680
Discussion The 1.esi11ts indicate that a t pressnreq of nitric oxide of 4 mm. 01: less, the nitrogen coiltailled 95% or more of Nl5xl4 and the nitrous oxide was l&olly x2110. 111 this narrOr$r range of conditioi1s1 the mechanLsnl of sere\l-icy. and Eoyesz may be said to be confirnled. Theie rvorkel:s hasre obserircd that with at1 illcrease in the pressure of nitric oxide, the btoichiometrp of the products is illincreasingly unsatisfactory. The preqent study points out why this is so. As indicated in Fig. 1, up to 4 mm. of nitric oxide, the rate of formation of X2I4is 110 more than the uncertainty in the results. But a t higher pressures of nitric oxide this is no lorlger It found Serewice and NOYes3 :Lnd crlllfirmed in the Preserlt instance, that at room tcniperatnre, ~ n d e rotherwi-ise constant conditions, tlie rate of production of i3itroge.n xvith an increase ill the nitric oxide pressure If the light abqorhi12g species is Only ammonia, and the reactions 1-6 are the only important steps, this trend cannot be explained. On the other hand, the rate of production of h~15~14 alone is seen to be constant with nitric oxide pressure within experimental error. QThile it is clear that in addition to step 2a, nitric oxide gir;es rise to llitrogellin a second Tvaysince this i i the only way W214 may be formedit is not obvious by what reaction this takes place. Both nitric oxide and ammonia absorb in the same general m\-e length region3 but most of the light is absorbed hy the ammonia since its extinctioll coefficierltis Mla173r times greater than that of nitric oxide. since the absorption s p e c t r y of ammonia is discrete down to a t least 2175 A., it is possible that some particular mercury line or lines may be absorbed bp nitric oxide even in the presence of ~ ~to lead ~ to ~ very ~ ammonia. H ~ )this~appears little del-omposition in View Of the meager yield of K,I4 in the photolysis of nitric oxide with light filtered bv ammonia. The trjnsferof ellcrgy from an excited ammonia mol~culeto a nitric oxide through a collision 3Ppears to be exuc~liitledby the fact that in its excited state, the ammonia molecule has a very short lifetime. On the mode Of formation Furth”r of nitrogen from nitric oxide in this system must await a more detailed analysis of the products of the photolysis a t high nitric ovide pressures. Acknowledgment.-The author wishes to thank Professor W. Albert NOYPS, Jr., for his advice and ei7couragemenl.
’”‘
( 3 ) For r r f v e n c e s and a summary see R. A. Noyes, Jr., and P. A. Leiehton, “The Photochemistry of Gases ” Reinhold Publ. Corp., Wew York, Y. T , 1311, p ~ 135 . 370.
A SPFC’I‘ROPHOTONETRIC STUDY OF THE HY1)ROLYSTS OF PLUTOSIUM(1V) BY S. K. RABIIIEAU AKD R. J. KLIXE Universzty 01 C c l t f o r i i n , Los Blamos Scientafic LaLoratoiy, Lo8 Alamos, New ilf extco Rereiiifd Soiember 19, 1959
The hJrdrolysis quotients Of Some Of the tetravalent actinides have hren determined both Tvith
electromotive force method&3 and with spectrophotometric procedure^.^-^ Also, from kinetic btudies of the reactioii betweeu I’u(V1) and PU(111) a value has been adduced for the hydrolysis quotient of Pu(IV) in DzO which is larger than that in HzO.’ Since with a spectrophotometric method the K K / K Dratios were found to be greater than unity for U(IV) arid Kp(1T’) in perchlorate solution,6this method mas employed in a similar study ?f the effect of sollrent 11~011the hydrolytic v i l h u m for pu(ITT). The temperature coefficients of the equilibria both in and in n20 have heen measured. Experimental The Cary Model 14 recording spectrophotometer was used together with a &uble-chambered 10 cm. spectrophotometric absorption cell described previously.* The Pu(IV) stocsk solutions, prepared as described below, were added from a weight buret to one leg of the mixing cell. To the other leg, a pipetted volume of filtered sodium perchlorate solution of known concentration was added. Thp sodium perchlorate solution was used to maintain the ionic strength a t a constant value of two. After the solutions had attained temperature equilibrium in a water thermostat, they were mixed quickly, and the cell was placed in a temperaturecontrolled mater thermostat in the cell compartment of the spectrophotometer. First readings of optical densities were made within 20 sec. of the time of mixing. In general, the diminution of the optical densities with time as a result of the disproportionation of Pu(1V) was small, and extrapolatioIls were made to time with little difficulty. Spectrophotometoric measurements were made a t the Pu(IV) peak of 4692 A. and also on the shoulder at 3?00A. In D20solutions, the Pu(1V) peak shifted to 4684 A . Better precision in the valucs computed for the Pu(1V) hydrolysis quotieat was achieved in the use of the measurements a t 3300 A. At this wave length, the observed molar absorptivities increased with decreasing acidity. Thus, the hydrolyzed form, p u o ~ + 3has , a molar absorptivity greater than that of PU7 4 a t this wave length. Stock solutions of Pu(1V) were prepared fresh daily by,the ,dissolution of purified Plutonium metal in the aPPrOPrlate weight of standardized 71yoperchloric acid, followed by dilution and by the oxidation of pu(111) to pu(1v) with a we‘ hed aliquot of a standard potassium dichromate solution. %nly enough oxidant was added t o convert 90% of the Pu(111) to Pu(IV) to avoid both the formation of higher oxidation states and the presence of excess dichromate. Although no slowness has been reported previously9 in the oxidation of Pu(II1) to Pu(1V) by dichromate, it has been found in the present study that in solutions of Pu(II1) in which the perchloric acid concentration was 0.1 hi’ or less, the yellow color of the dichromate disappefred gradually, and the maximum optical density a t 4692 A. was reached only after about a minute at 2.40. Accordingly, the stock solutions of Pu(1V) were made either 2.000 or 4.000 ,tf in perchloric acid, since under these conditions no slowness is observed and also the rate of disproportionation of Pu(1V) is negligibly small. The final acidity of the Pu(1V) solutions in the spectrophotometer cell was calculated from the weight of the combined solutions in the cell together with the measured solution density and the volume and acidity of the pu(IV) stock (1) This work was done under the aiispicer of the C. S. Atomic Energy Commission. (2) S. TV. Rabideau and J. F. Lemons, J. Am. Chem Soc., 73, 2895 (1951). (3) S. ’T Rabideau, ibid., 79, 3875 (1957). ( 4 ) K. A. Kraus and F. Kefson, abid., 72, 3901 (1950). ( 5 ) R. H.Betts, Can. J. Chem., 3 3 , 1775 (1955). (6) J. C. Suili\an and J. C. Hmdman, Tms JOTRNAL,63, 1332 (1959). (7) S.W.Rabideau and R. J. Kline, iLtd , 62, 617 (1958). (8) S. W. Rabideau, ?bid., 62,414 (1958). (9) R. E. Connich. “ T h e Actinide Elements,” Natl. Kuclear Energy Ser., AlcGraw-II111 Dooh Co., Inc.. Nea Yoih. X. P I Div. I V , I ~ - A IDN, [). ztx.
68 1
NOTES solutions. The acid consumed in the solution of the plutonium metal and in the oxidation of Pu(II1) to Pu(1V) bv rlichromatcx and the acid produced by thr hydrolysis of Pu(1V) were all taken into consideration in the computation of the acidity of the diluted Pu(1V) stock solution. The Pu(1V) stock solutions in DzO were prepared as for the Hi0 solutions with the exception that DzO was substituted for all the dilutions. The sodium perchlorate solutions were prqnred from the anhydrous salt and the mole fraction of DzO was about 0.95 in all those experiments in which DLO was used as the solvent. In the computation of the corrected optical densities of ttir Pu(IV) solutions, the contributions of Cr+3, P u + ~and , t hc acwi-salt blank wei e stibtracted. These corrections were determined separately tinder evperimental conditions of temperature and solution composition as nearly as possible simulating those used in the hydrolysis quotient deterniinetions.
Calculations Under appropriate conditions of metal ion conceiitration and acidity, the hydrolysis of Pu+4 caii be represented by the reaction Pu +4( aq )
+ H 2 0 = PuOII
+3(sq.)
+ H +(aq )
f 1)
The hydrolysis equilibrium quotient for equation 1 is K = [ P ~ o ~ I + ~ I [ H + I / [ P ~ + (~2I)
It can be shown easily that the relationship between the experimentally observed optical density, 011, and the molar absorptivities of the unhydrolyzed and hydrolyzed forms of Pu(IT'), €1 and t2,respectively, is OD = ( e i [ € I + ] E ~ K~O[PU(IV)]/( ) [H+] K ) 13) where [Pu(IY)] represents the total concentration of plutonium in the tetravalent state and the factor 10 corresponds to the cell length in centimeters. The unknowns in equation 3 are €1, €2 and K . Sets of equations of the type shown in equation 3 were obtained, one corresponding to each hydrogen ion concentration used. Instead of using a trial and error procedure in the evaluation of K , use was made of a computer program developed a t Los hlamos Scientific Laboratory for the least squares solution of non-linear equations1° using the IBJI-704 machine. Values of €1, c2 and K together with their standard deviations were obtained which minimized the sum of the squares of the difference, ( O D c a l c d - OD,,,tl). Although it is possible that other minima may exist in the parameter space, the requirement of positive values for el and c2 arid K , together with the availability of reasonable estimates of thece quantities, seriously limits the number of such minima. Moreover, the procedure used was verified by supplying significantly different initial estimates of the parameters, yet through the iterative machine calculation, identical fin:d values of the parameters were obtained. Results Medium Effe&-In the spectrophotometric del ermination of the Pu(IV) hydrolysis quotient, it is necessary to assume that the molar absorptivities of P u +and ~ P U O H + ~€1,and t2, respectively, are not functions of the medium. A series of measurepents waq made a t 2.4" a t a wave length of 2692 A. in perchlorate solutions of ionic strength
+
+
( I O ) R H L\Ioorr a n d J? K Zeigier, "The Solution of t h e General Leait Squares Problc 111 sit11 Special Reference t o H i g l i - S ] ~ ~Com~d r i l l t i rs," Iley,oit I A-2367, 0 c t i l ) r r I:, 19j9.
4.00 in an attempt t o examine the validity of this assumption. The acidity of the Pu(1V) solutions was varied between 2.000 and 4.000 M . The total Pu(1V) concentration was held constant and sodium perchlorate mas added to maintain the ionic strength a t a constant value of 4.00. Yalues of 59.86, 59.72, 59.67 and 59.30 Jf-l cm.-' were obtained for four different solutions in the above acid concentration range. It can be concluded that c1 is not sensitive to the substitution of Na+ for H+ under these conditions. Little can be said of the constancy of t2 from thew measurements because of the small fraction of this species present in the solutions studied. In the experimental measurements of the Pu(1V) hydrolysis quotient in solutions of ionic strength two, it can be considered that the medium changes from 1.90 M NaC104-0.10 M HC1O4 to 1.99 111 NaC104-0.01 J P HClO,, rather than from 2 111 HCIOl to essentially 2 LIP NaC1O4 since, in general, values of el, and K were not significantly different with or without the inclusion of data above 0.10 M HC104. Thus, large medium effects do not seem to be mesent in the sDectroDhotometric evaluation of K. Beer's Law Tests.-Conformity to Beer's law within 1% was observed for Pu(1V) solutions a t 2.4' in either perchloric acid solutions of 0.050 or 2.000 M a t constant ionic strengths of two a t 3300 and 4692 8. The Pu(1V) concentrat,ions were varied by as much as fourteen-fold. Pu(1V) Hydrolysis Quotients.-In the a,nalysisof the specbrophotometric data, it was found t~hatt,he resu1t.s obt'ained a t 3300 8.were especially useful in the evaluat,ion of the Pu(1V) hydrolysis quotient's. The precision of the values of K as measured by the standard deviation was better by nearly an order of magnitude from data obtained a t 3300 t'han a t 4692 8. This result occurs because bhe molar absorptivity of the hydrolyzed species is much larger a t the lower wave length ( € 2 = ca. 300 M-I cm.-' a t 3300 8.t ~ca. . 20 AI-' cm.-' a t 4692 8.).In Table I are given the results of measurements a t 25 and 15.4" in the solvent3 H20 and
nzo.
TABLE I PU( I\') HYDROLYSIS QUOTIENTS IN PERCHLORATE SOLUTIONS O F I O N I C STRENGTH Two Solvent HzO
Dz0
No. NO. UC. expts.b K(3300A.) co expts. K(4092-k.) uo 25.0 6 0.0226 0.0038 F 0.0234 0.014 6 ,0183 ,0004 F ,0438 ,012 4 ,0196 ,0011 5 ,0347 ,013 15.4 8 .Oil2 ,0013 8 .O220 ,015 6 ,0122 ,0018 ,0011 7 ,020 25.0 6 ,0101 ,008 6 ,0134 ,0013 4 ,015 ,006 6 .O103 .O10 15.4 6 ,0096 ,0048 6 ,0038 ,0005
Temp.,
Estimated standard deviation obtained by least squares methods. * Kumber of measurements made a t acidities between 0.01 and 2 M with a given Pu(1V) stock solution.
Although the data are presented for both the 3300 and 4692 A. mare lengths, the results obtained a t the latter wave length have such large uncertainties that they define K only poorly. With veights assigned to the 3300 A. data of Tahlc I in propor-
tion to 1/02, weighted mean values of 0.0185 f 0.0004 and 0.0115 f 0.0008 were calculated for the Pu(IT') hydrolysis quotients a t 25" in the H20 and D:O solvents, respectively. The uncertainties of the weighted means correspond to 2a = 2(1/ &J~)'/?, where tu1 is the weighting factor. Thus, a K E / K Dratio of 1.6 is obtained a t 25" from these spectrciphotometric measurements. This result is similar to observations6in the cases of Kp(IV) and U(IV) for which K H / K Dratios &-ere found to be 1.6 and 1.2, respectively. This result is not in agreement with the conclusion reached in the kinetic studies of the Pu(T'1)-Pu(II1) reaction.' In this latter n.ork. the results seemed to indicate a K H , / K Dratio less than unity. This discrepancy is unresolved. Iiraus and Nelson4 give spectrophotometric data ohtained at 4700 A. from which it is possible to calculate 2 value for K in perchlorate solutions of ionit: strength 0.5 a t 25" with the criterion that the sum of the squares of the difference, (tcalcdEexptl), is made a minimum. Whereas these authors found an average value of 0.0249 for K by a trial and error procedure, with the use of their data and the above criterion together with the computer program wed in this study, a value of 0.019 f 0.002 was obtained. The parameters €1 and ez whicih together with the above value of K produced this minimum least squares sum had values of 54.7 f 0.3 and 9.9 f 3 J1-l em.-', respectively. From the temperature dependence of the Pu(IT) hrdrolysis quotient as given in Table I, a value of 8.5 f 0.9" kcal./mole is computed for AH for reaction 1 in H20. This result is in agreement with the value of 7.3 f 0.5 kcal./mole obtained for AH from electromotive force measurements qf the acidity dependence of the Pu(I11)Pu(IT') c ~ u p l e . A ~ somewhat higher AH is found for the hydrolysis reaction in D20, namely, 19.4 & 1.3" kcal./niole. X wider temperature difference over which K mas evaluated would have been desirable. Attempts were made to evaluate K in H20and in D,O a t a temperature of 2.4"; holvever, it was not possible to obtain a sufficient change in thc values of the observed molar extinction coefficients to provide meaningful data. At temperatures greater than 25O, the disproportionation of Pu(1T') makeq the extrapolation to zero time more difficult. In the evaluation of the hydrolysis quotient of Pu(1T') from measurements of the formal potentials of the PLI(III)-Pu(IV) couple, a value of 0.054 was obi nined for K a t 2.3" in perchlorate solutions of ionic strength t w 0 . 3 It is not clear why such discordmt values of the hydrolysis quotient have been obtained hy the spectrophotometric and potentiometric methods. Acknowledgments.-The authors wish to express their appreciation to Drs. J. F. Lemons and C. E. Holley, Jr., for helpful discussions and to R. H. r\Ioore irid R . E. J70gel for their assistance in the computer programming and in the statistical treatment of' the data. (11) Th, uncertainties g l \ e n f o r the ralues of A H were computed by t h e rne.liod of propagation of errors u1th t h e use of t h e exprewon, 1 I(1/T1 - 1 / T 1 ) 2 [ U Z K I / K I z U21~2/K22])1/2. zkUAh
+{
+
THERMOCHEMISTRY OF DIMETHOXYBORASE BY W.J. COOPER . ~ S DJ. F. MAPI Callerg Chemrcal Company, Callery, Pennsylaania Received ?Jotember 90, 1059
Dimethoxyborane was first prepared by Burg and Schlesinger' from the reaction of diborane with methyl alcohol. They found the compound to be unstable above Oo, disproportionating according to the equation 6HB(OCH3)?
BzHs
+ 4B(OCH3)3
(1)
Nore recently Cchida,? have reported the kinetics of the vapor phase disproportionation in which the reaction was found to be second order with respect to dimethoxyborane and surface dependent. A determination of the heat of formation of dimethoxyhorane would not only enable one to calculate the heat of the disproportionation reaction but also afford a means of estimating the strength of the B-H bond in a somewhat different manner than from the heat of formation of BH3.3 Preliminary investigation of the heat of combustion of dimethoxyborane showed that the combustion was insufficiently complete for an accurate determination. Since the hydrolysis was reported to proceed readily,' the heat of hydrolysis was studied and found to be a satisfactory measurement for heat of formation determination. Experimental Sample Preparation and Purity.-The dimethoxyborane was prepared by reaction of diborane and methyl alcohol and was purified by distillation in a lowtemperature Podbielniak column. The purity, as determined by hydrolyzable hydrogen, was found to be greater than 98y0. As an additional check on the purity, the 0" vapor pressure (275 mm.) was measured prior to filling the hydrolysis sample bulbs. The sample bulbs of thin-walled Pyrex were filled by vacuum condensation and sealed while frozen a t liquid nitrogen temperature. To minimize disproportionation, the sample bulbs were stored in liquid nitrogen until they vere weighed prior t o introduction into the calorimeter. Calorimeter.-The calorimeter was of a design similar to that described by Van Brtsdalen and B n d e r ~ o n . ~The calorimeter vessel was a 250-ml. silvered dewar flask. The calorimeter heater, approximately 25 ohms, was of No. 32 constantan wound non-inductively on the stirrer well. This heater was used both for electrical calibration and for adjustment of the calorimeter temperature. The platinum resistance thermometer had a resistance a t 0' of 25.506 ohms and was calibrated by the National Bureau of Standards. The calorimeter was immersed in a water-bath maintained a t 25 i 0.05'. An outlet tube connected to a gas measuring buret provided for collection and measurement of the evolved hydrogen. Calibration and Hydrolysis Reaction .--A series of electrical calibrations of the calorimeter system was made, as sho%nin Table I. In addition, a calibration was run in conjunction with each hydrolysis experiment. S o significant differences in energy equivalents nerr observed hetwecn initial and final states. The amount of dimethoxyborane which reacted was determined from the weight of sample, with completeness of (1) A. B. BuIg and H. I. Schlesinger, J . A m . Ciiem. Soc., 66, 4023 (1933). (2) H. S. Uchida, H.B. Xreider, -1. 1Iurcliisiin and $1. F. Xlasi, THISJOURXAL, 63, 1414 (1959) (3) R. E. hlcCoy and S . H. Bauer, J . Am. Chtrn. S O C . , 78, 2001 (1950). (4) E. R. Van Art?dalPn and K. P. Anrlrrson.
