LITHIUMIN TRIETHYLSILANE-ETHYLAMINE REACTION
May 5 , 1964
in the above calculations may reflect the differences of the spin states of cobalt(I1) and cobalt(II1) as well as the free energy of the elementary electron-transfer step.?’ However,. no such variation of oxidation potential seems sufficient to account for the discrepancy observed between values of kz2 calculated for the V 2 + and the Ru(”3)6’+ reactions. Marcus’ theory is not invalidated by failing to account for the rates of reaction in a few particular cases. It does offer promise’4,25of correlating rates of electron transfer in a large number of “normal” cases, suggesting t h a t some special effects are operating in the reactions of C0(”3)6~+ with reducing agents; one will need to go beyond the theory t o discover what these special effects are. I n the reaction between Co(”3)e3+ and Ru(NH3),j2+ the over-all free energy of activation should reflect some compromise between the standard oxidation potential ( 2 2 ) This point has been discussed with regard t o t h e standard oxidation potential of Co(NHz)aOH22+-Co(NH3)sOH9 +.*a ( 2 3 ) A Haim and H. Taube, J . A m . Chrm. Soc.. 86, 499 (1963). (24) G. Dulz and K.Sutin, I n o r g . Chem., 1 , 917 (19F3) ( 2 5 ) R. J Campion, N . Purdie, and PI’. Sutin, ibid., in press.
[COSTRIBUTIOSFROM
THE
1691
of C O ( N H ~ ) ~ * + - C O ( N Hand ~ ) ~the ~ +standard oxidation potential (- 1.8 v.I3) of C O ( H Z ~ ) ~ ~ + - C O since (H~~)~~+ the ultimate cobalt(I1) product is Co’+(aq). Evidently the immediate cobalt(I1) species, once formed, decays rapidly into the equilibrium cobalt(I1) species. The other ruthenium(I1)-cobalt(II1) reactions reported in this paper seem to be more favorable thermodynamically. 26 Acknowledgments.-We wish to express our appreciation to Messrs. L. Bennett, R. Butler, D. Huchital, C. Andrade, R . Jordan, R. Linck, and J . Stritar, and Dr. E. S. Gould, who supplied a number of the complex compounds used in this study. The research described was supported by the National Science Foundation under Grant 6-20954. Funds for the purchase of the Cary Model 14 spectrophotometer were supplied by Grant G-22611 from the National Science Foundation. (26) For example, the standard oxidation potential of
C O ( N H ~ ) ~ O= H~ C *O~( S H ? ’ ~ ~ O+He~- ~ +
has been estimated t o be -0.33 v.27 (27) R. G . Yalman, Znorg. Chem., 1, 16 (1962).
GEORGEHERBERT JOKES LABORATORY, THEUSIVERSITY OF CHICAGO, CHICAGO, ILLINOIS]
The Role of Dissolved Lithium in the Reaction between Triethylsilane and Ethylamine BY ARLENVISTE AND HENRYTAUBE RECEIVEDDECEMBER 18, 1963 The reaction of triethylsilane with ethylamine is not significantly catalyzed by dissolved lithium itself. Instead the reaction appears t o be catalyzed by small amounts of a strong base such as lithium ethylamide, presumably present as an impurity in the lithium-ethylamine solution.
Introduction In 1934 Kraus and Nelson’ reported that triethylsilane reacts with ethylamine in the following way. ( CaHj)&H
Our main interest was in the possibility of genuine catalysis by lithium; if this were to take place, the reaction would offer a convenient means of studying the kinetics of a reaction involving a metal in solution.
Li + C 2 H b S H 2+ ( C z H s ) a S i h ” C ~ H+ s HZ
Experimental The stoichiometry with respect to Et3SiH and Ht Reagents.-Anhydrous ethylamine (Eastman No. 506) was was established, and the product EtaSiNHEt was isodried with metallic lithium and then distilled on the vacuum line lated and identified. By adding the lithium to a soluinto a storage vessel where it was kept over lithium wire until use. After each use, the ethylamine was distilled back t o this tion of Et&H in E t N H 2 , they found t h a t the reaction storage vessel. For most experiments, triethylsilane (Peninsular begins as soon as the blue color due to lithium develops, ChemResearch N o . 436) was sealed off in fragile Pyrex ampoules and (qualitatively) the rate of reaction increases as the under vacuum. Triethylfluorosilane (Peninsular ChemResearch Li concentration is increased. It was shown t h a t an S o . 2747) was redistilled. When used, it was outgassed on the initial mole ratio of Li to Et3SiH of either 1 or 0.1 vacuum line and condensed into the reaction flask. Small pieces of lithium metal (Lithium Corporation of America) were sealed results in catalysis. Accordingly, they concluded off in fragile Pyrex ampoules under vacuum. that “the reaction is a homogeneous one and the Apparatus and Procedure.-About 95% of the 420-ml. volume lithium evidently serves merely as a catalyst.” available t o the reaction system consisted of a glass vessel which Later, in a paper devoted to preparative work could be submerged in a thermostated bath in a dewar. T h e utilizing this and similar resctions, Dolgov, et d.,* reaction vessel was connected t o a manometer for pressure measurements. The solution and the gas above it were stirred with a offered the view that the reaction is instead basemultipole magnetic stirrer enclosed in an evacuated glass jacket catalyzed, the base being the alkali amide formed by and driven by a variable speed motor. Two evacuated glass the reaction of the alkali metal with the solvent. tubes served as ampoule breakers. Each was sealed to a reagent ampoule (Li or Et3SiH)and contained a small bar magnet allowPresumably the possibility of base catalysis was suging it t o be manipulated by a horseshoe magnet outside the reacgested in part by kinetic studies of the reaction between tion vessel. The reaction vessel was thermostated in a stripa trialkylsilane and alcohol or water, beginning with silvered dewar. The bath liquid (usually methanol) was periodithat of Price. cally circulated through a cooling coil in a Dry Iceemethanol (1) C. A. Kraus and W . K. Selson, J . A m . Chem. Soc., 66, 195 (1934). (2) B. N Dolgov, N . P. Kharitonov, and M . G. Voronkov, Z h . Obshch. K h i m . , a i , 678 (1954). (3) F. P. Price, J . A m . Chem. Soc., 69, 2600 (1847).
slush in a separate dewar by means of a self-priming pump, which was controlled by a thermistor temperature controller (Yellow Springs Instrument Co., Model 63RB). For most of the experiments, temperature stability was about 1 0 . 1 ” . Total drift in
ARLENVISTE A N D HENRYTAUBE
1692
the course of several months of use was 0.2-0.3'. The temperature for the kinetic experiments was -31.5'. The folloMing procedure was followed in most of the experiments. If EtsSiF was t o be used, a known quantity was condensed into the reaction vessel. Ethylamine was distilled into the reaction vessel. A known weight of Li was subsequentljadded by breaking one of t h e ampoules. A known weight o f EtsSiH was added by breaking t h e other ampoule, thus starting the reaction. Pressure was measured a s a function of time after EtsSiH addition. After completion of the experiment, the ethylarnine was distilled back t o its storage vessel on the vacuum line.
Results A set of experiments was done in which the concentrations of Li and Et3SiH were varied, but no other reactant was deliberately added. If t h e reaction were catalyzed solely by Li and not a t all b y LiNHEt, this set of experiments should have established the rate law. In Table I , til, (specifically the first half-life) is recorded for most of the experiments of this set The wide scatter in the results is obvious. Pseudofirst-order kinetics (in Et3SiH) were roughly followed during the major portion of most runs, though appreciable deviations were sometimes observed particularly in the early portion of the runs.
Three experiments with EtsSiF present which were followed to complete consumption of EtaSiH are summarized in Table 111. (There was some possibility of contamination in expt. 19.) TABLE I11 KINETICEXPERIMENTS WITH EtaSiF ADDED;FOLLOWED TO COMPLETIOS [Lila X lo2, M [EtaSiHIo X lo2, Jf [EtsSiFl0 X lo2, JI t a t 57; reaction," min t a t 90co reaction," min Experiment
DIRECTKINETICEXPERIMENTS [Li]
a
x
102,
1 2 2 2 2 2 2 3 3 5
2 1 1 2 2 4 5 1 7 5
M
[EtsSiHIo
x
102, Ma
First
11/%,
min
62 4 10 52 0.5 1.1 60 12 1.1 15
6.7 7.6 5.8 7.2 8.1 8.6 7.0 7.4 13.5 7.3
Another set of experiments was done in which the procedure followed was essentially the same as in the ordinary kinetic experiments above, except that now some Et3SiF was added to the E t N H s before adding Li and EtsSiH. Two runs with EtsSiF present which were not followed to completion of consumption of EtaSiH are summarized in Table 11. In expt. 17, TABLE I1
KINETICEXPERIMENTS WITH EtsSiF ADDED;NOT FOLLOWED TO COMPLETION 2.7 8.0 19.0 0.0 113 17
3 1 9 3 0 61 266 281 20
3 9 12 8 0 42 121 137 21
For these same Et3SiF experiments, Table I V summarizes data concerning the extent of Hz production during the first 60 min. of reaction. TABLE 11=
60 MIX. FOR
[EtaSiFIo/ [ L i l ~ Expected Om, cm. APa if reaction of Li with E t S H 2 were complete, cm . Total A P b a t t = 60 min. (including increase before EtaSiH addition), cm ,
qAP8 L l , / Experlment
x
THE
EXPERIMENTS WITH EtaSiF ADDED
7 . 1 0 . 2 5 0 13 0 . 2 0 0 . 1 1 2 3 . 3 2 6 . 2 23 6 22.8 27.8
3.9
3 2
4.0
3 8
4.2
0.1
0.2
0.2
0.2
0.3
105
2 1
1 7
2 8
102, 'If 17
18
19
20
a Represents approximate lower limit for [LI] a t t See text
Subscript 0 indicates initial concentration
[Lila X lo2, M [EtsSiH]o X lo2, Jf [EtsSiFIo X lo2, M Final c'< reactiona Final 1, min. Experiment
3 1 9 1 0 40 194-216' 224 19
a iZpparent per cent reaction, taken as A P t A P m X 100 Readings had not been taken during this interval
DATAFOR t TABLE I
Vol. 86
2.4 9.7 0.59 1.3 241 18
a Apparent per cent reaction when experiment was halted taken as final h'?calcd. A P m X 100.
t h e blue color was observed to fade by t = 65-72 min. ( t is measured from the time of addition of EtaSiH to the solution of E t N H 2 ,Li, and EtaSiF).
21 =
60 niin
Discussion It is apparent that the results of the direct kinetic experiments without EtaSiF scatter widely (Table I ) The conclusion to be drawn from these results is that although there may perhaps be a Li-catalyzed path available to the reaction, it is certainly not the only available path I n principle, the question of the existence of a disinct Li-catalyzed reaction path in the case of a solution containing both Li and a strong base such as LiKHEt could be resol1Ted by observing the decrease in reaction rate brought about by removing either Li or LiNHEt from solution without disturbing the other As a practical matter, however, removing the LiNHEt is likely to be preferable; since Li is ordinarily expected to be in large excess over LiNHEt, the amount of any reagent used to remove Li must also be large in comparison to the amount of LiNHEt present, and the danger exists that low-level impurities in the reagent may consume significant amounts of L i S H E t (In fact, before carrying out the Et3SiF experiments. several experiments were done using Hg to try to determine the effect of extracting Li from the solution
May 5 , 1964
LITHIUMIN TRIETHYLSILANE-ETHYLAMINE REACTIOX
The rate was retarded in varying degrees, but since the Hg was present in very great excess over LiNHEt, these results presumably have little significance and will not be discussed further.) I n the series of experiments with EtsSiF present, i t was hoped that EtiSiF might consume strong base impurities such as LiNHEt, while leaving most of the Li untouched. The basis for trying Et3SiF was the observation by Gierut, et al.,4that an attempted analysis of Et3SiF for fluorine failed because Et3SiF did not react with Na in liquid N H 3 a t -33'. The first experiment of the Et3SiF series (Table 11) shows t h a t EtsSiF is not entirely unreactive toward Li in EtNH2. However, even with 0.19 M EtaSiF the 0.027 M Li was not consumed until about an hour after the EtpSiH was added. Since the Et3SiF concentrations used in the other Et&F experiments (Tables 11-IV) were only 2-4y0 as large as this, the direct reaction between Li and Et3SiF should be relatively slow for the rest of the EtsSiF series. In all of the EtgSiF experiments except the first, 0.11 5 [ E ~ ~ S ~ F ] O6/ [ 0.23. L ~ ] ~ Thus Li is present in substantial excess over EtsSiF. In Table IV is given an approximate lower limit for the Li concentration a t t = 60 min. for the EtsSiF series
This makes allowance for possible consumption of the Li. The actual Li concentration seems unlikely to be materially less than this value (with the possible exception of expt. 19) and may well be larger. These lower limits are about the same size as Li concentrations in the ordinary kinetic experiments of Table I. Thus direct comparison of half-lives seems fair, projecting the initial rate observed'up to t = 60 min. in expt. 18--21. T h e projected half-lives of more than 2000 min. in expt. 18-21 are more than 30 times as long as the longest half-life in Table I. Thus the Et3SiF experiments show t h a t Li by itself does not significantly catalyze the reaction of EtpSiH with EtNH2. Not only does the negligible rate in the early portion of these experiments show t h a t Li does not significantly catalyze the reaction, b u t the rather abrupt onset of reaction a t an accelerating rate after 2-4 hr. provides evidence t h a t LiNHEt or some similar strong base is a good catalyst for the reaction. A reasonable interpretation is the following: Li slowly reacts with EtNHz to give LiNHEt and Hz. As long as EtsSiF is present, this L i N H E t apparently is rapidly consumed by the EtsSiF, probably by means of the reaction EtaSiF
+ LiYHEt +EtsSiNHEt + LiF
and thus cannot serve as a catalyst. However, as this reaction proceeds, Et3SiF is slowly consumed. ( I t is probably somewhat more slowly consumed by direct reaction with Li.) When the Et3SiF has been used up, the L i N H E t starts to build up and can begin to catalyze the reaction. Thus the point a t which the reaction takes place a t an appreciable rate is taken to be the point a t which EtsSiF has been used up. (In (4) J. A , Gierut, F. J. Sowa, and J. A. Nieuwland, J . A m . Chem. S o i . , 18, 897 (1936).
1693
connection with Tables I1 and 111, the H 2 pressure increase associated with complete consumption of Et3SiF by LiNHEt would correspond to an apparent per cent reaction of less than 5% in all cases except expt. l i , ) Although the above analysis suggests that a major catalyst is produced by reaction of Li with the solvent EtNH2, identification of this catalyst as LiNHEt is admittedly incomplete. For example, it is possible t h a t after a short period of reaction, LiXHEt might be converted to LiEt3SiNEt, if EtsSiNHEt is sufficiently more acidic than EtXIlz. This point was not pursued further since our primary interest was in whether or not the lithium was an effective catalyst. For simplicity, however, the base catalyst under discussion will be referred to as LiNHEt. A very rough estimate can be made of the catalytic effectiveness of the base catalyst. In expt. 20, at t = 278 min., the specific rate constant corresponds to t l I 2 = 2.4 min. Also, it appears that the pressure started increasing a t an appreciable rate a t about t = 265 min. If i t is assumed that the rate of LiNHEt production from 265 to 278 min. is approximately the same as during the first 265 min. of the reaction, judged by the pressure increase up to t = 265 min., then it appears t h a t the til, = 2.4 min. is due to a LiNHEt concentration of about 4 X M . Since in the experiments reported, [Li] was generally about (2-5) X M , this indicates t h a t a concentration of LiNHEt only 1 or 2y0 as large as t h a t of Li could, in these experiments, give a half-life of the order of 2-3 niin. Small, irreproducible amounts of LiNH E t would be a logical impurity in Li-EtNH2 solutions. Thus the active catalyst in the reaction between Et3SiH and EtNHz is not Li, but apparently rather is a strongly basic impurity such as LiNHEt. Recently there has been a revival of interest in a modification of the cavity model for metal-ammonia solutions, in the form of an ion-quadrupole model of the M P d i r ~ e r . ~If, ~such a picture applies also to solutions of lithium in ethylamine, i t is interesting to note t h a t the solvated electron pair appears to act as a substantially weaker base than LiNHEt in the catalyzed EtpSiH-EtNHz reaction and apparently also in the reaction with EtpSiF. A similar comparison holds in a t least one instance in metal-ammonia solutions, since the NH3-NH2- exchange proceeds much more rapidly than the decomposition reaction of alkali metal-ammonia ~ o l u t i o n s* . ~The reactions to be compared are KH2-
+
r\"3
= SHs
+ KH2-
and 2e- t XHI = H -
+ XH2-
(H- once formed would presumably rapidly form H2). Acknowledgments.-Fellowship support provided for A. V. by the National Science Foundation is gratefully acknowledged. This research was supported by the Atomic Energy Commission, Contract A T ( 11-1)-378. ( 5 ) M . Gold, W. L.Jolly, and K . S. Pitzer, ibid.,84, 2284 (1962). (6) M. Gold and W L Jolly, I n o r g . Chem., 1, 818 (1962). (7) R . A . O g g . Discussions F a r a d a y Soc , 17, 215 (19,541. ( 8 ) D . M. Yost and H . Russell, J r . , "Systematic Inorganic Chemistry," Prentice-Hall. Inc., N e w York. N. Y . , 1946, p. 137.
