BROMATE SOLUBILITY I N PROTIUM AYD DEUTERIUM OXIDES1 B Y RICHARD W.RAMETTE b N D EDWARD A. DRATZ
Department of Chemssti y, Unavetszty of Colorado, Boulder, Colorado
Depaitment of C h e m t s t i y , Caileton College, X o r f h f i e l d , llznnesota
Recezved September 24, 1962
Recent experimental work by Snyder and Roberts has shown that in a large number of allenes and acetylenes the spin-spin splittings bet\\-een protons separated are between 2 and 3 c.p.s.l These by four carbons (J14) experimental results can be nicely correlated by Karplus' theory of a-electron coupling of nuclear spins. I n this theory the a-electron contribution to the spinis given approximately by spin splitting, &H'(n),
=
Rererbed October
4
1962
Except for studies of weak acid dissociation, w r y little attention has been given to ionic equilibria in deuterium oxide. One recent paper deals with oxalatocomplexes of cadmium and copper.2 Almost as soon as deuterium oxide was isolated it was discovered3 that sodium and barium chlorides are less soluble in this solvent than in ordinary water, and since that time a number of solubilities have been determined (1) From the Senior Thesis submitted by E A Dratz, Carleton College,
1961. (1) E. I. Snyder and J. D. Roberts, J . Am. Chem. Soc., 84, 1682 (1962). (2) M. Karplus, J . Chem. Phvs., 33, 1846 (1960).
(2) D L Moldasters ( t al. J Phys. Chem , 66, 219 (1962) (3) H Tajlor, E. Caley, and H Eyring J A m . Chem S a c , 66, 43.34
(1917)
April, 1963 by Menzies and co-worker~,~ N ~ o n a n Chang ,~ and coworkers,'j and a few others.' However, these studies have been concerned with rather soluble salts for the purposes of straight coniparisoii of solubilities, observation of temperature effects, and the use of solvent mixtures. I n an intensive series of investigations Lange and co-workers8 contributed a body of knowledge on the integral heats of solution and dilution of a large number of salts in both maters, but otherwise the literature lacks thermodynamic data, and it remains to apply equilibrium theory and activity coefficient coiisiderations to electrolyte solubility in heavy water. Silver bromate was chosen for the present and initial study because it can be prepared in pure form without water of crystallization, is soluble enough to permit accurate determination of concentration of the saturated solutions, which nevertheless are sufficiently dilute to permit application of the Debye-Huckel theory, and in general seems to be an excellent representative of 1:I completely dissociated electrolytes.v Experimental Crystalline silver bromate was prepared by slow and simultaneous addition of solutions of silver nitrate and potassium bromate to mechanically-stirred distilled water a t room temperature, over a, five-hour period. Gravimetric analysis of the washed and dried product (as silver bromide) gave 99.8y0 purity, while iodometric titration with thiosulfate using ammonium molybdate aa catalyst gave 99.7% purity. Recrystallization caused no improvement in purity. It has been shown that the use of sodium bromate in the precipitation of silver bromate leads to an impure product because of solid solution formation,10 but that this is not a problem when potassium bromate is used.11 Deuterium oxide was obtained from the Liquid Carbonic Division of General Dynamics Corporation, with their specification of greater than 99.5% purity. All other chemicals n7ei-e of analytical reagent grade. Lithium perchlorate was used as the inert electrolyte for the purpose of ionic strength variation, and each solution also contained 1 X M perchloric acid t o repress any tendency for silver to form hydroxy species. A large excess of silver bromate was added, and the solutions were rotated in borosilicate glass bottles for two days in a thermostated air-bath a t 38", and were then transferred to a water-bath controlled a t 35.0' where they stood for three days with occasional manual shaking. The waterbath was equipped with an electrically-heated cover to keep air above the bottles slightly above 35O, thus preventing distillation of water into the tops. With the aid of an apparatus described (4) F. Miles, R. Shearman, and A. Menzies, Nature, 138, 121 (1936); R. Shearman and A. Menries, J . Am. Chem. SOC.,69, 185 (1937); F. Miles and A . Menzies, ibid., 59, 2392 (1937); R. Eddy and A. Menzies, J . Phys. Chem., 44, 207 (1940); R. Eddy, P. hlachemer, and A. Menzies, ibid.. 45, 908 (1941). ( 5 ) E. C. h'oonan, J . Am. Chem. Soc., TO, 2915 (1948). (6) T. Chang and T. Chu, Z. physik. Chem., A184, 411 (1939); T. Chang and Y. Hsieh, Sci. Rept. Natl. Tsing Hua Univ., AS, 252 (1948), J. Chinese Chem. Soc., 16, 10, 65 (1949); T. Chang and E. Tseng, Sci. Rec., 3, 101 (1950); T. Chann and K. Wang, H u a Hsueh Hsueh Pao, 22, 414 (1956); T. Chang and Y. C h a m , Sci. Sinica, 4, 555 (1965). (7) R. Kingerly and V. LaWIer, J . A m . Chem. SOC.,63, 3256 (1941); L. Brickwedde, J. Res. Natl. Bur. Std., 36, 377 (1946); H. E. Vermillion, B. Werbel, J. Saylor, and P. Gross, J . Am. Chem. Soc., 63, 1346 (1941); J. Curry and C. Hazelton, ibid., 60, 2771 (1938); F. Hein and G. Biihr, Z. phgsik. Chem., B38, 270 (1937). (8) E . Lange, W. Martin, and H. Sattler, 2. ges. Naturw., 1, 441 (1936); Chem. Zentr., I, 4882 (1936); E . Lange and W. Martin, Z. EZektroehem., 42, ti62 (1936), 2. physik. Chem., A178, 214 (1937); A180, 233 (1937); E. Larigo and €I. Sattler, ibid., A119, 427 (1937); W. Birnthaler and E . Lange, Z. Elektrochem., 43, ti43 (1937); 44, 679 (1938). (9) It has been proposed, by analogy to other salts, that silver bromate might be very slightly associated in solution (P.B. Davies and C. B. Monk, J. Chem. Soc., 71, 2718 (1951)). This idea has not been applied in the present model because it is uncertain; the effectwould be very small in any case, and if the evidence ever becomes definite corrections can then be applied t o our calculations. (10) J. Ricci and J. Aleshnick, J . Am. Chem. Soc., 66, 980 (1944). (11) J. Rioci and J. Offenbaoh. ibid.. 1 8 , 1897 (1981).
KOTES
94 1
elsewhere12 the saturated solutions were filtered through fine porous glass by pressure, and five-ml. samples were pipetted into titration flasks. The entire sampling apparatus, including pipet, was maintained a t 35". The dissolved bromate was determined by iodometric titration using sodium thiosulfate, after adding 25 ml. of water, 3 ml. of 6 M sulfuric acid, about 1.5 g. of potassium iodide (sufficient to redissolve silver iodide), and allowing to stand in the dark for ten minutes. The bottles and remaining solutions were then equilibrated in a similar manner at 25.0" and finally at 14.7", and samples were taken a t each temperature.
Results Table I shows the average values (from two closely agreeing results in each case) obtained for solubility (moles/liter). TABLE I SILVERBROMATE SOLUBILITIES HzO solutions LiClOa, M
sx
103
De0 solutions
sx
103
5.56 4.54 4.94 5.99 5.16 6.26 14.7" 6.43 5.30 5.42 6.55 6.76 8.09" 8.73 7.33 7.65 25.0" 9.09 7.89 9.35 8.02 .loo 9.52 9.57 .000 11.22 0.26 ,025 12.03 0.62 .050 12.48 35.0' 0.92 .075 12.84 1.12 ,100 13.09 a This value is identical with that reported by C. B. Monk, Trans. Faraday SOC.,47, 292 (1951), and is consistent with the range of other values. See ref. 10 and 11, and I. Tananacv, et. al., Zhur. Obschei Khim., 19, 1207 (1949). 0.000 .025 .050 ,075 ,100 .000 ,025 ,050 ,075
Discussion If the solubility process is represented by the reaction AgBr03(c)
+ nH20 (or mD20) Ag+(aq.)
+ BrOs-(aq.)
(1)
where the ions are solvated, then application of tlic solubility product principle gives
K = Q.fz (2) where K is the equilibrium product of the ion activities, Q is the product of the equilibrium molarities, and f is the mean activity coefticient. I n logarithmic form, with the introduction of the Debye-Huckel equation for a 1:1electrolyte (3) The parameter A varies oiily with the temperature and the dielectric constant. Values of the latter for protium oxide were taken from Harned and whih the equation proposed by &Ialmberg14was used 10 calculate the values for deuterium oxide. At the teinperatures of 14.7, 25.0, and 35.0' the corresponding (12) R. W. Ramette, E. A. Dratz, and P. W. Kelly, J . Phzls. Chem., 66 527 (1962).
(13) H. S. Harned and B. B. Owen, "The Physical Chemistry of Eleatrolytic Solutions," Reinhold Publ. Corp., Kew York, N. Y..31d Ed.. 1938, p. 161. (14) C. G. Malmberu, J . Res. Natl. B u r . Std., 60, 609 (1958).
NOTES
942
values of the quantity 2A are, for deuterium oxide; 1.014, 1.030, and 1.050; and for protium oxide; 1.002, 1.018, and 1.038. It is seen from equation 3 that a plot of the experimental quantity, pQ, us. the ionic strength function ~ l ' ~ / ( lBap1i2) should be linear with a slope equal to -2A and an intercept equal to the logarithmic thermodynamic value of the solubility product, pK. A difficulty arises in the choice of a proper value for the ion-size parameter Ba. This mas treated as an empirical constant for each solvent at each temperature, and with the help of an IBM 610 computer successive values for Ba were tried, each time fitting the line by the method of least squares, until the straight line had the predicted value for the slope. Table I1 summarizes the values of Ba used and the values of pK found. The uncertainties in the pK values, estimated for 90% confidence, are of the order of 0.006 log unit.
+
TABLE I1 THERMODYNAMIC VALUES FOR SILVER BROMATESOLUBILITY PRODUCT IN PROTIUM A N D DEUTERIUM OXIDES PK
Ba
14.7 25.0 35 0
4.577 4.265 3.994
1.71 1.50 1.57
PK
4.749 4.414 4 126
Ba
1.56 1.56 1.82
A more regular progression of Ba values niight be expected, but it should be pointed out that a t the low ionic strengths used the pK values are not particularly sensitive to the choice of the value for Ba. Thus, the extrapolated value of pK varies about 0.002 log unit for each 0.1 unit change in Ba. The standard free energy changes for the processes of equation 1 were calculated from the relationship AFo = - RT In K , while the standard enthalpy changes were determined by means of plots of pK ( K expressed in molarity units) us. 1/T, fitted by least squares. Standard entropy changes followed directly from the free energy and enthalpy changes. These thermodynamic quantities are listed for the temperature of 25.0' in Table 111. The uncertainties shown are estimated for 90% confidence. Probable (50% confidence) errors are only one-third as large. TABLE 111 THERMODYNAUIC FUNCTIONS FOR SILVER AT AFQ, kcal./mole
H2O De0
5.83 dc 0 02 6.02 i 0 01
BRO'vL4TE SOLUBILITY
25' AHO, koal./mole
11 7 f 0 . 3 12 5 z t 0 . 4
for the standard entropy a t 25' of the aqueous silver ion, and 38.5 f 1.0 for that of the aqueous bromate ion. The present results for the entropy of solution in protium oxide can be subtracted from the sum of these ion entropies to give 36.3 f 1.5 for the entropy of crystalline silver bromate, which agrees with a calculated value of 37.3 i 0.6 reported by Kireev.17 Acknowledgment.-This work was supported by a National Science Foundation grant for scientific research. We are grateful to Robert F. Broman for performing the computer calculations, and to John B. Spencer for the initial experimental work. (17) V. A . Kireev, Zhur. Obsckei Kkim., 16, 1199 (1946).
THERMODYNAMICS OF COPPER(I1) IODATE SOLKBILITY IN P R O T K M AND DEUTERIUM OXIDES' BY RICHARD W. RAUETTEAND ROBERT F. BROX~K Department of Chemistry, Carleton College, Northfield, Minnesota Received October 6 , 1062
D20
Ne0 t, oc.
1'01. 67
4x0, cal /de& mole
1 9 7 f 1 2 21 8 f l 3
From these results it is seen that the lower solubility of silver bromate in deuterium oxide (Le,, the more positive free energy change) is due to an enthalpy effect which is larger and thus overconies an entropy effect which would favor higher solubility in deuterium oxide. This would indicate that the hydration energy is larger in protium oxide, although the water structure is broken down to a larger extent in deuterium oxide. This is consistent with the theoretical results of Swain and Bade+ for alkali halides, but further discussion awaits a larger body of data. Kelley and King16adopt 17.54 f 0.15 cal./deg.mole (I5) C. G. Swain and R. F. W.Bader, Tetrahedron, 10, 182 (1960). (16) IC. K. Kelley and E. G. King, U. S. Bureau of Mines, Bull. 591. 1962.
The lack of information on ionic equilibria in deuterium oxide solution and a bibliography of previous solubility studies have been reported in an earlier paper2 dealing with the thermodynamics of silver bromate solubility. The present work differs from the latter in two respects : a 2 : 1 charge type electrolyte is used, and the crystalline form includes water of solvation. For the latter reason the isotope effects, as characterized by the differences in the thermodynamic functions, are not merely dependent upon the state of the ions in the infinitely dilute solution. h more thorough comparison could be made if the thermodynamics of the reaction
were known, but such information is not yet available. Experimental The copper(I1) iodate monohydrate was prepared3 by adding 0.6 mole of iodic acid in 200 ml. of water and slightly over 0.3 mole of copper sulfate in 300 ml. of water simultaneously and dropwise into 400 ml. of hot water with constant mechanical stirring. The product was washed thoroughly with a total of about 2 liters of hot water, using a repetitive process of partial sedimentation and decantation t o remove smaller particles. After drying, the =say according t o iodometric titration with thiosulfate gave 99.9% for the purity. A monodeutrate form of the salt was prepared in a similar manner but on a smaller scale: 125 ml. of 0.3 &I potassium iodate in D20and 125 ml. of 0.15 M copper sulfate in DzO were added dropwise to about 30 ml. of D2O a t 60' with constant stirring. The precipitate was digested at this temperature for an hour, was collected on a glass filter, and dried a t 100". Titrimetric assay with EDTA indicated a purity of 99.9%. The D2O was specified as greater than 99.5% pure (from General Dynamics Gorp.) and all chemicals were of reagent grade. The lithium perchlorate solutions, which also contained M perchloric acid to repress the formation of hydroxy-copper species, were equilibrated with solid copper iodate and sampled as previously described2 at 14.7, 25.0, and 35.0'. Analysis for dissolved copper was made by adding, to each 5-ml. sample of saturated solution, 2 ml. of 0.1 M acetic acid, 0.1 M sodium (1) From the Senior Thesis submitted by Robert F. Broinan, Carleton College, 1961. (2) R. W. Ramette and E. A. Dratz, J . Phvs. Chem., 6 7 , 940 (1963). (3) J . L. Sudmeier, Senior Thesis, Carleton College, 1959.
