J. Pbys. Cbem. 1981, 85, 3905-391 1
3905
Chemical Kinetics of the Gas-Phase Reaction between Uranium Hexafluoride and Hydrogen A. L. Myersont and J. J. Chludrlnskl, Jr." Exxon Research and Engineering Company, Linden, New Jersey 07036 (Received: June 9, 1980; In Final Form: August 3, 1981)
A study was made of the chemical kinetics of the homogeneous gas-phase reaction between uranium hexafluoride and hydrogen by measuring the rate of disappearance of UF6. It has been concluded that the rate-limiting step for which the kinetics have been measured is UF6 + Hz = UF5+ HF + H (2). The reaction has been studied in a steady-state flow system over a temperature range of -625-825 K. Various surface-to-volumeratios were employed to aid in distinguishing gas-phase reactions from surface reactions. The steady-state concentration of the UFGafter reaction with H2was monitored in a special multipass infrared spectrophotometerat the 626-cm-' absorption band of UFs. The principal problems were corrosion, plugging, surface intrusion, and the deleterious effects of minute traces of water; these problems have greatly slowed progress in this field. Several series of measurements involving different initial species concentrations and residence times, with each series at constant temperature, show that the rate is first order in UF,. Our measurements yield a dependable Arrhenius curve in terms of a second-order expression for k , the bimolecular specific reaction rate constant for the disappearance of UF,. It is believed that this overall rate of the disappearance of UF6 is somewhat less than twice that of the critical reaction step 2, indicated above, so that the specific reaction rate constant k2 is -8.7 X 1014exp(-34550 kcal/RT) cm3 mol-' s-l. Conclusions have been reached concerning the relative importance of the various elementary reaction steps involved in the chemical mechanism.
Introduction The literature on the chemical kinetics of homogeneous gas-phase UF, reactions is, in general, sparse indeed, let alone that on a specific reaction such as UF6 + H2 which is the subject of the present study UF, + [(6 - n)/2]Hz = UF, + (6 - n)HF (1) where n is usually considered to be 4 and/or 5. The only indication of UF6 H2 kinetics apparent in the open literature is that which is incidental to an investigation by Dawson, Ingram, and Bircumshawl (DIB) of catalytic aspects of the reaction. The mixtures of UF6 and H2 were reacted in sealed bulbs, and the kinetic results were neither considered conclusive by the authors nor included in the paper. These incidental and preliminary results of DIB were later used by Tumanov and Galkin (TG) to calculate the only experimental rates present in their relatively recent (1972) review2 of the UF6-H2 reaction. The reaction mechanism proposed by TG is strictly speculative and includes the following gas-phase reactions: UF, H2 = UF,+ H F H (2)
+
+
+
UF5 + H = UF4 + HF UF6
(3)
+ H = UFB + HF
(4)
H+H=H2
(5)
UFI (g) = UF4 (SI (6) The findings of this work provide reliable rate data and lead to what we believe is a plausible mechanism for reaction 1. With respect to the above scheme (TG), only reactions 2 and 4 appear to participate. The critical elementary reaction step that emerges in the present case is the abstraction of an F atom from UF, by H,: UF6 + Hz = UF, H F H (2)
+
+
which is also the rate-limiting step, as we shall see later. The elimination of the brief participation of an FH2 in+ Mote
Marine Laboratory, Sarasota, FL 33577. 0022-3654/81/2085-3905$01.25/0
termediate complex is also discussed. The experimental method chosen for the present study was a gaseous flow system which included a Monel reactor in a nitrogen furnace. The measurements consisted of determining the decrease in UF, concentration by monitoring {UF,}at steady state with the strong IR absorption O f UF6 at 626 Cm-'. By finally circumventing problems associated with wall reactions, the large effect of minute traces of water, condensation phenomena, and plugging, we have obtained dependable values of the rate of loss of UF6 in the gasphase reaction 2 over a temperature range of -625-825 K. Product analysis of the solids formed has been carried out. Because of the well-known instabilities of UF5, it has not been possible to identify UF5 in the final product, but it is shown that kinetic and other arguments point to reaction 2.
Experimental Section As was mentioned in the previous section, an all- (noncorrosive) metal flow system was developed. The apparatus, a schematic diagram of which is shown in Figure 1, was fluorinated with 2 torr of Fz which was left in the system for 20 min to complete the conditioning process. The entire system was then pumped down to 0.1 hm, through such chemical traps as granular CaO and MgO. The concentration or partial pressure of UF6 was monitored, before and after reaction, in a custom-made Wilks 1.0-m (or less, by external adjustment) path-length monochromator. The internal parts of the aluminum cell which were to be exposed to UF6 and F2were plated with nickel while the aluminum mirror optics (for path reflectance) were successively coated with nickel and gold. The lens and window optics were made of AgC1. All O-rings, gaskets, and seals were Viton. The resistance to corrosion of this system was quite satisfactory. Only the mirrors underwent a gradual loss of reflectance whereby perhaps (1) J. K. Dawson, D. W. Ingram, and L. L. Bircumshaw, J. Chem. Soc., 14,21 (1950). (2) Yu. N. Tumanov and N. P. Galkin, At. Energ., 32, 1, 21 (1972).
0 1981 American Chemical Society
3906
The Journal of Physical Chemistry, Vol. 85, No. 25, 198 1
Myerson and Chludzinski Beginning of flow of 1.O I m i n - l of Ar (10)(zero-attenuation)
P R
-
PJlnp ng fac I I, Rotameter
e - N a n e l or
I I
ia1w
Idlapllragm or h e l l o v i i ) T = cold trap G = Emidongauge \ c - VacuLm DI"(JP
KD
-P
-K
n3 Dryer
A,
I
MOkCCHROl!IATOR
G
-
F6 Ar
FLO!!
Flgure 1. Apparatus for gas-phase UF,-H,
chemical kinetics.
30-40% loss of reflectance took place over a period of several months of heavy use. The principal reactant carrier was a mixture of 1000-2000 ppm of pure UF6 (Allied Chemical Co.) in ultrahigh-purity argon. The mixture was prepared accurately ( f l % ) in situ by using an MKS Baratron precision pressure gauge to measure the partial pressure of UF6 and stored a t up to 4 mPa in a stainless-steel cylinder. The partial pressure of UF6, after its reaction with Hz, was determined by measuring the IR attenuation due to its v3 band at 626 cm-l. Calibration of (UF,) was achieved strictly through the use of experimentally determined curves. The latter were obtained by measuring the attenuation of a specific well-defined initial beam intensity, Io (in arbitrary recorder units), in the spectrometer cell. This was usually for a path length of 0.5 m, which provided sufficient attenuation for our purposes, and for a given slit width and beam intensity. The concentration, (UF,], corresponding to a given signal attenuation, Il/Io, was obtained by precise dilution with Ar of the original accurately prepared mixture of UF6 in h. while plots of log lo/lvs. (UF,] provided smooth curves with little scatter, departures from the linearity required by Beer's law were present. This was not surprising in view of the use of a multipass cell in which the special optical surfaces were never completely clear of the U compounds. However, this had no effect on the data since, as indicated, the experimental calibration curve was derived under conditions identical with the runs and updated when deemed necessary. The flow rate of the UF6-Ar mixture and that of the ultrahigh-purity Hz in Ar mixture were monitored by Pyrex and Teflon rotameters, all experimentally calibrated. It was found early in the research that very special care was necessary in drying the Hz in order to prevent the UF6-H20 reaction from becoming predominant. To that end, all Hz mixtures, just before entering the preheater, were passed through a drying trap containing 6-16 mesh SX-144-2 silica gel, immersed in a dry-ice bath. The flowing mixture of Hz in Ar was analyzed for H 2 0 to levels of less than 5 ppm by a Du Pont 510 moisture analyzer. The sieve material was regenerated periodically by a reverse flow of ultradry helium at room temperature. Because of the danger of ignition of Hz in air, a constant-temperature, propellerstirred, Nzbath was used. The two reactants first flowed separately through 36 in. of Monel-tubing preheaters before going through a mixing capillary and finally into the Monel reactor itself, which was 25 cm (ca.) in length. Three inner diameters were used for the reactors--1.09, 2.21, and 2.67 cm-which provided residence times, T,,,, of from 0.5 to 12 s. The reacted gas mixture was cooled rapidly after exiting the heated reactor in order to freeze it chemically. The gas then passed through two Teflon filters to remove all particulate matter before entering the IR cell, in order to minimize coating of the internal optics
"1
T = 748K
70 v)
z
0 II! 2
60-
E
50-
0
w 0 a:
40-
> k2 since is close to 4 kcal mol-1 and Eact-Pis certainly no less than 30 kcal mol-' (as measured here and as would be expected). It can be concluded, since the specific rate constant of reaction 4 is so much faster (by roughly 8 orders of magnitude) than that of reaction 2, that the total rate of dis-
+
+
(5) JANAF Thermochemical Tables, 2nd ed., Nutl. Stand. Ref. Data Ser. (U.S., Nutl. Bur. Stand.), 37 (1971).
3910
Myerson and Chludzinski
The Journal of Physical Chemistry, Vol. 85, No. 25, 1981
7.0
\
REACTOR i.d.=2.67cm (largesl i , d . ) ; I=23.2cm DATA material: monel
{
A
\A \A \A
\
1 A\
4*0
t
3.0 0.5
While the vapor pressure of the UF, is sufficiently high to keep it in the gas phase, that of UF4 is so very low that it will condense out as soon as formed, aiding the progress of the reaction to the right. A major question then becomes “How fast is reaction 7?”, and our answer can be only a rough estimate which indicates that its effect is perhaps 1 order of magnitude less than that of reaction 2. If this is even close to correct, very little of the final UF4 found can be formed during the measured reaction as a result of reaction 7. However, there is ample opportunity for UF4 formation after the measurement. First, there is the long exposure of the deposits to higher temperatures after the measured reaction where, e.g., the deposits may remain 1 h or more while cooling before removal. Secondly, UF5 also hydrolyzes readily (e.g., in the air) to U(1V) and U(VI), and only moderate precautions were taken during transfer for analysis. Thus, there is little likelihood for retention of the UF5 compound. We now consider reactions which may reduce UFs to UF4 during the reaction residence time. The principal possibilities are
+ Hz = UF41 + H F + H UF5 + H = UF41 + HF
(8) (3) Reaction 8 can be eliminated immediately, since the rate of reduction of UF, by this step is slower than that of UF6 by reaction 2 by roughly 8 orders of magnitude. This conclusion follows from recent thermochemical data4 which indicate that the bond energy of UF5-F is 68.3 kcal mol-’ (reaction 9, Table 11) as compared to 101.0 kcal mol-’ for UF4-F (reaction 10, Table 11). Hence, since even reaction 2 is endothermic, the activation energy of reaction 8 is some 30 kcal mol-’ greater. Similarly, the rate of reduction of UF5 due to reaction 3 is slower than the rate of reduction of UF6 due to reaction 4 (to which reaction 3 is analogous), and in this case the k is estimated to be slower by l/z order of magnitude (the ratio is approximately t h a t of exp(-6000/RT)/ exp(-4000/RT) where 6000 and 4000 are Eact-3 and respectively (as estimated by an extension of the Hirschfelder rules). Thus, the loss of UF5 via H will be secondary to the loss of UF6 via H. Paraphrasing, relatively little of the UF5 formed during the true kinetic residence time in the reactor, will have the opportunity of being reduced to UF4. This is because most of the H atoms are involved with UF6. As previously discussed, this implies the conversion of UF5 to UF4 via other routes and subsequent to the real or kinetic residence time. The direct recombination of atomic hydrogen H + H + M = Hz + M (5) UF,
1 .o
1.5
2 .o
I / T (K)X103
Figure 6. Final Arrhenius plot for the gas-phase UF,-H?.eaction. The specific reaction rate constant for k 2 = 8.7 X 10 exp(-34550 kcallRT) cm3 mol-‘ s-‘,which is half the k calculated from the best Arrhenius line as discussed in the text.
appearance of UF6, as measured, is probably essentially twice that of the rate of reaction 2 alone; i.e., one additional molecule of UF6 is decomposed immediately by reaction 4 for each UF6 molecule that reacts with an Hz molecule. The rate of loss of H atoms by recombination also has been estimated and found to be far too slow (by at least 2 or 3 orders of magnitude) to provide any competition for reaction 4. This is in accord with the equation for d{UF,]/dt readily obtained by writing out the rate equations for reactions 2 and 4 and assuming steady state for {H),which yields first order for (Hz}and the factor of 2 mentioned. Thus, the actual rate constant of the elementary reaction 2 would be kz = 8.7 X 1014exp(-34550 kcal/RT) cm3 mol-l s-l However, this is made less straightforward by the fact that analyses of the product found in the reactor indicate that the final product is UF4. The analyses were carried out in part by LeDoux Co. and in part by ER&E’s Analytical and Information Division (AID). Final resolution, correlations, and interpretation of the data were carried out by Gaines? who concluded that the great weight of the data indicated that practically all of the final, stable product found was UF4. Now this does not a t all mean that the immediate product of the chemical kinetics which we are measuring is UFI. UF5, which may well be the final and immediate product of the primary steps with which we are concerned, is unstable. First of all, it disproportionates thermally7~sto UF6 and UF4. 2UF5 = UF4(s)4 + UF6 (7) (6) P. Gaines, ‘‘Elemental Analysis of Materials Formed from the Reaction of UF, with Hzat Elevated Temperatures”, Exxon Research and Engineering Co., Ref. No. 77AN443, April 5, 1977. (7) E. Jacob, Z. Anorg. AElg. Chem., 400, 45-50 (1973).
does not affect the scheme, as previously mentioned, since it is several orders of magnitude too slow. As previously indicated, our measured value of the activation energy of kz is 34.55 kcal mol-l, not too far off from the endothermicity (-36.4 kcal mol-l) of the rate-limiting abstraction reaction: UF6 + H2 = UF5 + H F + H (2) On the basis of endothermicity, its Eactshould be >36.4 kcal mol-l, and certainly at least several kcal more than that, e.g., 40. Thus, the minimum estimated activation energy of the reaction of complete decomposition, reaction 2, is higher than the measured value by somewhat more than 5 kcal, which could easily be accounted for by an (8) A. S. Wolf, W. E. Hobbs, and K. E. Rapp, Znorg. Chem., 4, 755 (1965).
J. Phys. Chern. 1981, 85, 3911-3916
391 1
found to be first order in u F 6 and Hz. The specific reaction rate constant measured for the disappearance of UF6 due only to the principal rate-limiting elementary step, reaction 2, is given by the expression k2 = 8.7 X exp(-34500)/RT cm3 mol-I s-l with an error limit well within 2~50%.This rate is slightly more than one-half that of the net rate of disappearance of UF, because one of the products of reaction 2, the H atom, reads quickly with UF6 to produce another molecule of UF5. It is believed that UF5 is the final product of the chemical kinetics studied and that UF4 is formed subsequently.
-25-5070 error in the Arrhenius slope and the greater uncertainty of the estimation of EaCt.The possibility of FH2 being a short-lived but final product (rather than H F H) has been considered briefly but the reaction is then, as follows from the basic thermodynamic calculation, so endothermic (-70 kcal) that the activation energy becomes impossibly high. In summary, the rate of reaction 2 is well within the limits allowed by experimental and estimated uncertainties, so that it is concluded that kz = 8.7 x ioi4 exp(-34550/RT) cm3 mol-1 s-l.
+
Conclusions The chemical kinetics of the homogeneous gas-phase reaction between UF, and H2 have been measured and
Acknowledgment. We are grateful to Dr. John Horsley for his helpful advice and discussions.
A Study of Rearrangement Energies of Redox Species Karl W. Frese, Jr. SRI International, Menlo Park, California 94025 (Received: June 23, 1981; In Flnal Form: September 3, 1981)
A large body of rate constant data from homogeneous electron-transfer reaction kinetics studies was analyzed by use of the complete Marcus theory. The goal of the analysis was to obtain reorganization energies to be used in kinetic analyses of reaction rates at semiconductor electrodes. In general, quite consistent results were found further supporting the basic tenets of the Marcus theory.
Introduction The theory of semiconductor electrode kinetics that is most used by workers in semiconductor electrochemistry today is due to Gerischer.l The general concepts of this theory have been widely used to qualitatively explain relative reaction rates on semiconductor electrodes.2 Much less quantitative testing, however, has been done. The main reason for this is that certain parameters in the theory, most notably the reorganization energy, A, have not been known very accurately. The reorganization energy is the change in free energy for the process in which the environment of an electron (ligands and solvent) is changed from its initial state to its final state while the electron remains associated with its initial core. These environmental changes correspond to libration and vibration of solvent molecules and charged ligands as well as bond distortions within a reacting molecule itself. Some very recent3v4and some older techniques5s6have been developed to determine X by using semiconductor electrodes. Hbwever, these techniques have been applied to a limited number of redox species. There is an extensive literature on electron-transfer rates in homogeneous solution, The rate have been measured for both the exchange, e.g,, F ~ z + / F ~and ~ +so-called , reactions, e,g., ~ ~ 2 + ce4+.These ~omogeneousrate constants are usually in-
terpreted with the rate theory given by Marcus.’ This theory contains the same reorganization energy term X as does the Gerischer model. In principle, both theories distinguish between A,, for oxidized form and A r d for the reduced form of a redox pair. It is the purpose of this paper to obtain values of the reorganization energy for various redox species by using the Marcus theory and the experimental rate constants for homogeneous reactions. These X values can be used in the Gerischer theory to better understand the quantitative aspects of similar reactions occurring a t semiconductor electrode surfaces. The outline of the paper is as follows. We will first give the pertinent equations for the models of Gerischer and Marcus. Then, reorganization energies obtained by using the Marcus equations and the homogeneous solution rate constants will be tabulated. When possible these results will be compared to those previously obtained.
/
Theoretical Models Gerischer Model. The role of reorganization energies in electron-transfer reactions on semiconductors is illustrated by Gerischer’s equation which is outlined below. The rate of electron or hole transfer between the semiconductor bands and the redox species in solution is mainly controlled bv the relative enerev levels involved. For example, in thk case of an n-type semiconductor, where reaction between photogenerated to the energy level a t the top of the valence band, EVB, and a reducing agent characterized by the energy level Bredis ‘Onsidered, the anodic current density is Of the forms’Q VI
(1)H. Gerischer in “Physical Chemistry, An Advanced Treatise”, Vol. IX-A, Erymg, Anderson, and Jost, Ed., Academic Press, New York. Also
see references cited therein. (2) S. R. Morrison, “Electrochemistry at Semiconductor and Oxidized Metal Electrodes”, Plenum Press, New York, 1980. (3) K. W. Frese. Jr.. M. J. Madou. and S. R. Morrison. J. Electrochem. SOC.; 128,1527 (198ij. Ser., (4)No. S. R. 146 Morrison, (1981). M. J. Madou, and K. W. Frese, Jr., ACS Symp. ( 5 ) R. Memming and F. Mollers, Ber. Bunseges. Phys. Chem., 76,475 (1972). ’ (6)R. A. L. Vanden Berghe, F. Cardon, and W. P. Gomes, Surface Sci., 39, 843 (1973).
(TEd
jVB+ = A[h+][Red] -)’2e-U%-Elll * / 4 b k T (7) R. A. Marcus, J. Chem. Phys., 43, 679 (1965).
0022-3654/81/2085-3911$01.25/00 1981 American Chemical Society
(1)
